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Formation of manganese oxides on early Mars due to active halogen cycling

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Nature Geoscience
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https://doi.org/10.1038/s41561-022-01094-y
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Formationofmanganeseoxidesonearly
Ma...
Nature Geoscience
Article https://doi.org/10.1038/s41561-022-01094-y
durations.ThesolubilityofO2 islow,andifthesystemwerec...
Nature Geoscience
Article https://doi.org/10.1038/s41561-022-01094-y
100 mmol l−1
Mn(ii) and 10 mmol l−1
bromate in chlori...
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Formation of manganese oxides on early Mars due to active halogen cycling

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In situ rover investigations on Mars have discovered manganese oxides as
fracture-filling materials at Gale and Endeavour craters. Previous studies
interpreted these minerals as indicators of atmospheric oxygen on early
Mars. By contrast, we propose that the oxidation of manganese by oxygen
is highly unlikely because of exceedingly slow reaction kinetics under
Mars-like conditions and therefore requires more reactive oxidants. Here we
conduct kinetic experiments to determine the reactivity of the oxyhalogen
species chlorate and bromate for oxidizing dissolved Mn(ii) in Mars-like
fluids. We find that oxyhalogen species, which are widespread on the
surface of Mars, induce substantially greater manganese oxidation rates
than O2. From comparisons of the potential oxidation rates of all available
oxidants (including reactive oxygen species peroxide and superoxide),
we suggest that the oxyhalogen species are the most plausible manganese
oxidants on Mars. In addition, our experiments precipitated the manganese
oxide mineral nsutite, which is spectrally similar to the dark manganese
accumulations reported on Mars. Our results provide a feasible pathway to
form manganese oxides under expected geochemical conditions on early
Mars and suggest that these phases may record an active halogen cycle
rather than substantial atmospheric oxygenation.

In situ rover investigations on Mars have discovered manganese oxides as
fracture-filling materials at Gale and Endeavour craters. Previous studies
interpreted these minerals as indicators of atmospheric oxygen on early
Mars. By contrast, we propose that the oxidation of manganese by oxygen
is highly unlikely because of exceedingly slow reaction kinetics under
Mars-like conditions and therefore requires more reactive oxidants. Here we
conduct kinetic experiments to determine the reactivity of the oxyhalogen
species chlorate and bromate for oxidizing dissolved Mn(ii) in Mars-like
fluids. We find that oxyhalogen species, which are widespread on the
surface of Mars, induce substantially greater manganese oxidation rates
than O2. From comparisons of the potential oxidation rates of all available
oxidants (including reactive oxygen species peroxide and superoxide),
we suggest that the oxyhalogen species are the most plausible manganese
oxidants on Mars. In addition, our experiments precipitated the manganese
oxide mineral nsutite, which is spectrally similar to the dark manganese
accumulations reported on Mars. Our results provide a feasible pathway to
form manganese oxides under expected geochemical conditions on early
Mars and suggest that these phases may record an active halogen cycle
rather than substantial atmospheric oxygenation.

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Formation of manganese oxides on early Mars due to active halogen cycling

  1. 1. Nature Geoscience naturegeoscience https://doi.org/10.1038/s41561-022-01094-y Article Formationofmanganeseoxidesonearly Marsduetoactivehalogencycling Kaushik Mitra 1,3 , Eleanor L. Moreland 1,4 , Greg J. Ledingham1 & Jeffrey G. Catalano 1,2 InsituroverinvestigationsonMarshavediscoveredmanganeseoxidesas fracture-fillingmaterialsatGaleandEndeavourcraters.Previousstudies interpretedthesemineralsasindicatorsofatmosphericoxygenonearly Mars.Bycontrast,weproposethattheoxidationofmanganesebyoxygen ishighlyunlikelybecauseofexceedinglyslowreactionkineticsunder Mars-likeconditionsandthereforerequiresmorereactiveoxidants.Herewe conductkineticexperimentstodeterminethereactivityoftheoxyhalogen specieschlorateandbromateforoxidizingdissolvedMn(ii)inMars-like fluids.Wefindthatoxyhalogenspecies,whicharewidespreadonthe surfaceofMars,inducesubstantiallygreatermanganeseoxidationrates thanO2.Fromcomparisonsofthepotentialoxidationratesofallavailable oxidants(includingreactiveoxygenspeciesperoxideandsuperoxide), wesuggestthattheoxyhalogenspeciesarethemostplausiblemanganese oxidantsonMars.Inaddition,ourexperimentsprecipitatedthemanganese oxidemineralnsutite,whichisspectrallysimilartothedarkmanganese accumulationsreportedonMars.Ourresultsprovideafeasiblepathwayto formmanganeseoxidesunderexpectedgeochemicalconditionsonearly Marsandsuggestthatthesephasesmayrecordanactivehalogencycle ratherthansubstantialatmosphericoxygenation. ConcentratedoxidizedmanganesedepositsonMarsoccurinveinsor fractureatGale(>25 wt%)(refs.1–3 )andEndeavour(>2.3 wt%)(refs.4,5 ) craters.Thesemineralsserveasindicatorsofpastredoxconditionsthat promoted abiotic oxidation of Mn(ii). Detailed understanding of the redox state of aqueous solutions, the oxidants that facilitated Mn(ii) oxidationandtheprocessesformingmanganeseoxidesarecriticalto geochemically constrain past environmental conditions and provide essential insight into the habitability of early Mars. Manganese oxides occurring in fracture fills, such as in Gale and Endeavourcraters,requiretransportinafluidandchemicalfractiona- tion from iron. Weathering of ferromagnesian silicates (for example, olivine)istheanticipatedsourceofdissolvedMn(ii).Althoughmolecu- lar oxygen (O2) is currently considered the primary Mn(ii) oxidant on Mars1,2,6,7 , slow reaction kinetics makes O2 an implausible Mn(ii) oxidant.Paststudiesat25 °CdemonstrateincrediblyslowMn(ii)oxi- dationrateswithnoobservableoxidationfor>7 yratpH 8.4in0.2 bar O2 (ref. 8 ). In addition, pO2 in the present Martian atmosphere is low (~10−5 bar)9 althoughitperiodicallyreached0.001to0.05 baronearly Mars6 . The rate of Mn(ii) oxidation by O2 decreases linearly with pO2 (refs.8,10–13 ),furtherslowingreactionratesonMars. OwingtoequilibrationwiththeCO2-rich(0.5 bar)atmosphere,the pHofatmosphericallyconnectedaqueoussystemsonMarswasweakly acidic (pH < 6.3) during late Noachian to early Hesperian periods14 (SupplementaryDiscussionsection1).Althoughwater–rockreaction mayneutralizetheacidityfromatmosphericCO2 andgeneratealkaline systems in the Martian subsurface15 , these processes consume all O2 while conditions are still acidic (Extended Data Fig. 1). The role of O2 in manganese oxide formation on Mars is further challenged by the Received: 24 January 2022 Accepted: 26 October 2022 Published online: xx xx xxxx Check for updates 1 Department of Earth and Planetary Sciences, Washington University, St Louis, MO, USA. 2 McDonnell Center for the Space Sciences, Washington University, St Louis, MO, USA. 3 Present address: Department of Geosciences, Stony Brook University, Stony Brook, NY, USA. 4 Present address: Department of Earth, Environmental and Planetary Sciences, Rice University, Houston, TX, USA. e-mail: catalano@wustl.edu
  2. 2. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y durations.ThesolubilityofO2 islow,andifthesystemwereclosedafter initialO2 entrythenafullysaturated1-cm-widefracturecouldgenerate a MnOOH coating only 20 nm thick for pO2 of 0.02 bar and 3 nm thick for pO2 of 0.003 bar. These are orders of magnitude too thin to have producedtheconcentrationsobservedviaX-rayspectrometry1–4 .There arethussubstantialkineticandthermodynamicbarrierstoO2 having servedastheoxidantthatproducedfracture-fillingmanganeseoxide depositsonMars. Abundant oxychlorine species (perchlorate and chlorate) have beenfoundubiquitouslyonthesurfaceofMars19,20 andwereprobably present since ~4 billion years ago (Ga) (ref. 21 ). Chlorate (ClO3 − ) can readily oxidize dissolved iron(ii) orders of magnitude faster than O2 orviaultravioletphoto-oxidationunderawiderangeofMars-relevant conditions,includinginacidicfluidswheretheseotheroxidationpath- waysareexceedinglyslow22–24 .ChlorateoxidizesMn(ii)at~160 °C(ref. 25 ) but previous work has not studied such reactions under ambient conditions.AssessmentofpotentialMn(ii)oxidationratesattempera- tures relevant to Mars through analogy with Fe(ii)24 are not possible because these species display reaction rates with other oxidants that varywidely,althoughMn(ii)oxidationratesaregenerallyslower11 . Oxychlorine-forming processes on Mars also produce oxybro- mine compounds since halogens share similar chemistry19 . Mars is a halogen-rich planet26 , with chlorine and bromine concentrations in bulkMarsabout4timesgreaterthanonEarth26,27 .Bromineconcentra- tions on the Martian surface are 1–10% of the total chlorine concen- tration27 . Bromine was detected in all samples analysed by the Spirit, Opportunity5 andCuriosityrovers3,27 ,withthehighestconcentrations in the manganese-rich Dillinger member at Gale crater3 . In addition, bedrockshows~80%enrichmentinbrominecomparedwithGalecrater soil28 . Bromine volatilizes and oxidizes faster than chlorine, thereby experiencing more rapid production–consumption cycles on Mars andoccurspredominantlyasbromate(BrO3 – )(refs.19,29 ).Whilebromate hasbeenshowntooxidizedissolvedMn(ii)atapproximatelypH –0.5 (ref. 30 ), no studies have explored its reactivity with Mn(ii) under less extreme chemical conditions. In this Article, we show that the rate of Mn(ii) oxidation by oxyhalogens is orders of magnitude faster than by O2 in Mars-like conditions and that manganese oxides indicate an activehalogencycleonMars. Oxidationofmanganesebyoxyhalogensin Mars-relevantfluids We conducted laboratory experiments to determine the reactivities of the oxyhalogen species, chlorate and bromate, towards oxidizing dissolved Mn(ii). The experiments were conducted in magnesium chlorideandmagnesiumsulfatebackgroundsolutionsatpH ~3,5and 7 to emulate Mars-like fluids. We studied the rate of Mn(ii) oxidation inkineticexperimentsbymonitoring[Mn(ii)]andpHinsystemscon- taining individual oxidants, either bromate or chlorate, and in mixed systemscontainingbothoxidants(seeMn(II)oxidationbyoxyhalogen species for details). TheresultsshowthatMn(ii)oxidationbychlorate,althoughther- modynamically favourable (Extended Data Table 2), is not observed on the timescales of weeks to months (Supplementary Fig. 1). The negligiblepHchangeintheseexperimentsindicatesthat<0.1%ofthe dissolvedMn(ii)oxidized.Bycontrast,solutionscontainingbromate displayed extensive Mn(ii) oxidation within six to eight weeks (Fig. 2 andExtendedDataTable3).ThepHofthesystemsdecreasedtobelow 2 in all experiments, regardless of the initial pH (Fig. 2). Mn(ii) in sul- fatesystemsdisplayedsloweroxidationthaninchloridesystemsdue to greater complexation of Mn(ii), decreasing the concentration of the free Mn2+ ion (Supplementary Fig. 2). In all experiments, fuming brownish-red gas volatilized from the solutions upon de-capping the reactors,implyingincompletebromatereductionbyMn(ii),forming Br2 rather than Br− . This also suggests a redox-mediated volatiliza- tion mechanism for bromine on Mars. Batch experiments containing exceedinglyslow(andoftenunmeasured)kineticsofMn(ii)oxidation belowpH ~6(refs.8,10–12,16,17 )(ExtendedDataTable1).Applicationofan existingratemodel11 showsthatpotentialpO2 conditionsonearlyMars (0.05to10−5 bar(ref.6 ))requirehundredstomillionsofyearstooxidize manganese (Fig. 1). Focusing specifically on the period of Gale crater formation, sedimentation and exhumation, peak O2 concentrations ranged from 0.003 to 0.02 bar (ref. 6 ). These O2 levels yield a half-life (t1/2)ofMn(ii)oxidationof840–5,600 yratpH 6and7,600–51,000 yrat pH 5(seeSupplementaryDiscussionsection2forfurtherdiscussion). If the pH was as low as 4, which is plausible given the occurrence of akaganeiteintheunitsurroundingtheGalecraterfractures18 ,thenman- ganese oxidation by O2 was not even thermodynamically favourable. Furthermore,thesesubsurfacefracturesystemswouldhaveneededto remainhydrologicallyconnectedtotheatmosphereforthesefulltime 4.0 4.5 5.0 5.5 pH 6.0 6.5 7.0 7.5 10 8 10 7 10 6 10 5 10 4 10 3 10 2 10 1 10 0 t 1/2 (yr) t 1/2 (yr) 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 pH pH limit for 0.5 bar CO 2 10 –5 10 –2 0.03 O2 (bar) Water availability Rhodochrosite saturation 10 –4 10 –3 0.05 pH limit for 0.5 bar CO 2 10 –5 10 –2 0.03 Water availability Rhodochrosite saturation 10 –4 10 –3 0.05 a b O2 (bar) 10 8 10 7 10 6 10 5 10 4 10 3 10 2 10 1 10 0 Fig.1|TimescalesofMn(ii)oxidationbyoxygen. a,b,Thehalf-life(inyears)for Mn(ii)(1 mmol l−1 )oxidationbyO2 plottedasafunctionoftheinitialpHat25 °C. Allsimulationswerebufferedwith0.5 barCO2 andO2 settothespecifiedpartial pressures.CalculationsconsideredbothfixedpHsystems(a)andfree-driftpH systemsinwhichthepHisallowedtovaryinresponsetoH+ generationduring Mn(II)oxidation(b).TheverticallinesatpH 6.34demarcatethepHabove whichcarbonatemineralsmustbepresent.ThedashedverticallineatpH 5.56 demarcatesthepHofrhodochrositesaturation,abovewhich1 mmol l−1 Mn(ii) isnotfullysoluble.Thehorizontallinedemarcatesthemaximumtime(1million years)expectedforliquidwateravailabilityonMars6 .Conditionsforwhich thermodynamicslimitMn(ii)oxidationto<50%orforwhichMn(ii)oxidation cannotoccur(SupplementaryTable1)arenotrepresentedinthefigure.
  3. 3. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y 100 mmol l−1 Mn(ii) and 10 mmol l−1 bromate in chloride-rich systems showed considerable Mn(ii) oxidation after ~650 days at pH 3.48 and 7.53 (Extended Data Table 4), which is consistent with the expected kinetic behaviour of the reaction if the rate is first order with respect tobromateconcentration31 . OxyhalogensproduceMn(iii/iv)oxidensutite Batch experiments were conducted to investigate the mineral prod- ucts formed when bromate oxidizes Mn(ii) in Mars-relevant fluids (ExtendedDataTable5).Mn(ii)oxidationbybromateyieldsdarkbrown to black manganese coatings inside the reactor walls (Extended Data Fig. 3), which are confirmed by X-ray diffraction (XRD) to be the min- eral nsutite (γ-MnO2) (Fig. 3a,d, Extended Data Fig. 2). Nsutite is a dis- ordered, nanoscale intergrowth of the MnO2 polymorphs pyrolusite andramsdelliteandcanincludeMn(iii)substitutionsforMn(iv).X-ray photoelectronspectroscopy(XPS)resultsshowthatthesolidsformed in sulfate fluids and at higher pH values generally contained greater Mn(iii)contents(Fig.3b,eandSupplementaryTable3).Mixed-valent Mn(iii/iv)mineralsarethereforemorelikelytoformthanpureMn(iv) in sulfate-rich fluids common on Mars. The visible and near infrared (VNIR) reflectance spectra of the precipitates have characteristically lowreflectancevalues(Fig.3c,f).Thespectraaresimilartothemanga- nese oxide rock coatings at Endeavour crater4 (Extended Data Fig. 4). TheslightlybrighterspectraonMarsareprobablycausedbythepres- ence of magnesium sulfate or other salts intermixed with manganese oxides.Nsutiteadsorbsmetals(forexample,Pb,Co,Cu)32 similarlyto othermanganeseoxides33,34 .Itsformationisthusnotinconsistentwith thecorrelationofMnwithCuandNiatGalecrater(Stephen)1,2 ,between Ni and Mn at Endeavour crater (Pinnacle Island and Stuart Island)4 . Comparisonofplausiblemanganeseoxidantson Mars HomogeneousMn(ii)oxidationbyO2 hasexceedinglyslowkineticsbelow pH~7(ref.35 ),withexperimentsreportinglittletonooxidationovertime- scales of months to years at 0.2 bar O2 and 20 °C (refs. 8,10–12,16,17,36 ). The order of Mn(ii) oxidation by O2 with respect to [OH− ] is ~2 (refs. 8,10–13 ). TheMn(ii)oxidationratethusdecreases~100×witheachunitdecreasein pH13 ,makingthereactionlessfavourableandprobablyunfeasibleatacidic pH35 .Heterogeneousredoxreactionssuchassurfaceandelectrochemi- cal catalysis promote faster Mn(iii/iv) mineral formation than during oxidationinhomogeneoussolution16,17 .Studiesreport22%,51%and63% Mn(ii)oxidationwithin60daysinthepresenceofmagnetite,two-line ferrihydrite and goethite, respectively, at pH 7 and 0.2 bar O2 at 25 °C (ref.17 ).However,notethatsurfacecatalysisdependsonMn(ii)adsorp- tiononmineralsurfaces,whichdropsprecipitouslybelowpH 6(ref.37 ), eliminatingthisoxidationpathwayunderweaklyacidicconditions.No studieshavedemonstratedtheabioticoxidationofMn(ii)byO2 below pH6atambienttemperaturesforeitherhomogeneousorheterogeneous systems.ThehighpCO2 ofearlyMarsrequiresthatmostwatersincontact withtheatmosphere,thesolesourceofO2,wouldhaveapH < 6(ref.14 ). TherateofMn(ii)oxidationbyO2,bothhomogeneousandsurface catalysed, is slower than by bromate (Figs. 1 and 2 and Extended Data 0 20 40 60 80 100 0 15 30 45 60 0 1 2 3 4 5 6 7 0 15 30 45 60 0 20 40 60 80 100 0 15 30 45 60 0 1 2 3 4 5 6 7 0 15 30 45 60 0 20 40 60 80 100 0 15 30 45 60 0 1 2 3 4 5 6 7 0 15 30 45 60 (Mn( II )) (mmol l –1 ) (Mn( II )) (mmol l –1 ) (Mn( II )) (mmol l –1 ) pH pH pH (Mn(II)) (Mn(II)) (Mn(II)) a b c (Mn( II )) (mmol l –1 ) (Mn( II )) (mmol l –1 ) Time (d) Time (d) Time (d) Time (d) (Mn( II )) (mmol l –1 ) Time (d) Time (d) Time (d) Time (d) Time (d) Time (d) Time (d) Time (d) d e f (Mn(II)) (Mn(II)) (Mn(II)) 0 20 40 60 80 100 0 15 30 45 60 0 15 30 45 60 0 20 40 60 80 100 0 15 30 45 60 0 1 2 3 4 5 6 7 0 15 30 45 60 0 20 40 60 80 100 0 15 30 45 60 0 1 2 3 4 5 6 7 0 15 30 45 60 0 1 2 3 4 5 6 7 pH pH pH Fig. 2 | Rates of Mn(ii) oxidation by bromate. a–f, Dissolved Mn(ii) concentration and pH versus time in systems containing approximately 100 mmol l−1 Mn(ii) with 100 mmol l−1 bromate in 100 mmol l−1 MgCl2 (a–c) at initial pH 2.95 (a) 4.77 (b) and 6.91 (c) and in 100 mmol l−1 MgSO4 (d–f) at initial pH 3.03 (d), 4.92 (e) and 6.87 (f) at 24 °C.
  4. 4. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Table1).AssumingaratelawthatisfirstorderwithrespecttoMn(ii)and bromate31 , it would take 120, 400 and 900 days to oxidize about 20%, 50% and 90% Mn(ii), respectively, in a system containing 24 mmol l−1 Mn(ii)and10 mmol l−1 bromateatpH 7at25 °CwiththepHdecreasing to2.50,2.10and1.82,respectively.Bycontrast,onlyabout8%Mn(ii)is oxidizedin120daysby0.2 barO2 undersimilarreactionconditions16 . In heterogeneous systems containing ferrihydrite, a 30–50% drop in Mn(ii) is observed within 120 days (Extended Data Table 1)16,17 . Note that rates on Mars would be at least an order of magnitude slower because of the lower pO2. For reference, 0.21 bar O2 would produce Intensity (counts) pHinitial 2.96 2.49 5.15 5.95 6.65 γ-MnO2 γ-MnO2 γ-MnO2 γ-MnO2 0 1,000 2,000 3,000 4,000 5,000 6,000 7,000 8,000 15 25 35 45 55 65 a Intensity (counts) 0 1,000 2,000 3,000 4,000 5,000 6,000 15 25 35 45 55 65 pHinitial 2.86 3.69 4.30 6.72 7.59 γ-MnO2 γ-MnO2 γ-MnO2 γ-MnO2 0 0.04 0.08 0.12 0.16 0.20 0.24 0.28 0.40 0.90 1.40 1.90 2.40 Wavelength (µm) Reflectance (+0.02) (+0.06) (+0.09) (+0.16) pHinitial 3.69 2.86 4.30 6.72 7.59 e f Reflectance (+0.05) (+0.10) (+0.15) (+0.20) pHinitial 2.96 2.49 5.15 5.95 6.65 0 0.06 0.12 0.18 0.24 0.30 0.36 0.40 0.90 1.40 1.90 2.40 Wavelength (µm) °2θ (Cu Kα) °2θ (Cu Kα) b c d 0 500 1,000 1,500 2,000 45 47 49 51 53 pHinitial 3.69 4.30 7.59 Binding energy (eV) Arbitrary units Mn(III) Mn(IV) Mn(II) 50.6:49.4:0 48.8:51.2:0 33.7:66.3:0 (IV):(III):(II) 0 500 1,000 1,500 2,000 45 47 49 51 53 2.96 5.15 5.95 Binding energy (eV) Arbitrary units Mn(III) Mn(IV) Mn(II) 100:0:0 43:57:0 (IV):(III):(II) 76:24:0 pHinitial Fig.3|MineralproductsofMn(ii)oxidationbybromate. a–f,TheXRD patterns(a,d),XPSspectra(b,e)andVNIRreflectancespectra(c,f)ofthe mineralsprecipitatedfollowingtheoxidationofdissolvedMn(ii)bybromate inmagnesiumchloride(a–c)andmagnesiumsulfate(d–f)fluids.Patternsand spectraareoffsetvisuallyforclarity.ThelabelsindicatetheinitialpHofthe samples.SeeExtendedDataTable5forcompletesampledetails.Diagnostic peaksintheXRDscans(a,d)arelabelledandindicatedbydashedlines.TheMn 3pXPSspectra(b,e)includemineralstandards(dottedlines)corresponding toMn(iv),Mn(iii)andMn(ii).ThepercentagesofMn(iv),Mn(iii)andMn(ii)in eachsampleareindicated(seeSupplementaryTable3forfittingparameters). ThedashedlineoverlyingtheVNIRspectra(c,f)indicatesthepositionofthe 1.92 μmabsorptionband.TheverticaloffsetsoftheVNIRspectraareindicatedin parentheses.
  5. 5. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y <0.1%oxidationofMn(ii)undertheconditionsofourexperimentswith oxyhalogensthatcompletelyoxidizeMn(ii)atpH ≤ 7(ExtendedData Table 6). As demonstrated by the results, bromate can oxidize Mn(ii) downtopH ~3whereMn(ii)oxidationbyO2 becomesabout108 ×slower (100× decrease in rate with each unit pH decrease). Unlike bromate, chlorate produced no observable Mn(ii) oxidation in homogeneous systems in this study. Such reactions could possibly be accelerated in a heterogeneous system via surface catalysis, analogous to the faster reactionbetweenMn(ii)andO2. Arecentpaper6 suggestedO2 asaneffectiveMn(ii)oxidantonMars during bolide impacts on the surface. Thermal and kinetic modelling using extrapolated data36 showed that at 50–75 °C, dissolved Mn(ii) can be readily oxidized (within ~tens of days) by 0.03 bar O2 in 0.5 bar CO2 atpH 8.However,maintainingapH 8fluidisunachievableundera 0.5 barCO2 atmosphereasthiswouldrequireadissolvedbicarbonate concentration in excess of 1 mol l−1 (See Supplementary Discussion section 5 and Supplementary Table 4 for details). Dissolution of CO2 into alkaline fluids produced by water–rock reaction would decrease the pH to between 4.56 and 6.95, depending on the temperature and assumptionsofthereactionsbufferingpH.Owingtothesecond-order rate dependence on [OH- ], the rate of reaction will decrease between 125× and 7,600,000×, requiring decades to millennia to appreciably oxidize Mn(ii) using O2 as an oxidant. In addition, extrapolation of the rate constant for Mn(ii) oxidation by O2 to 75 °C in the study6 was potentially erroneous because of imprecise rate data at 25 and 37 °C, as previously cautioned36 . Therefore, Mn(ii) oxidation by O2 during thermal transients associated with bolide impacts6 is many orders of magnitudeslowerthanestimatedinthenotedpreviousstudybecause thepHconsideredisunattainableunderhighpCO2. Other species are ineffective or implausible oxidants of Mn(ii) in subsurface fracture systems. Hydrogen peroxide (H2O2) reduces Mn(iii/iv)toMn(ii)34,38 .Whilethereactiveoxygenspeciessuperoxide (O2 − )(refs.39,40 )doesoxidizeMn(ii)41 ,theresultingH2O2 mayeffectively negatetheoxidationbyO2 − .Otheroxychlorinespecies,suchaschlorite (ClO2 – ),hypochlorite(ClO− )andchlorinedioxide(ClO2)gasaswellasan arrayofradicals,areprobablyproducedasshort-livedintermediateson Mars42,43 andareknowntorapidlyoxidizedissolvedMn(ii)44,45 .However, theirnon-detectiononMars20 indicatesthattheydonotaccumulateon thesurface,unlikemoreoxidizedspecies,andtheirtransientnaturein water46 suggeststhattheymaydecomposebeforereachingsubsurface fracturesystemswheremanganeseoxidesformed1,2,4 .Notethatwedo notdiscountthepotentialroleofthesereactiveoxychlorinespeciesin oxidizingMn(ii)intheshallowsubsurfaceofMars.Thefinelydissemi- natedmanganeseoxidesthatwouldresulthavenotbeenobservedto date but warrant investigation as active missions and future sample returnenableidentificationviamicroanalysis.However,bromateisthe firstspeciesidentifiedthatplausiblyoccursonMarscapableofbeing transported into subsurface fracture systems and oxidizing Mn(ii) at measurable rates under neutral to acidic conditions. Oxyhalogens are thus more likely the primary oxidants of manganese on ancient Mars than O2, and manganese oxides in sedimentary units cannot be interpretedasreliableindicatorsofearlyMartianatmosphericoxygen. Manganeseoxidesasasignatureofhalogen cyclingonMars Manganese oxides in fractures on Mars formed in weakly acidic flu- ids.TheKimberlyformationthathoststhefracture-fillingmanganese oxidesinGalecratercontainsakaganeiteandpossiblyjarosite18 ,which require pH values of 5 or lower to form. Crater rim fractures in Gale craterthatcontainmanganeseoxidesalsocontainmagnesiumsulfates andwereprobablyhydrologicallyconnectedtofluidsfromtheBurns formation, which was also acidic. Mineralogical data are lacking for the manganese oxide fracture-filling material, which inhibits further evaluation of the formation pH. However, acidification associated Groundwater table HCl Cl (Deposit on the surface) Chloride/ bromide salts HCl, HBr Fe(II) Fe(III) P r e c i p i t a t i o n Mineral transformation Fe(III) Br Chlorate/ bromate salts Dissolutio n Groundwater interaction with basalt F e ( I I ) a n d M n(II)-rich groundwater M n(II) M n(III/IV) Precipitation B r O 3 – B r – Chloride/ bromide salts E v a p oration H2/CH4 H2O/CO2 CO3 2– BrO3 – Br– ? Fractures HBr ClOx (ClO3 – , ClO3 – , ClO2) Atmospheric pathways of oxyhalogen formation Surficial pathways of oxyhalogen formation (for example, D, E, F) ClO3 – /BrO3 – Oxyhalogen brines Tentative destruction of reduced atmospheric gases Cl– ClO4 – /ClO3 – /ClO2 Mn(III/IV) oxides Fe(III) oxides, sulfates Halide salts recycled to produce oxyhalogens (deposited) (B) (A) (C) (G) (for example, H, I) (J) (A) Ref. 51 (B) Ref. 52 (C) Ref. 56 (D) Ref. 53 (E) Ref. 54 (F) Ref. 55 (G) Ref. 22 (H) Ref. 57 (I) Ref. 58 (J) Ref. 50 BrO3 – Cl– /Br– Fig.4|PotentialpathwaysofhalogencyclingonMars,includingreactions withironandmanganese.HalogenreleasedfromtheinteriorofMarsserves astheprimarysourceontheMartiansurface51 .Oxyhalogenspecies,especially perchlorate,chlorateandbromate,areproducedbyatmospheric52 andsurface processes53–55 .TheseaccumulateonthesurfaceofMarsuntildissolvinginwater andbeingtransportedintostandingbodiesofwaterorpercolatingdownward intothesubsurfaceasbrines.Oxyhalogensinstandingbodiesofwatermayserve asoxidantsofdissolvedFe(ii)22 andformironoxidemineralsaswellaschloride andbromide,whichuponevaporationcouldformhalidesalts56 .Oxyhalogen brinesmayinteractwithdissolvedMn(ii)infractureandveinsandforminsoluble manganese(iii/iv)oxides.Reducedgases(forexample,H2 andCH4)inthe atmospheremaypotentiallyreactwithoxyhalogens,decomposingthesegases andgeneratingchlorideandbromide50 .Thehalidesaltsformedfrommultiple processesundergooxidationtorejuvenatetheoxyhalogenpool57,58 .
  6. 6. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y withanatmospherewithhighpCO2 demonstratesthattheseprobably formed at pH 6.5 or lower. Accordingtotheprecedingcomparisonofoxidantsandtheweakly acidic pH conditions associated with the fracture-filling manganese oxide deposits on Mars, oxyhalogen species produce far greater oxi- dation rates than is possible with O2. This indicates that manganese oxides in fractures on Mars are indicators of past halogen cycling (Fig. 4). We propose that manganese oxides were formed when dis- solved Mn(ii) was oxidized by downward-percolating oxyhalogen brines in the Martian subsurface rather than by O2-saturated waters. Halogensundergophotochemicaloxidationviavariousatmospheric and surficial processes to produce oxyhalogens such as perchlorate, chlorateandbromate.Subsequentreductionordecompositionofoxy- halogensproduceschlorideandbromide,whichcanbephoto-oxidized again,therebyformingahalogencycle.ThevolatilizationofBr2,formed duringbromatereductionbyMn(ii),couldleadtolessbromineinthe fluid,therebydecouplingmanganeseandbromineconcentrationsin sedimentsuponfluidevaporation.Inaddition,Br– issolubleandwould concentrate in places where the fluid evaporated and not necessarily where poorly soluble manganese oxides precipitate. Halogens are inter-connected to the redox cycling processes of manganese and irononMars21,47 asFe(ii)andMn(ii)canplayactiverolesindrivingthe reductionportionofthehalogencycle22,23,48 .WhileFe(ii)hasmultiple plausible oxidation routes on Mars beyond the oxyhalogens, Mn(ii) oxidationisslowforallplausibleoxidantsexceptbromate. Chlorine stable isotopes in apatites in Martian meteorites have been interpreted to indicate an active chlorine cycle on Mars dating back to ~4 Ga21 . It is unclear whether this cycling operated continu- ously or was inhibited during periods with reducing atmospheres containing H2 or CH4 (ref. 6 ). Reactions between reduced gases and oxychlorinespeciesdonotoccurinaqueoussystemswithoutmicro- bial activity49 but have been observed with ultraviolet-irradiated perchloratesalts50 .Chlorineandbrominecyclingaredrivenbysimilar processesandthereforeoperatedintandem.Manganeseoxidemin- erals directly observed in ancient sedimentary rocks1,2,4,5 by rovers provideaninsiturecordofactivehalogencyclingonearlyMars.The kinetically feasible production of manganese oxides proposed here suggests a complex and active Martian halogen cycle and advocates closer and more detailed geochemical study of halogens and oxyh- alogen species on early Mars. Onlinecontent Anymethods,additionalreferences,NaturePortfolioreportingsum- maries, source data, extended data, supplementary information, acknowledgements,peerreviewinformation;detailsofauthorcontri- butionsandcompetinginterests;andstatementsofdataandcodeavail- abilityareavailableathttps://doi.org/10.1038/s41561-022-01094-y. References 1. Lanza, N. L. et al. High manganese concentrations in rocks at Gale crater, Mars. Geophys. Res. Lett. 41, 5755–5763 (2014). 2. Lanza, N. L. et al. Oxidation of manganese in an ancient aquifer, Kimberley formation, Gale crater, Mars. Geophys. Res. Lett. 43, 7398–7407 (2016). 3. Berger, J. A. et al. Elemental composition and chemical evolution of geologic materials in Gale crater, Mars: APXS results from Bradbury landing to the Vera Rubin ridge. J. Geophys. Res. Planets 125, e2020JE006536 (2020). 4. Arvidson, R. E. et al. High concentrations of manganese and sulfur in deposits on Murray Ridge, Endeavour crater, Mars. Am. Mineral. 101, 1389–1405 (2016). 5. Mittlefehldt, D. W. et al. Diverse lithologies and alteration events on the rim of Noachian‐aged Endeavour crater, Meridiani Planum, Mars: in situ compositional evidence. J. Geophys. Res. Planets 123, 1255–1306 (2018). 6. Wordsworth, R. et al. A coupled model of episodic warming, oxidation and geochemical transitions on early Mars. Nat. Geosci. 14, 127–132 (2021). 7. Liu, Y. et al. Manganese oxides in Martian meteorites Northwest Africa (NWA) 7034 and 7533. Icarus 364, 114471 (2021). 8. Diem, D. & Stumm, W. Is dissolved Mn2+ being oxidized by O2 in absence of Mn-bacteria or surface catalysts? Geochim. Cosmochim. Acta 48, 1571–1573 (1984). 9. Trainer, M. G. et al. Seasonal variations in atmospheric composition as measured in Gale crater, Mars. J. Geophys. Res. Planets 124, 3000–3024 (2019). 10. Davies, S. H. & Morgan, J. J. Manganese (ii) oxidation kinetics on metal oxide surfaces. J. Colloid Interface Sci. 129, 63–77 (1989). 11. Morgan, J. J. Kinetics of reaction between O2 and Mn(ii) species in aqueous solutions. Geochim. Cosmochim. Acta 69, 35–48 (2005). 12. von Langen, P. J., Johnson, K. S., Coale, K. H. & Elrod, V. A. Oxidation kinetics of manganese(ii) in seawater at nanomolar concentrations. Geochim. Cosmochim. Acta 61, 4945–4954 (1997). 13. Madden, A. S. & Hochella, M. F. A test of geochemical reactivity as a function of mineral size: manganese oxidation promoted by hematite nanoparticles. Geochim. Cosmochim. Acta 69, 389–398 (2005). 14. Baron, F., Gaudin, A., Lorand, J. P. & Mangold, N. New constraints on early Mars weathering conditions from an experimental approach on crust simulants. J. Geophys. Res. Planets 124, 1783–1801 (2019). 15. Kite, E. S. & Daswani, M. M. Geochemistry constrains global hydrology on early Mars. Earth Planet. Sci. Lett. 524, 115718 (2019). 16. Wang, X. et al. The presence of ferrihydrite promotes abiotic formation of manganese (oxyhydr)oxides. Soil Sci. Soc. Am. J. 79, 1297–1305 (2015). 17. Lan, S. et al. Mechanisms of Mn(ii) catalytic oxidation on ferrihydrite surfaces and the formation of manganese (oxyhydr) oxides. Geochim. Cosmochim. Acta 211, 79–96 (2017). 18. Treiman, A. H. et al. Mineralogy, provenance, and diagenesis of a potassic basaltic sandstone on Mars: CheMin X‐ray diffraction of the Windjana sample (Kimberley area, Gale crater). J. Geophys. Res. Planets 121, 75–106 (2016). 19. Zhao, Y. Y. S., McLennan, S. M., Jackson, W. A. & Karunatillake, S. Photochemical controls on chlorine and bromine geochemistry at the Martian surface. Earth Planet. Sci. Lett. 497, 102–112 (2018). 20. Sutter, B. et al. Measurements of oxychlorine species on Mars. Int. J. Astrobiol. 16, 203–217 (2017). 21. Bellucci, J. et al. Halogen and Cl isotopic systematics in Martian phosphates: implications for the Cl cycle and surface halogen reservoirs on Mars. Earth Planet. Sci. Lett. 458, 192–202 (2017). 22. Mitra, K. & Catalano, J. G. Chlorate as a potential oxidant on Mars: rates and products of dissolved Fe(ii) oxidation. J. Geophys. Res. Planets 124, 2893-2916 (2019). 23. Mitra, K., Moreland, E. L. & Catalano, J. G. Capacity of chlorate to oxidize ferrous iron: implications for iron oxide formation on Mars. Minerals 10, 729 (2020). 24. Mitra, K., Moreland, E. L., Knight, A. L. & Catalano, J. G. Rates and Products of Iron Oxidation by Chlorate at Low Temperatures (0 to 25°C) and Implications for Mars Geochemistry. ACS Earth Space Chem. 6, 250–260 (2022). 25. Zheng, D., Yin, Z., Zhang, W., Tan, X. & Sun, S. Novel branched γ-MnOOH and β-MnO2 multipod nanostructures. Cryst. Growth Des. 6, 1733–1735 (2006). 26. Dreibus, G. & Wanke, H. Mars, a volatile-rich planet. Meteoritics 20, 367–381 (1985).
  7. 7. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y 27. Rampe, E. B., Cartwright, J. A., McCubbin, F. M. & Osterloo, M. M. in D.E. Harlov, L. Aranovich (eds.) The Role of Halogens in Terrestrial and Extraterrestrial Geochemical Processes 959–995 (Springer, 2018). 28. VanBommel, S., Gellert, R., Berger, J., Yen, A. & Boyd, N. Mars science laboratory alpha particle X-ray spectrometer trace elements: situational sensitivity to Co, Ni, Cu, Zn, Ga, Ge, and Br. Acta Astronaut. 165, 32–42 (2019). 29. Wang, X. et al. Multiphase volatilization of halogens at the soil– atmosphere interface on Mars. J. Geophys. Res. Planets 126, e2021JE006929 (2021). 30. Thompson, R. C. Reduction of bromine(v) by cerium(iii), manganese(ii), and neptunium(v) in aqueous sulfuric acid. J. Am. Chem. Soc. 93, 7315–7315 (1971). 31. Birk, J. P. Kinetics and mechanism of the reduction of bromate ion by hexachloroiridate(iii). Inorg. Chem. 17, 504–506 (1978). 32. Le, N. C. & Van Phuc, D. Sorption of lead(ii), cobalt(ii) and copper(ii) ions from aqueous solutions by γ-MnO2 nanostructure. Adv. Nat. Sci. Nanosci. Nanotechnol. 6, 025014 (2015). 33. Post, J. E. Manganese oxide minerals: crystal structures and economic and environmental significance. Proc. Natl Acad. Sci. USA 96, 3447–3454 (1999). 34. Noda, N. et al. Highly oxidizing aqueous environments on early Mars inferred from scavenging pattern of trace metals on manganese oxides. J. Geophys. Res. Planets 124, 1282–1295 (2019). 35. Luther, G. W. The role of one- and two-electron transfer reactions in forming thermodynamically unstable intermediates as barriers in multi-electron redox reactions. Aquat. Geochem. 16, 395–420 (2010). 36. Hem, J. D. Rates of manganese oxidation in aqueous systems. Geochim. Cosmochim. Acta 45, 1369–1374 (1981). 37. Coughlin, B. R. & Stone, A. T. Nonreversible adsorption of divalent metal ions (Mnii , Coii , Niii , Cuii , and Pbii ) onto goethite: effects of acidification, FeII addition, and picolinic acid addition. Environ. Sci. Technol. 29, 2445–2455 (1995). 38. Jacobsen, F., Holcman, J. & Sehested, K. Oxidation of manganese(ii) by ozone and reduction of manganese(iii) by hydrogen peroxide in acidic solution. Int. J. Chem. Kinet. 30, 207–214 (1998). 39. Yen, A., Kim, S., Hecht, M., Frant, M. & Murray, B. Evidence that the reactivity of the Martian soil is due to superoxide ions. Science 289, 1909–1912 (2000). 40. Zent, A. P., Ichimura, A. S., Quinn, R. C. & Harding, H. K. The formation and stability of the superoxide radical (O2 − ) on rock‐forming minerals: band gaps, hydroxylation state, and implications for Mars oxidant chemistry. J. Geophys. Res. Planets 113, E09001 (2008). 41. Learman, D. R., Voelker, B. M., Madden, A. S. & Hansel, C. M. Constraints on superoxide mediated formation of manganese oxides. Front. Microbiol. 4, 262 (2013). 42. Liu, D. & Kounaves, S. P. The role of titanium dioxide (TiO2) in the production of perchlorate (ClO4 – ) from chlorite (ClO2 – ) and chlorate (ClO3 – ) on Earth and Mars. ACS Earth Space Chem. 3, 1678–1684 (2019). 43. Rao, B., Anderson, T. A., Redder, A. & Jackson, W. A. Perchlorate formation by ozone oxidation of aqueous chlorine/oxy-chlorine species: role of ClxOy radicals. Environ. Sci. Technol. 44, 2961–2967 (2010). 44. Hamilton, G., Chiswell, B., Terry, J., Dixon, D. & Sly, L. Filtration and manganese removal. J. Water Supply Res. Technol. Aqua 62, 417–425 (2013). 45. Chen, L., Zhang, J. J. & Zheng, X. L. Coupling technique for deep removal of manganese and iron from potable water. Environ. Eng. Sci. 33, 261–269 (2016). 46. Adam, L. C., Fabian, I., Suzuki, K. & Gordon, G. Hypochlorous acid decomposition in the pH5–8 region. Inorg. Chem. 31, 3534–3541 (1992). 47. Farley, K. A. et al. Light and variable 37 Cl/35 Cl ratios in rocks from Gale crater, Mars: possible signature of perchlorate. Earth Planet. Sci. Lett. 438, 14–24 (2016). 48. Brundrett, M., Yan, W., Velazquez, M. C., Rao, B. & Jackson, W. A. Abiotic reduction of chlorate by Fe(ii) minerals: implications for occurrence and transformation of oxy-chlorine species on Earth and Mars. ACS Earth Space Chem. 3, 700–710 (2019). 49. Miller, L. G., Baesman, S. M., Carlström, C. I., Coates, J. D. & Oremland, R. S. Methane oxidation linked to chlorite dismutation. Front. Microbiol. 5, 275 (2014). 50. Zhang, X. et al. Reaction of methane and UV-activated perchlorate: relevance to heterogeneous loss of methane in the atmosphere of Mars. Icarus 376, 114832 (2022). 51. Smith, M. L., Claire, M. W., Catling, D. C. & Zahnle, K. J. The formation of sulfate, nitrate and perchlorate salts in the Martian atmosphere. Icarus 231, 51–64 (2014). 52. Catling, D. C. et al. Atmospheric origins of perchlorate on Mars and in the Atacama. J. Geophys. Res. Planets 115, E00E11 (2010). 53. Turner, A. M., Abplanalp, M. J. & Kaiser, R. I. Mechanistic studies on the radiolytic decomposition of perchlorates on the Martian surface. Astrophys. J. 820, 127 (2016). 54. Wu, Z. C. et al. Forming perchlorates on Mars through plasma chemistry during dust events. Earth Planet. Sci. Lett. 504, 94–105 (2018). 55. Carrier, B. L. & Kounaves, S. P. The origins of perchlorate in the Martian soil. Geophys. Res. Lett. 42, 3739–3745 (2015). 56. Melwani Daswani, M. & Kite, E. Paleohydrology on Mars constrained by mass balance and mineralogy of pre‐Amazonian sodium chloride lakes. J. Geophys. Res. Planets 122, 1802–1823 (2017). 57. Kim, Y. S., Wo, K. P., Maity, S., Atreya, S. K. & Kaiser, R. I. Radiation-induced formation of chlorine oxides and their potential role in the origin of Martian perchlorates. J. Am. Chem. Soc. 135, 4910–4913 (2013). 58. Kang, N., Anderson, T. A., Rao, B. & Jackson, W. A. Characteristics of perchlorate formation via photodissociation of aqueous chlorite. Environ. Chem. 6, 53–59 (2009). Publisher’s note Springer Nature remains neutral with regard to jurisdictional claims in published maps and institutional affiliations. Springer Nature or its licensor (e.g. a society or other partner) holds exclusive rights to this article under a publishing agreement with the author(s) or other rightsholder(s); author self-archiving of the accepted manuscript version of this article is solely governed by the terms of such publishing agreement and applicable law. © The Author(s), under exclusive licence to Springer Nature Limited 2022
  8. 8. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Methods Geochemical modelling of Mn(ii) oxidation by O2 TherateofoxidationofdissolvedMn(ii)bymolecularoxygenat25 °C wasimplementedinTheGeochemist’sWorkbenchmoduleReact59 using amodifiedLawrenceLivermoreNationalLaboratorydatabase60,61 .The hydrolysisandcomplexationconstantsofMnwereadjustedtomatch thosereportedbyMorganetal.11 .ThegassolubilityofCO2 andtheacid dissociationconstantsinthedatabasewereverifiedtobeequaltothose reportedbyref.11 .Weimplementedtheratelawfromref.11 −d [Mn (II)] /dt =4 [O2] { k1 [MnOH + ] + k2 [Mn (OH)2 (aq)] + k3 [Mn (CO3) 2− 2 ] (1) wherek1,k2,andk3 aretherateconstantsforeachkinetictermwithlog k1 = −1.78,logk2 = +1.32andlogk3 = −1.09. Theoxidationof1 mmol l−1 Mn(ii)intherangeofexpectedO2 con- centrations(0.03,0.01and10−5 bar)(ref.6 )wasmodelledinabackground fluidof100 mmol l−1 MgCl2 thatisconstantlybufferedbyO2 and0.5 bar CO2 asafunctionofpH(7–4).TwotypesofpHsystemswerealsostudied, fixed pH systems and free-drift pH systems. The fixed pH systems are idealscenariosinwhichthereisperfectpHbufferingandthepHnever undergoesanychangeduringreaction.Thefree-driftpHsystemsdisplay pHbufferingonlyfromthedissolvedspeciespresent.Allnaturalsystems fallinbetweenthesetwoendmemberscenarios.Purelyaqueoussystems behaveasfree-driftsystemswhilereactionsinfracturesorporespaces behaveasanintermediatebetweenfree-driftandpHbufferedsystems owingtothepHbufferingarisingfromwater–rockinteraction. Rhodochrosite(MnCO3)wassuppressedinthesecalculationsbecause itsformationabovepH5.56lowersthedissolvedMn(ii)concentrationin ourmodelandthusslowstherateofoxidation62 .Simulationsabovethis pH value should thus be considered lower bounds for the timescale of Mn(ii)oxidationbyO2.NotethatthesolublerangeofMn(ii)extendsto pH6ifdissolvedMn(ii)concentrationsarelimitedto~150μmoll−1 andto pH7iflimitedto~4μmoll−1 .Inaddition,bixbyite(Mn2O3)wassuppressed becauseitisnotreportedinstudiesofMnoxidation11 . Mn(ii)oxidationbyoxyhalogenspecies Laboratoryexperimentswereconductedtoinvestigatewhetherchlo- rate and/or bromate are capable of oxidizing Mn(ii) in Mars-relevant fluids at ambient conditions (24 ± 1 °C, 1 atm). To isolate the effects of chlorate or bromate on dissolved Mn(ii), all reactions were con- ducted inside a Coy Laboratory Products vinyl anaerobic chamber (N2 = 97%, H2 = 3%), with <1 ppmv O2 concentration, maintained using palladium catalysts, to inhibit O2 interference in Mn(ii) oxidation. All reactors were wrapped with Al foil to inhibit any probable Mn photo-oxidation. Stock solutions of manganese(ii) chloride tetrahy- drate[MnCl2·4H2O],manganese(ii)sulfatemonohydrate[MnSO4·H2O], magnesium chloride hexahydrate [MgCl2·6H2O], magnesium sulfate hexahydrate[MgSO4·6H2O],sodiumchlorate[NaClO3]andsodiumbro- mate[NaBrO3]werepreparedusingAmericanChemicalSociety-grade FisherScientificreagentsindeionized,deoxygenatedwaterandkept insidetheanaerobicchamberindarkreagentbottles.Thestocksolu- tionsweredilutedinappropriateamountstopreparetheexperimental solutionsandachievethedesiredstoichiometry. Twotypesofexperimentswereconductedforstudyingtheeffect ofchlorateand/orbromateondissolvedMn(ii):(1)kineticexperiments and(2)batchexperiments.Whilethekineticexperimentsweresetup to track the concentration of Mn(ii) and pH as a function of time, the batchexperimentsweresetuptostudythemineralproductsofMn(ii) oxidation by chlorate and/or bromate. The batch experiments were greater volume (150 ml) replicates of the kinetic experiments (50 ml) to extract a substantial amount of mineral product for identification and characterization. Reactors of batch experiments were sealed in air-tightserumbottlesusingbluebutylrubberstoppersandbrought outside the anaerobic chamber to be placed on a shaker table set to ~170 rotations per minute to promote mixing and avoid gravimetric settling of any mineral produced. The kinetic experiment reactors were kept inside the anaerobic chamber on end-over-end rotators. The kinetic experiments were sampled by extracting a small aliquot fromthesereactorstodeterminethe[Mn(ii)]andpHinregularweekly timeintervals.Thesolutionsandthemineralprecipitatesofthebatch reactorswereanalysedonlyattheendoftheexperiments. Theexperimentalsolutions(bothkineticandbatch)weresetupwith threetypesofoxidants:(A)onlybromate,(B)onlychlorateand(C)both bromate and chlorate. All experiment types contained Mars-relevant backgroundsaltsmagnesiumchlorideormagnesiumsulfate63,64 toserve asionicstrengthbuffers(~100 mmol l−1 ).Thebackgroundsaltsalsopro- videanionswithdifferentabilitiestocomplexdissolvedMn(ii),thereby having the capacity to affect reaction rates and mineral products. The initialpHoftheexperimentalsolutionswassetusing1 mol l−1 solutions ofhydrogenchlorideandsodiumhydroxideto7,5and3usingaThermo ScientificOrionStarpHmeterwithanAgClelectrode.Theexperimental pH was allowed to drift freely as a response to Mn(ii) oxidation and to determine the effect of Mn(ii) oxidation on the acidity of the solution. Analogousoxidant-freecontrolexperimentswerealsopreparedatpH 7 and 3 to verify the absence of inadvertent oxidation by stray oxidants (forexample,O2)inthereactors(SupplementaryFig.5). Oxidation experiments using only bromate sought to determine whetherthisoxyhalogenspeciescanoxidizeMn(ii)andproduceMn(iii/ iv)mineralsinMars-relevantfluids.Solutionscontaining~100 mmol l−1 (or10mmoll−1 )Mn(ii),aseitherMnCl2 orMnSO4,werepreparedinback- groundsaltmixturesof0.1moll−1 Mg-chlorideorMg-sulfate,respectively. Stocksolutionofsodiumbromate(2moll−1 )wasaddedtotheexperimen- talsolutionstobringthebromateconcentrationapproximatelyequalto Mn(ii)(~100or10 mmol l−1 ),asperrequirement. Similarly, oxidation experiments using only chlorate sought to determine whether this oxyhalogen species can oxidize Mn(ii) and produce Mn(iii/iv) minerals in Mars-relevant fluids. Oxidant type A. Solutions containing ~100 mmol l−1 Mn(II), as either MnCl2 or MnSO4, werepreparedinbackgroundsaltmixturesof0.1 mol l−1 Mg-chlorideor Mg-sulfate,respectively.Stocksolutionofsodiumchlorate(2 mol l−1 ) wasaddedtotheexperimentalsolutionstobringthechlorateconcen- trationapproximatelyequaltoMn(ii)(~100 mmol l−1 ). A third set of experiments evaluated Mn(ii) oxidation in mixed chlorate/bromatesystems.Experimentalsolutionswith100 mmol l−1 chlorateand10 mmol l−1 bromatewasmixedwith100 mmol l−1 Mn(ii) in both Mg-chloride and Mg-sulfate fluids starting at initial pH 7, 5 and 3. All other experimental conditions remain identical except the inclusion of both chlorate and bromate together to replicate Martian fluidconditions.Theconcentrationofbromatecouldrangeashighas 10%ofthechlorateinMartianfluids27,65,66 ,andthereforeexperimental solutionsinvestigatedMn(ii)oxidationwith100 mmol l−1 chlorateand 10 mmol l−1 bromate. Control experiments investigated Mn(ii) oxida- tionin100 mMchlorateand10 mMbromateseparatelyforcomparison. Analyticalmethods Thebromateconcentrationsintheexperimentalsolutionsweremeas- ured at the start of the reaction and at the end of the experiment with Thermo Fisher Scientific Dionex Integrion high-pressure ion chro- matograph using an IonPac AS11 analytical column equipped with an AG11 guard column, ADRS 600 suppressor, a conductivity detector and an AS-DV automated sampler. Analytical conditions consisted of 12 mMKOHeluentataflowrateof1 ml min–1 .Bromateconcentrations atthestartoftheexperimentweredeterminedbeforeadditionofdis- solved Mn(ii) and were corrected for the dilution resulting from this final solution component. The Mn(ii) concentration was measured periodicallyusingiCAP7400Duoinductivelycoupledplasmaoptical emission spectroscopy. The first data point was taken immediately afterthestartoftheexperiments,anditrepresentstheinitialamount ofdissolvedMn(ii)presentinthesolutions.
  9. 9. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y The minerals precipitated in the reactors were collected at the end of the experiment by filtration using a 0.22 µm pore size mixed cellulose ester membrane. The filtered solid products were dried in a vacuum desiccator for 3–4 days before characterization using XRD, XPS and VNIR spectroscopy. XRD was performed using a Bruker d8 Advance diffractometer configured with a Cu source operating at 40 kVand40 mAandaposition-sensitive,energy-dispersiveLynxEye XEdetector.Thedriedmineralsamplesweregroundusingamortarand pestle before analysis, and a zero-background silicon sample holder was used because the amount of mineral precipitated in these batch reactionswaslimitedinsomesystems.TheXRDscanswereperformed usingcontinuousscanningwitha0.05°stepsizeand1 scounttimeper step. The Bruker Diffrac.Eva programme was used for processing the rawdataandidentificationofmineral(s)basedonpatternmatching. TheXPSanalyseswereconductedusingaPhysicalElectronics5000 Versa Probe II Scanning ESCA microprobe equipped with a monochro- matic Al Kα X-ray source. The binding energy was calibrated using the C1s248.8eVadventitiouscarbonpeak.TheMn3ppeakhasbeenshownto providerobustquantificationofMnoxidationstate67 ;thusthispeakwas usedtoquantifytheratioofMnoxidationstatesinsamples.Peakswere fittedinCasaXPSsoftwarewithanasymmetricLorentzian(LA)function andaniteratedShirleybackground.Otherpeakshapefunctions,suchas SumGaussian–Lorentzianwerealsoexplored,butthesefunctionsyielded lower quality fits with larger errors. The proportion of Mn(ii), Mn(iii) andMn(iv)ineachsamplewasdeterminedbyfittingacomponentpeak for each of these oxidation states and determining its relative area, an approachthathasbeenemployedtoquantifyMnoxidationstatesinfiltra- tionmediasamplesfromwatertreatmentplantsandMn(hydr)oxides67,68 . Eachofthecomponentpeakswasfixedattheprimarypeakposition(>80% of total peak area) identified for Mn(ii), Mn(iii) and Mn(iv) standards. Thesethreestandards,manganese(ii)chloride[MnCl2],manganese(iii) oxyhydroxide[MnOOH]andmanganese(iv)oxide[MnO2],wereselected tobestrepresentpossibleformationproductsfromoxidationbychlorate. MonteCarloanalysiswasusedtoassessthestabilityofpeakfitsandobtain associatederrorsforeachcomponent.AVoigtlineshapewasusedforall fits(LA(1.53,243)).SincesamplepeakscouldbefitwithonlyMn(iv)and Mn(iii)components,thestandarddeviationwasthesameforbothofthese components;thatis,anincreaseintheproportionoffittedMn(iii)by4% wouldresultina4%decreaseintheproportionoffittedMn(iv). TheVNIRreflectancespectraofthemineralprecipitatesproduced were collected using an Analytical Spectral Devices (ASD) portable VNIR spectrometer. The powdered samples of the mineral precipi- tates were measured using an ASD Muglight, a bench-top probe from 0.35–2.50 μm spectral range with a fixed illumination angle of 30°. The spectra were normalized using a Spectralon diffuse reflectance standardtoradiancecoefficients. Dataavailability The data associated with the manuscript are available at: https://fig- share.com/s/e12d62da416302225cf3. Codeavailability The code to model Mn(ii) oxidation by O2 in the Geochemist’s Work- bench is available at https://figshare.com/articles/online_resource/ MnII_oxidation_species_Morgan_V7_rea/21066232. References 59. Bethke, C. M. Geochemical and biogeochemical reaction modeling. (Cambridge University Press, 2007). 60. Catalano, J. G. Thermodynamic and mass balance constraints on iron-bearing phyllosilicate formation and alteration pathways on early Mars. J. Geophys. Res. Planets 118, 2124–2136 (2013). 61. Delany, J. & Lundeen, S. R. The LLNL Thermochemical Database Report UCRL-21658 (Lawrence Livermore National Laboratory, 1990). 62. Sternbeck, J. Kinetics of rhodochrosite crystal growth at 25°C: the role of surface speciation. Geochim. Cosmochim. Acta 61, 785–793 (1997). 63. Vaniman, D. T. et al. Magnesium sulphate salts and the history of water on Mars. Nature 431, 663–665 (2004). 64. Fox-Powell, M. G., Hallsworth, J. E., Cousins, C. R. & Cockell, C. S. Ionic strength is a barrier to the habitability of Mars. Astrobiology 16, 427–442 (2016). 65. Gellert, R. et al. Alpha Particle X‐ray Spectrometer (APXS): results from Gusev crater and calibration report. J. Geophys. Res. Planets 111, E02S05 (2006). 66. Marion, G., Catling, D. & Kargel, J. Br/Cl partitioning in chloride minerals in the Burns formation on Mars. Icarus 200, 436–445 (2009). 67. Ilton, E. S., Post, J. E., Heaney, P. J., Ling, F. T. & Kerisit, S. N. XPS determination of Mn oxidation states in Mn (hydr) oxides. Appl. Surf. Sci. 366, 475–485 (2016). 68. Cerrato, J. M., Hochella Jr, M. F., Knocke, W. R., Dietrich, A. M. & Cromer, T. F. Use of XPS to identify the oxidation state of Mn in solid surfaces of filtration media oxide samples from drinking water treatment plants. Environ. Sci. Technol. 44, 5881–5886 (2010). Acknowledgements This research was funded by NASA Science Mission Directorate Future Investigators in NASA Earth and Space Science and Technology (FINESST) programme through award no. 80NSSC19K1521. J.G.C. was supported by the NASA Exobiology programme through award no. 80NSSC18K1292. G.J.L. was supported by the National Science Foundation Graduate Research Fellowship Program under grant no. DGE- 1745038 and DGE-2139839. Discussions with B. Jolliff, R. Arvidson and J. Hurowitz improved this manuscript. P. Carpenter is thanked for assistance with XRD data collection and Rietveld refinements. R. Arvidson and A. Knight are thanked for assistance in VNIR data collection. Authorcontributions K.M. and J.G.C. designed the study. K.M. conducted the experiments with assistance from E.L.M. G.L. contributed X-ray photoelectron spectroscopy analyses. J.G.C. and K.M. performed the thermodynamic and kinetic modeling. K.M. analyzed the results and wrote the original manuscript, with additional text provided by J.G. and further editing by G.L. and E.L.M. All authors have read and agreed to the published version of the manuscript. Competinginterests The authors declare no competing interests. Additionalinformation Extended data is available for this paper at https://doi.org/10.1038/s41561-022-01094-y. Supplementary information The online version contains supplementary material available at https://doi.org/10.1038/s41561-022-01094-y. Correspondence and requests for materialsshould be addressed to Jeffrey G. Catalano. Peer review information Nature Geoscience thanks Yasuhito Sekine and the other, anonymous, reviewer(s) for their contribution to the peer review of this work. Primary Handling Editor: Tamara Goldin, in collaboration with the Nature Geoscience team. Reprints and permissions informationis available at www.nature.com/reprints.
  10. 10. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y ExtendedDataFig.1|ReactionofMartianbasaltwithwaterequilibratedwith0.5 barCO2 and0.03 barO2.(a)RelationshipofdissolvedCO2 andO2 concentrationstopH.(b)Massfractionofmineralalterationproductions.
  11. 11. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y ExtendedDataFig.2|Mineralogyofthesolidsproducedbyoxidationof dissolvedMn(II)by10 mmol L−1 bromate.XRDpatternsofthesolidsproduced byreactionof10 mmol L−1 Mn(II)with10 mmol L−1 bromatein100 mmol L−1 magnesiumchloridefluids.Patternsarevisuallyoffsetforclarity.Diagnostic peaksarelabeledandindicatedwithdashedlines.SeeExtendedDataTable4for completesampledetails.
  12. 12. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y ExtendedDataFig.3|ImagesofthesolidsproducedbyoxidationofdissolvedMn(II)bybromate.Mineralsprecipitatedfollowingreactioninmagnesiumchloride (a-d)andmagnesiumsulfate(eandf)fluids.
  13. 13. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y ExtendedDataFig.4|ComparisonofVNIRspectraofthemanganeseoxidemineraldetectedatEndeavorcrater,Mars,andproducedinMn(II)oxidation experiments.SolidsproducedfromoxidationofdissolvedMn(II)bybromatein(a)magnesiumchlorideand(b)magnesiumsulfatefluids.Thelabelsindicatethe initialpHofthesamples.SeeExtendedDataTable5forcompletesampledetails.
  14. 14. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 1 | Homogeneous and heterogeneous Mn(II) oxidation rates by oxygen
  15. 15. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 2 | Fluid composition of the kinetic experiments shown in Fig. S1 with [Mn(II)]≈[ClO3 − ] ≈ 100mmolL−1
  16. 16. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 3 | Fluid composition of the kinetic experiments shown in Fig. 2 with [Mn(II)]≈[BrO3 − ] ≈ 100mmolL−1
  17. 17. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 4 | Fluid composition of the mineral precipitation experiments shown in Extended Data Fig. 2 with approximately 100mmolL−1 [Mn(II)], and 10mmolL−1 [BrO3 − ]
  18. 18. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 5 | Fluid composition of the mineral precipitation experiments shown in Fig. 3 with [Mn(II)]≈[BrO3 − ] ≈ 100mmolL−1
  19. 19. Nature Geoscience Article https://doi.org/10.1038/s41561-022-01094-y Extended Data Table 6 | Percentage of Mn(II) oxidation in open systems buffered by 0.21bar O2 in presence of 4 ×10−4 bar CO2 under identical conditions as the experiments with oxyhalogens

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