.- 001CES-391 Self-Sterilizing Properties of Martian Soil: Possible Nature & Implications

A. 1. Tsapin, M. G. Goldfeld, K. H. Nealson Jet Propulsion Laboratory, California Instituteof Technology, MS 183-301, 4800 Oak Grove Drive, Pasadena, California 91 009 K. M. Kemner Argonne National Laboratory, Argonne, Illinois B. Moskowitz Institute of Terrestrial Magnetism, University of Minnesota, Minneapolis, Minnesota

Copyright 0 2000 Society of Automotive Engineers, Inc.

ABSTRACT unaffected.hours, whilelargely remained0, formation These observations were interpreted as an-indication of the As a result of the Viking missions in 1970s, the presence of presence of a strong oxidant on the Martiansurface, or, a strong Oxidant in Martian’Oil was suggested‘ Here we most probably, several different types of oxidants [3-5]. No present a testable, by near-term missions, hypothesis that site on Earth has been foundwith such an oxidantpresent, contributesto that OxidizingPool. strong enough to decomposewater. This is not unexpected, were studied for their spectral and oxidative properties and when one takes intoaccount the abundance of reductants biological activities. Ferrate(V1) has distinctive spectroscopic that make unlikelythe accumulationof such a strong oxidant features making it available for detection by remote sensing anywhereon the surfaceof the Earth. reflectance spectra andcontact measurements via .-

Commonoxidation states of are (+2) and (+3).How- some dioxygen and ozone. In these circumstances, ever, at certain conditions, higher oxidation states can be ferrate(V1) can serve as a form of stabilization and storage formed, including Fe(lV), Fe(V), andFe(VI). These are quite of activeoxygen, but only if its decomposition is slow unstable under most usual conditions on Earth, where water enoughtopermit the accumulation of the product in is abundant, but they can form and persist in dry systems. question. Besides, Fe(VI) known as tetrahedral Fe0; anion, although a verystrong oxidant, is a well-characterizedchemical FERRATE(V1) STABILITY AND PRESERVATION species, with its standardelectrochemical potentials E" = +2.2 V in acidic, +0.78 V in alkaline solutions[14]. It is rather In aqueoussolutions, ferrate(V1) is ratherstable in alkali stable in strong alkaline solutions in the absence of efficient only, at pH>10, and even then it is readily reduced by most reductants. Ferrates(V1) salts with various cations, such as organic materials. In a pH-E" diagram, where E" is reduction K', Na', Ba", Li', Rb', Cs', Ag', and a few tetralkyVaryl potential,thestability region can be approximately ammonium cations, have been described [15-181. Currently, presented as in Fig. 1. Beyondthat region, ferrate(V1) is there is a burst of interest to ferrate (VI) as a promising reduced according to the following equilibria: oxidizing reagent for organic synthesis [16,19] and material

for rechargeablealkaline batteries of increasedcapacity Fe0:- + 8H+ + 3e- w Fe3++ 4H,O POI. Fe0; + 4H,O + 3e- w Fe(OH),(s)+ 50H' In order to assess the possible contribution of ferrate(V1) to the oxidizingpool in the Martiansoil, we address the following questions: 1. Is therean opportunity to form ferrate (VI) onMars, accordingto what is knownabout its surface compo- sition and environmental conditions? T P~oH~' 2. Would such a compound be stableenough to persist and accumulate under those conditions? 3. Would it display the essential reactivity that was found in the samples of Martiansoil in the Viking experiments, i.e. produce gas while moisturized, and 4. Would it produce carbon dioxide when in contact with the organic materials that constituted nutrient solution in those experiments? 5. Would these chemical activities be impaired at heating in a waysimilar to the inactivation seen in Viking's 2 4 6 8 10 12 14 experiments? PH 6. With Marsexploration program in mind, what are the possible approaches to identification of ferrate(Vl), and Fig. 1. Eo-pH stability diagram for iron compoundsin in particularwhat are the spectralfeatures of this aquatic systems. species thatwould permit its characterization by both contact measurements and remote sensing? At a lower pH, not only organic materials, but also water is oxidized. So, thermodynamically one wouldneed a highly HIGHER FERRATES: CHEMISTRYy BIOLOGICAL alkalineenvironment for ferrate(V1) to be stored. This ACTIVITY, AND SPECTRAL FEATURES requirement is not so limiting for Mars soil, as it seems at first glance. Indeed, there are strong reasons to assume that ROUTES TO FERRATE(V1) in the absence of biogeniccalcium carbonatedeposits The regular way to ferrate(V1)is through the wet oxidation of which now serve as a powerful buffer on the Earth surface, Fe(lll) with hypochlorite [15, 161: we used this approach in bothpre-Cambrian Earth, and Mars at some stage of its ourexperiments. However, formation of ferrate(V1)was geological history, were covered with a "soda ocean" of its reported in dry, elevated-temperaturereactions of iron- pH>9 [23, 241. containingmaterials with somealkaline peroxides and Different scenarios, with ratheracidic evolution pathways superoxides [21, 221. On the other hand, there is a strong havebeen suggested forMars [25]. However, in the beliefthat active oxygen speciessuch as peroxides, absence of conclusive pH measurements, the only experi- superoxides,singlet oxygen, hydroxyl radicals, and ozone mentalevidence now available, from the same Viking UV are formed under irradiation in Martian atmosphere, and program, as Oyamapointed out [l], favorsan alkaline affect its soil [6-111. Overall, Martian atmosphere is highly environment:indeed, the solutionsproduced by mixing oxidized, dominated by carbon dioxide and with presence of samples of Martiansoil with water displayed a short-term I.

absorption of carbon dioxide, as any basic solution would do. From simulation experiments, Quinn and Orenberg [26] also concluded that Martian soilmaterial is mostlikely at -16.177% least mildly alkaline.

Ferrates of alkaline metals are unstable in the presence of 1 moisture.However the Martian surface is dry andcold. -8.75646 Thus, evenalkaline ferrates couldbe stabilized there. Besides, otherferrates, such as ferrate(Vl), are much less water soluble and consequently fairly stable in 295.81C humid milieu. It is essential that ferrate(V1) anions are not sensitive to light [27]. Thus, overall, there is enough reason to suggest that the formation of ferrate (VI) and its preservation in soil are consistent with the present knowledge of chemical compo- sition and environmental conditionsat the Martian surface. CHEMICAL PROPERTIES OF FERRATE(V1) In the Viking Gas Exchange Experiments, it was found that samples of Martiansoil released oxygen gas upon introduction of watervapor into the sample cell [1,2]. Oxygen was also released from a sample of soil that had been pre-heated at 145"C, although the amount of gas was substantiallyreduced. When treated with a "nutrient" solution, i.e. an aqueous mixture of organic acids and amino acids,carbon dioxide was first rapidly released, thenthe rate of its production decreased after a small fraction of all organic carbon had been oxidized. Carbon dioxide release Fig. 2. Thermal decomposition of &FeO,: was completely prevented by preheating the soil samples at Thermogravimetry (TGA), and differential scanning 160°C, and a consumption was observed instead. In light of calorimetry (DSC). Heating rate: 5"C/min. these results,thermal decomposition of ferrate(V1) is relevant to the problem in hand. Oxvaen Release in the Reaction with Water Thermal Decomposition Dry ferrate is not immediately reactive with aprotic solvents Potassiumferrate(V1) as a dry powder is stable at room such as ether, chloroform, or benzene, which permits their temperature. However, on heating it decomposes, releasing applicationfor removal of traces of waterfrom ferrate oxygen gas. Decomposition slowly proceeds, starting about preparations. Neither is ferrate(V1) soluble in any of those 50°C, and is complete by 300°C. Its rate strongly depends solvents. on the traces of water: dry samples are less sensitive to heating.Thermal decomposition is a complex,multi-stage Upon addition of water to (VI), the mixture process, as can be seen fromthermogram and DSC bubblesindicating an intense gas evolution. In Viking's observations (Fig. 2). One may speculate that transient experiments, no attemptwas made exploreto the lower oxidation states of iron, such as Fe(+5) and Fe(+4), interaction of soil samples with pure liquid water.Oxygen are first produced,before the final product, Fe,O,, forms. evolution was recorded, however, when water vapor from These intermediate forms still display high oxidative power. the nutrient solution accessed the solid soil. The nature of A catalyticeffect of Fe(lll) on ferratedecomposition is thereaction was not explored further, and it remained anticipated and may significantly complicate the results of unclear if it were a stoichiometricoxidation of water,or DSC andthermogravimetry analysis. These results are decomposition of the solid material, catalyzed or initiated by essential to the understanding of Viking's data since they water. show that "sterilization" might result in a set of various iron products in a poorly predictable manner, depending on the The reaction of water with potassium ferrate, however, is a other components of the soil and possiblyon the rate of truewater oxidation: GC-MS analysis of the gas product heating as well. after reaction of potassium ferrate with liquid H,'*O revealed that the gas product formed in the head space of the reactor vessel is "O,, indicatingno significant oxygen exchange between ferrate and water. Water oxidation is enhanced by acids, and the true active oxidizing species is HFe04- anion, 0,Accumulation not Fe0,2' dianion per se [24]. Thereaction may be A 1400 presented as: I200 1000 HFe04-+ % H,O + FeOOH + 34 0, + OH' 800 1 800 400 Barium ferrate(Vl), BaFeO;H,O,formed by ion-exchange 200 precipitation from the sodium ferrate(V1) solutionwith barium 0 -m nitrate, was not reactive under comparable conditions, along Time, houn with its low solubility.

We explored the kinetics of dioxygen accumulation in two B O2Accumulation ways. First, as in the Viking experiments, oxygen formation wasmeasured in the setting wherein solidpotassium 2m ferrate(V1)contacted with watervapor only, while liquid 2000 water and ferrate were placed in separate open containers 115cQ in one and the same chamber of the gas-meter instrument 1000 (Micro-Oxymax respirometer, with paramagnetic sensor for 5cQ oxygen).Typical results onthe kinetics of dioxygen accumulation under these conditions are in the Fig. 3A. The reactiondevelops slowly and steadily, and the kinetics O1Accumulation obviously reflects the transferof vapor to ferrate(V1). C

In another set of experiments, potassium ferrate(Vl), in the 500 amount of ca. 50 mg, wasplaced in the respirometer 400 chamber; then 0.2 mL of water was injected providing direct 300 contact of the solid material with liquid water. In this setting, 200 oxygenwas released immediately (Fig. 34. Thereaction 0 wasnot significantly impaired by pre-heating the solid -100 potassium ferrate at 145"C, or even at 170°C for 2 hrs. In Time, hours that sense, theobservations were close, at leastat a qualitative level, to the Viking results, which also indicated that oxygen formation was rather resistant to pre-heating of Fig. 3. Accumulation of oxygen at the reaction of potassium the soilsample. When waterwas replaced by a liquid ferrate (50 mg) with water. A: Reaction with water vapor. 1 : Non-heated sample; 2: Sample pre-heatedat 170°C for "nutrient solution", a short burst of oxygen was typically first 2 hrs; 3: Sample pre- heated at185°C for 4 hrs. observed, although in a lesser amount than with pure water, B: Reaction with liquid water, 2 mL, injected through a but then oxygen consumption was recorded (Fig. 3C). More septum. 1 : Non-heated sample; 2: Solid pre-heated at thorough heating, at 185°C for 4 hrs, resulted in ferrate(V1) 145°C for 2 hrs; 3: Solid pre-heated at 170°C for 2 hrs. decomposition (color of the solid drastically changed from C: Reaction with 2 mL of nutrient solution (0.024 M in dark purple to greenish), andproduced solid material lactate, formate, alanine, glycine, glycolate,pH6). 1 : No pre- incapable of releasingoxygen (Fig. 3A, Graph 3 and Fig. heating; 2: Solid pre-heated at 145°C for 2 hrs; 3: Solid pre- 3C, Graph 3). Needless to say, a mixture of iron(ll1) oxide heated at 185°C for 4 hrs. with alkali did not produce any oxygen, and only fluctuations within instrument stability limits of approx. 550 pLof gas In our hands, addition of aqueous solutions of formate to volume could be seen. powderedpotassium ferrate(V1) resulted in the release of Carbon Dioxide Release/ConsumRtion in the Reaction with carbon dioxide in the head space of the reactor vessel (Fig. Nutrient Solutions 4, Graph 7). Formate was chosen for these preliminary experimentsbecause carbon dioxide is the onlypossible According to their respective redox potentials, ferrate(V1) is product of its oxidation. Therefore, even visual observations a stronger oxidant than permanganate and chromate. The of the change in color (from purple to yellow-brown) serves extremelystrong oxidizing power of ferrate(V1)was well an unambiguous indication of carbon dioxide formation. In documented in its reactions with a number of various further experiments, CO, was determined using IR sensor of compounds: it converts chromium(1ll) to chromate, oxidizes the Micro-Oxymax respirometer. When a complete nutrient ammonia,cyanide, hydrogen sulfide andother sulfur solution was added to solid ferrate(Vl), release of carbon compounds [29-341, andwas proposed as a valuable dioxidewas recorded as well Fig. 4, Graphs 2,3,8). oxidizingreagent for organic synthesis [16, 351 and However, it was completely inhibited and actually reversed wastewater treatment [36]. to a steady CO, consumption after pre-heating at 145°C or 170°C (Fig. 4, Graphs 4-7). Viking’s settings. Nevertheless, the overall set of results, i.e. COz AccumlationlConsumption evolution of oxygen on water vapor-ferrate contact and its

1500 I I relative lack of sensitivity to pre-heating, as well as carbon dioxide release and absorption,together with pre-heating 1000 effects, are in line with Viking’s observations.

500 BIOLOGICAL ACTIVITY OF FERRATE(V1) 4 1 0 Antisepticproperties of ferrate(V1) have beenpreviously, reported,and sodium ferrate has been suggested as a -500 replacementforchlorine and ozone in waste-water -1000 J I treatment [35]. Experimentsconducted on Viking landers Tlme, hours failed to detect any organic material on the Martian surface. It has been concluded therefore that oxidants in Martian soil Fig. 4. Carbon dioxide release/consumption at the reaction were responsiblefor the destruction of organicorganics. of potassium ferrate (50 mg) with 0.2 mL 2.4 M formic acid, Havingthat in mind, we decidedto study in some more [Graph 71; or pH 8 [Graphs 3,4, s]. 7-3 no or 2 mL nutrient detail the interaction between ferrate(V1)and a few most solutions, pH 6; [Graphs2, 5, 4, pre-heating; 4, 5 pre- important biogenic compounds. We applied two approaches heated at 145°C 2 hrs, at 170°C;6, 7: pre-heated 3 hrs. in order to elucidate the efficiency and mechanism of the Graph 8: mixture of 50 mg potassium ferrate with 150 mg biologicaleffects of ferrate(V1): the alterations of the UV silicon dioxide, nutrient solution pH 8. absorption spectra of solutions containing aromatic amino acids, nitrogen bases, andnucleotides were measured in the presence of potassiumferrate(Vl), and polymerase At the same time, some dioxygen was formed in all chainreaction (PCR) wasconducted in the nucleicacid experiments with diluted nutrient solutions, apparently due solutions beforeand after application of ferrate(V1). to a concomitant reaction with water, which is certainly a far more abundant reagent. It should be kept in mind that some AbsorDtion sDectra backgroundcarbon dioxide consumption is anticipated in each case due to the alkaline nature of the solid reagent. It wasfound that contact with ferrate(V1)invariably The balance between oxidation and neutralization reactions eliminated characteristic 280 nm absorption peak of trypto- may give variousnet concentrations of carbondioxide in phane and phenylalanine. Likewise, the peak at 260 nm in head space, and in the absence of oxidation (as after the solutions of the four nucleotides, A, T, G and C, was thermal treatmentof ferrate(Vl)), steady consumptionof CO, drastically decreased in the presence of ferrate(V1). is actuallyobserved. The product of ultimate thermal decomposition of potassium ferrate displays essentially the PCR Data same behavior in terms of CO, consumption, as a mixture of iron(ll1) oxide and KOH. Ferrate(V1) reacts not onlywith isolated nucleotides, but with theentire cDNAmolecules isolated from Shewanella The above observations are in line with the results of the putrifaciense bacteria. We found also that plasmids Viking LabeledRelease Experiments. However a wordof containing inserted genes of 16sRNA from different sources caution may be worthwhile concerning these similarities. We couldnot be used as templates forpolymerase chain do not actually know what the other components of Martian reaction (PCR) aftercontacting ferrate(V1). In that sense, soil are, or the way iron(Vl), if present, is chemically bound sterilization with ferrate(V1) is essentially different from other to other oxides, and how such a combination would affect agents: usually,enzymes catalyzing template synthesis of the reactions of ferrate. For instance, we do know already DNA are damaged, while the genetic material itself remains that a combination with (barium ferrate(V1)) is unaltered and can be used again for synthesis once enzyme far less reactivethan alkaline ferrates.Besides, ourlab inhibitor is removedor destroyed. Genetic material in the experimentswere designed in such a way as tomake form of bacterial DNA, was totally destroyed by ferrate(Vl), measurements convenient and distinctively visible. To this so thatno further polymerization of complementary chain end, e.g., concentrations of “nutrient solutions” were taken were possible. Although the genetic material itself remains hundred times those in standard microbiological media. As a unalteredand can be used again for template synthesis result, all kinetics were completed within several hours. No once the enzymeinhibitor is removedor destroyed. With attempt hasbeen made to stretchout observations for ferrate(VI),the situation is different: although ferrate(V1) is months, as was the case in the real Viking’s missions. easily destroyed and removed from the solution as iron(ll1) Experiments with more diluted solutions and different ratios oxide, no further DNA synthesis occurred until the solution of solid ferrate/solution volume, which gave inconclusive wasinfected by outer sources. For that reason, it seems results andapproached theinstability limits of the reasonable to believe that ferrate(V1) has some potential as instrument, were not taken into account, and further effort a deeply sterilizing agent, including decontamination of air would be necessary for a closer simulation of the original and water in spacecrafts or space stations. It is noteworthy .. thatno hazardous or harmful wastes are produced with ferrate, iron(ll1) oxide being the only byproduct. SPECTRAL CHARACTERIZATION OF FERRATE(V1) Ferrate(V1)displays multiple spectral features making it a convenient material for spectral detection and quantitative determination. We discusshere a fewspecific methods which are suitablefor remote sensing,contact measure- ments in situ, and analysis of returned sample. Absorption and reflectance spectra A deep purple color which appears in highly alkaline iron- containing solutions when treated with strong oxidants was found centuries ago. This characteristic color is due to the 25005002000 1500 1000 nm tetrahedral ferrate(V1) anion, FeO,", with its absorption peak at 505-510 nm, andmolar extinction of 1100-1300 M.'cm" at these wavelengths. This spectrum makes ferrate (VI) verydistinct from other forms of dissolvediron, and makesabsorption spectrum anappropriate tool for ferrate(V1) identification and quantitation when samples are 60 available for dissolution. 2 With Mars exploration in view, a more practical alternative 50 us the use of reflectance spectra that can be recorded by remote sensing fromorbit. The broad-range reflectance spectrum of crystalline potassium ferrateis presented in Fig. 40 5, together with other forms of iron. The band at 1600 nm is acharacteristic one, and it canbe used as a markerto 30 search for the presenceof Fe(VI) on the Martian surface.

Accordingto Viking's results,the content of oxidant in 20 Martian soil is rather low, most likely in the ppm range. If this reactivematerial is the only one with ferrate(V1)spectral features, it certainlywouldn't be visible in reflectance spectra. However,we now know that combinations of ferrate(V1) anion with various metals may considerably differ V in both their watersolubility and reactivity. Most of 1600 1200 2000 2400 nm ferrate(Vl), if it is present at all, may be in the form of some non-reactivecompounds, of which barium ferrate is an Fig. 5. Reflectance spectraof iron compounds. 1 : Broad example. Though chemically rather resistant, they still give range spectrum of solid potassium ferrate; 2-7: spectra in the characteristic peak atca. 1600 nm, making the detection the range 1200-2400 nm. 2: Solid K,FeO,. 3: Fresh of ferrate(V1) likely by means of remote sensing reflectance moisturized precipitate of Goethite, FeOOH. 4: CaO-Fe,O, spectroscopy. mixture, 1O:l by mass. 5: CaO-Fe,O,-K,FeO, mixture, 20:2:1 by mass. 6: Solid FeSO,. 7: Solid Fe,(SO,),. Mossbauer Spectra Mossbauerspectroscopy is the mostpowerful and tral components. One component consists of a singlet line straightfotwardapproach to the identification of various with a negative isomer shift of -0.91 mm/s, and the other is oxidation states of iron. It is important in the context of these aquadrupole-split doublet with an isomer shift of +0.32 studies since a compactMossbauer instrument has been mm/sand quadrupole splitting of 0.57 mm/s.The relative developed for the future Mars missions and will be a part of abundance of these two components can be estimated from the Athena payload [36]. TheMossbauer spectrum of a the relative area of the two sub-spectra. The singlet com- samplecontaining potassium ferrate(V1) and otheriron ponent has a negative isomer shift, which is opposite in sign species is shown in Fig. 6, together with a diagram of the to theisomer shifts for Fe(ll)and Fe(lll) phases and is ranges of isomer shifts in ironcompounds with various identified as the Fe(VI) phase. The doublet component has valences and spin states, as referencedto iron metal at an isomer shift and quadrupole splitting consistent with an 300K. It is produced by a superposition of two main spec- r 0- 5 -Fe(6+) cd *.F Fe(3+) 4d :'J - - - 2- d 1.5 :A 1 - - - - - Fe(2+) : ;., If*

-20 0 20 40 60 80 100 -10 -8 -6 -4 -2 0 2 4 6 8 '10 Energy Relative to Fe foil (E0=7112 eV) Source Velocity (mmls) Fig. 7. XANES of solutions of iron compounds: Fe(VI)as KFeO,, Fe(lll) as Fe(NO,),, Fe(ll) as FeCI,.

i! Fe(ll) S=l solutions and solid samples containingtetrahedrally - coordinated Fe(VI) in ferratedianion. Additionally, more ire(l1) s=d subtle changes in X-ray absorption edge energy have been i Fe(lll) $=5/2 - shownto enable the identification of Fe(ll)and Fe(lll) i -Fe(lll) S=3@ F~III)s=i/2 valence states. -j Fe(lv) S=2 i j - Fe(iv) S=I j CONCLUSION I -Fe(Vl) S=l , i....i,..,i -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 To summarize, the data presented are consistent with the Isomer Shift (mmls vs. aFe) hypothesis of ferrate(V1) and possibly other higher ironcontributing to the oxidizingpool in the Martian soil, and provide spectraland chemical approaches to make Fig. 6. Mossbauer spectrumof KFeO,, with an admixture this hypothesis testable in forthcoming Mars missions. of Fe,O,, and the diagramof isomer shifts of iron in various oxidation states. ACKNOWLEDGMENTS

Fe(ll1)-bearing phase which is paramagneticatroom This workwas partially performed atJet Propulsion temperature. Laboratory,California Institute of Technology, under a contract with National Aeronautic and Space Administration X-Rav SDectra grant 100483.344.50.34.01 and was supported, in part, by the DRDF grant100656-00888. Support for KMK was X-Rayabsorption near edgespectra (XANES) were provided by the US DOE NABlR program and the Argonne obtained using the Material Research Collaborative Access National Laboratory LDRD program. Support for the work Team Sectorat the AdvancedPhoton Source, Argonne done at MRCAT sector of the APS was provided by the US National Laboratory. This technique relies on heavy, large- DOE,Office of Science, Office of BasicEnergy Sciences scale facilities, and is not usable in any space mission in the andgrant # DE-FG02-94ER45525. We thankGene foreseeable future. However, it has great potential for the McDonaldsfor GC-MS measurements, Cindy Grovefor analysis of returned samples, due to its sensitivity to the access to somespectral facilities, and AnnaLarsen structuralenvironment ironof atoms, and thus to its (Chemistry Dept., UC-Riverside) for thermogravimetric and oxidation state as well. The focus, in this technique, is on DSC data. the spectral regionvery close (within about30eV) to the ionization threshold of the compound in question. A XANES REFERENCES spectrum of ferrate(V1)sample is shown in Fig. 7. 1. Oyama, V.I., Berdahl, B.J. The Viking gas exchange Comparison of XANES spectra of a sample with XANES experiment results from Chryse and Utopia surface samples. spectra of a known standardallows identification of the J. Geophys. Res. 82,4669-4675 (1 977) oxidation state andlocal geometry around iron. For iron, 2. Levin, G.V., Straat, P.A. 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