Descent without Modification? The Thermal Chemistry of H 2 O 2 on Europa and Other Icy Worlds

Abstract

The strong oxidant H₂O₂ is known to exist in solid form on Europa and is suspected to exist on several other Solar System worlds at temperatures below 200 K. However, little is known of the thermal chemistry that H₂O₂ might induce under these conditions. Here, we report new laboratory results on the reactivity of solid H₂O₂ with eight different compounds in H₂O-rich ices. Using infrared spectroscopy, we monitored compositional changes in ice mixtures during warming. The compounds CH₄ (methane), C₃H₄ (propyne), CH₃OH (methanol), and CH₃CN (acetonitrile) were unaltered by the presence of H₂O₂ in ices, showing that exposure to either solid H₂O₂ or frozen H₂O+H₂O₂ at cryogenic temperatures will not oxidize these organics, much less convert them to CO₂. This contrasts strongly with the much greater reactivity of organics with H₂O₂ at higher temperatures, and particularly in the liquid and gas phases. Of the four inorganic compounds studied, CO, H₂S, NH₃, and SO₂, only the last two reacted in ices containing H₂O₂, NH₃ making \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} and SO₂ making \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} by H⁺ and e⁻ transfer, respectively. An important astrobiological conclusion is that formation of surface H₂O₂ on Europa and that molecule's downward movement with H₂O-ice do not necessarily mean that all organics encountered in icy subsurface regions will be destroyed by H₂O₂ oxidation. Key Words: Europa—Laboratory investigations—Icy moons—Infrared spectroscopy—H₂O₂ resistance. Astrobiology 15, 453–461.

1. Introduction

The possibility of life within and below the surface ices of Europa has been the subject of numerous studies following reports from NASA's Galileo mission, suggesting that liquid water may exist below the surface ice (Carr et al., 1998). Life beneath the ice shell was treated by Gaidos et al. (1999), while life within the ice itself, perhaps aided by rising subsurface thermal plumes, was examined by Ruiz et al. (2007). Aside from harboring living organisms, it has been speculated that Europa's ices also could contain evidence for subsurface biology in the form of molecular by-products or organic material (Marion et al., 2003) brought upward by convection (Pappalardo et al., 1998). Hand et al. (2009) summarized and extended much of the earlier work on Europa astrobiology and described possible scenarios for biogeochemical cycles and the putative subsurface ocean's chemistry.

The persistence and detection of any material within or beneath Europa's ice shell depend strongly on such material's response to the local environment of high-energy radiation, strong oxidants, and low temperatures. In general, one can divide the solid-phase chemistry that is expected to occur on Europa and other icy moons of the Solar System into that which is driven by UV photons (E ∼ 10 eV), by kiloelectronvolt and megaelectronvolt external particle radiation, and by thermal processes at about 30–130 K. The chemistry induced by UV photons is largely confined to surface depths of less than a centimeter, whereas ionizing radiation can alter the chemistry for several meters below the surface (Barnett et al., 2012). Thermal chemistry, which does not rely on an external energy source, will be active at all depths below the surface but may be more important at increasing depths as temperatures rise.

To date, telescopic and spectroscopic observations of Europa have probed its surface composition to depths of at best a few centimeters, with deeper subsurface chemistry remaining hidden. We and others have examined Europa's ice chemistry through extensive photo- and radiation-chemical laboratory experiments on icy materials (Hudson and Moore, 2001; Baratta et al., 2002; Hand and Carlson, 2011; Johnson et al., 2012). However, low-temperature thermal processes relevant to Europa have been studied much less frequently. In a recent paper, we described a thermally driven chemical reaction that can occur at europan temperatures even in the absence of external photons and ionizing radiation (Loeffler and Hudson, 2010). Amorphous ices consisting of H₂O and SO₂ showed changes in their IR spectra on warming from 50 to 100 K, demonstrating that H₂O and SO₂ combined to form H₃O⁺ and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^{ -}$$ \end{document} (bisulfite): \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}2{ \rm H}_2{ \rm O} + { \rm SO}_2 \rightarrow { \rm H}_3{ \rm O}^ + + { \rm HSO}_3^{ -} \tag{1}\end{align*} \end{document}

That this reaction took place in the absence of far-UV photons and ionizing radiation means that it also can occur beneath Europa's surface. Thus, the results we reported can be regarded as “hidden chemistry” in that they are changes that will occur at depths greater than those which current instrumentation can probe.

Moore et al. (2007b) showed that the irradiation of H₂O+SO₂ ice mixtures at 86–132 K with 0.8 MeV protons resulted in oxidation of SO₂ into \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} (sulfate). The suspected overall reaction was \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}4{ \rm H}_2{ \rm O} + { \rm SO}_2 \rightarrow 2{ \rm H}_3{ \rm O}^{ +} + { \rm SO}_4^{2 -} + { \rm H}_2 \tag{2}\end{align*} \end{document}

This radiation-driven process is a plausible source of the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} that was identified on Europa's surface by way of observations from the Galileo spacecraft's Near Infrared Mapping Spectrometer (NIMS) instrument (Carlson et al., 1999b). Since the ion irradiation of H₂O-ice already has been reported (Moore and Hudson, 2000) to produce hydrogen peroxide (H₂O₂), we suspected that H₂O₂ acts as an oxidizing agent for Reaction 2. This, in turn, led to our most recent study (Loeffler and Hudson, 2013), in which we examined whether adding H₂O₂ to solids rich in H₂O-ice could promote oxidation-reduction chemistry solely by thermal means. In those studies, we found that warming a H₂O+H₂O₂+SO₂ ice from 50 to 130 K initiated the formation of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^{ -}$$ \end{document} and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} through the following reactions: \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}2{ \rm H}_2{ \rm O} + { \rm SO}_2 \rightarrow { \rm H}_3{ \rm O}^{ +} + { \rm HSO}_3^{ -} \tag{3}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2 { \rm O}_2 + { \rm HSO}_3^{ -} \rightarrow { \rm H}_3{ \rm O}^{ +} + { \rm SO}_4^{2 -} \tag{4}\end{align*} \end{document}

Having shown that H₂O₂ can produce sulfate through Reaction 3 followed by Reaction 4, we now consider whether thermally induced H₂O₂ reactions also are important in other systems. The ability of frozen H₂O₂ to oxidize other molecules is of considerable interest to astrobiologists for several reasons. Greenberg (2010) summarized the case for downward transport of surface ices on Europa. Such movement would transfer surface material toward subsurface liquid and higher temperatures, carrying with it Europa's radiolytically generated H₂O₂ (Carlson et al., 1999a). If the H₂O₂ reaches the subsurface liquid, then it could significantly alter the chemistry occurring there, as the liquid would become more acidic (Pasek and Greenberg, 2012). Were subsurface oxidation by frozen H₂O₂ possible, then it could weaken the argument that organics (e.g., CH₄), originating from endogenic or exogenic sources (Zolotov and Kargel, 2009), trapped in ices could be available as nutrients for subsurface life. Conversely, any SO₂ that is oxidized by H₂O₂ could provide a potential nutrient for subsurface sulfate-reducing bacteria or their extraterrestrial analogues. There also is a connection to planetary protection in our interest in low-temperature H₂O₂ reactions. Such reaction chemistry can be used to understand the fate of molecules deposited by accident on icy Solar System surfaces and in subsurface regions that are shielded from direct exposure to UV light or ionizing radiation but which still may contain H₂O₂.

In the present study, we examined how H₂O₂ in H₂O-ice might promote chemical changes in the absence of ionizing radiation and far-UV photons. Our goal was to determine whether relatively simple molecules, specifically those listed in Table 1, can react in the solid state with H₂O₂. We point out that for this study we chose a relatively simple reaction system to compare the propensity of hydrogen peroxide to react with a variety of astrobiologically relevant molecules. Including more reactants or adding a known room-temperature catalyst (e.g., metal ions) might have affected our results. Although H₂O₂ is a potent room-temperature oxidant, to our knowledge the present study and our previous one (Loeffler and Hudson, 2013) are the first to investigate its oxidizing abilities in ices at temperatures relevant to the outer Solar System. Here, we first briefly review our earlier SO₂ work and then extend it and move on to other new results obtained by in situ measurements with IR spectroscopy.

Table 1.

Compounds Combined with H₂O₂ in Ices

Molecule	Seen in the interstellar medium? ^a	Seen in cometary comae? ^b	Seen on Solar System moons? ^c	Which objects in the Solar System?
Inorganic
SO₂	Yes	Yes	Yes	Galilean satellites^d,e
H₂S	Yes	Yes	Yes?	Suspected on Galilean satellites^e
NH₃	Yes	Yes	Yes	Charon^f and possibly Miranda^g
CO	Yes	Yes	Yes	Triton^h and Plutoⁱ
Organic
CH₄	Yes	Yes	Yes	Triton^j and Pluto^k
C₃H₄	No	No	Yes	Titan's atmosphere^l
CH₃OH	Yes	Yes	No	Centaur object Pholus^m
CH₃CN	Yes	Yes	Yes	Titan's atmosphereⁿ

See Tielens (2013). ^bSee Mumma and Charnley (2011). ^cSee, for example, Dalton et al. (2010), ^dLane et al. (1981), ^eMcCord et al. (1998), ^fBrown and Calvin (2000), ^gBauer et al. (2002), ^hCruikshank et al. (1993), ⁱOwen et al. (1993), ^jCruikshank et al. (1976), ^kCruikshank and Silvaggio (1979), ^lMaguire et al., (1981). ^mBased on modeling of spectra (Merlin et al., 2012) and ⁿMarten et al. (2002).

2. Experimental Methods

Experiments were performed with a cryostat (T _min∼10 K) operating in a stainless steel high-vacuum chamber (P∼1×10⁻⁷ torr) interfaced to an IR spectrometer. Ices were prepared by co-deposition of H₂O, H₂O₂, and an organic or inorganic molecule, chosen from those listed in Table 1, onto a pre-cooled (10–50 K) gold-coated aluminum mirror (area≈5 cm²) by using three separate pre-calibrated gas lines. Mixtures that contained CO and CH₄ were deposited at 10 K, but all other samples were prepared at 50 K. Pure H₂O₂ was prepared in a glass manifold as previously described (Loeffler and Baragiola, 2011). During deposition, the increase in the sample's thickness was monitored by interferometry with a diode laser (670 nm), and the deposition was halted when a thickness of about 1.5 μm was reached.

After deposition of a sample, its IR spectrum was recorded before, during, and after warming at 1 K min⁻¹. A closed-cycle helium cryostat and a resistive heater served to maintain ices at the desired temperature within the 10–200 K range studied. Higher temperatures could be reached but were avoided, as they were accompanied by sample loss by sublimation, which complicated the interpretation of our experiments.

Spectra were measured from 7000 to 400 cm⁻¹ with a Bruker Vector 22 Fourier transform infrared spectrometer at 2 cm⁻¹ resolution and with 200-scan accumulations. To obtain a spectrum, the IR reflectance (R) from the ice-coated substrate was divided by the reflectance of the bare metal substrate (R ₀), taken before ice formation, and then converted to an absorbance-type scale, -log(R/R ₀), for the figures we show in this paper. See the work of Loeffler and Hudson (2010, 2012) for additional experimental details.

Each compound studied was examined in the solid phase mixed with frozen H₂O₂, both in the presence and absence of H₂O-ice. During experiments, ices were held as long as 20 h at the higher temperatures (e.g., 100–170 K) to check for thermal changes. At the end of each experiment, the sample was warmed to room temperature, but no residual material was observed either visually or through IR spectroscopy.

The sensitivity of our spectrometer to detect the products of a thermally induced reaction under the conditions employed was checked with appropriate background and calibration experiments. For example, when working with CO we sometimes observed a weak CO₂ feature before warming with H₂O₂ (or ¹³CO₂ when ¹³CO was being studied). Blank experiments without H₂O₂ present traced this initial CO₂ in our samples to low-level CO₂ contamination, on the order of 0.002%, in our CO source and a small background contamination in our vacuum system. Possible products, such as H₂CO, CH₃OH, or SO₂, from other reactants have weaker IR absorptions than CO₂ but were still strong enough to be easily identified if present. We estimate that these possible products were detectable down to ∼1×10¹⁵ molecules cm⁻² or better, corresponding to a number ratio of ∼2×10⁻³ (∼0.2%) when compared to the initial reactants. Thus, any nondetection we reported corresponds to an upper limit of these values.

It should be emphasized that the design of our experiments allowed all data to be collected in situ at the temperatures of choice. This avoided the need to raise samples to room temperature for chemical analyses. Also, the co-condensation method used to prepare our ices avoided the uncertainties in freezing room-temperature solutions, which will crystallize to give regions of varying H₂O₂ concentration in the sample. In short, the value of our results is enhanced both by the way the samples were prepared and by our method of analysis. For recent amino acid work in which similar experimental methods were used, see a study by Gerakines and Hudson (2013).

The reagents and suppliers used were H₂O₂ (Sigma-Aldrich, 50% by weight), SO₂ (Matheson, 99.98%), H₂S (Matheson, 99.5%), ¹²CO (Matheson, 99.998%), ¹³CO (Sigma-Aldrich, 99% ¹³C), NH₃ (Matheson, 99.9992%), CH₄ (Matheson, 99.999%), C₃H₄ (Sigma-Aldrich, 98%), CH₃OH (Sigma-Aldrich, >99.8%), and CH₃CN (Sigma-Aldrich, 99.93%). The water used was triply distilled with a resistivity greater than 10⁷ Ω cm.

Finally, we point out that our use of the terms oxidation and reduction refers to the loss and gain of electrons as deduced by both formal charges and oxidation states. The meaning and method for assigning oxidation states continues to be debated and discussed (Gupta et al., 2014; Karen et al., 2014), but for now we still employ the IUPAC-recommended convention (McNaught and Wilkinson, 1997). For a careful review of redox (oxidation-reduction) chemistry in a planetary setting other than Europa, see the work of Nixon et al. (2012).

3. Results

3.1. Inorganic compounds

Table 1 lists the four inorganic compounds we studied. Sulfur dioxide (SO₂) was selected since we already have data for it under other conditions (Loeffler and Hudson, 2010, 2013), and hydrogen sulfide (H₂S) was chosen as a possible precursor to SO₂. Carbon monoxide (CO) and ammonia (NH₃) were studied as they are among the simplest of carbon- and nitrogen-containing compounds, respectively. Moreover, each of these molecules is known to be extraterrestrial, with SO₂ having been reported for Europa and the other molecules suspected to be present either from cometary or meteoritic delivery or as primordial material.

Figure 1 shows the IR features of a H₂O+SO₂ ice on warming from 50 K. The most obvious change seen is the appearance of vibrational bands near 1050 cm⁻¹ due to the formation of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^{ -}$$ \end{document} . Figure 2 shows a similar experiment but where H₂O₂ has been added to the ice; the strongest unobscured H₂O₂ absorptions are at 2840 and 1454 cm⁻¹. As the sample was heated, the H₂O₂ and SO₂ bands decreased, and absorptions of sulfur oxyanions appeared. Figure 3 quantifies these changes in a plot of IR band areas over time for H₂O₂ (2840 cm⁻¹), SO₂ (1150 cm⁻¹), and the sulfur oxyanions (1170–940 cm⁻¹), whose main component is \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} (1068 cm⁻¹). It is seen that the H₂O₂ feature in the sample began to decrease by ∼70 K, dropped by a factor of two by ∼125 K, and was at the noise level about 10 min after reaching 130 K. We note that on warming anhydrous H₂O₂+SO₂ ices, binary mixtures, we found no evidence for oxidation to make \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} . This agrees with, and confirms, an earlier computational study that found that the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^{ -}$$ \end{document} → \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} conversion requires H₂O, H₂O₂, and SO₂ (Vincent et al., 1997).

FIG. 1.

Infrared spectra of an H₂O+SO₂ (6:1) ice mixture made at 50 K and then warmed at 1 K min⁻¹ to 130 K. The top spectrum was recorded after the sample had been at 130 K for 40 min. Spectra have been offset for clarity.

FIG. 2.

Infrared spectra of an H₂O+SO₂+H₂O₂ (80:14:6) mixture deposited at 50 K and then warmed at 1 K min⁻¹ to 130 K. The top spectrum was recorded after the sample had been at 130 K for 40 min. All spectra have been offset for clarity.

FIG. 3.

Normalized areas for IR bands of H₂O₂ (2840 cm⁻¹), SO₂ (1150 cm⁻¹), and sulfur oxyanions (940–1170 cm⁻¹) in a H₂O+SO₂+H₂O₂ (80:14:6) mixture deposited at 50 K and warmed at 1 K min⁻¹ to 130 K. A temperature of 50 K corresponds to time=0, and the vertical line corresponds to 130 K. The sulfur oxyanions feature has contributions from \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm S}_2{ \rm O}_5^{2 -}$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_3^{ -}$$ \end{document} , \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} , and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm HSO}_4^{ -}$$ \end{document} .

Previously, we showed that ion irradiation of H₂O+H₂S ices near 100 K results in the formation of SO₂ at the expense of H₂S (Moore et al., 2007a). A straightforward explanation is that reactions such as (5) and (6) convert H₂S first into elemental sulfur (S) and then into SO₂, with both processes requiring H₂O₂ and perhaps involving numerous mechanistic steps (e.g., Hoffmann, 1977). \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2{ \rm O}_2 + { \rm H}_2{ \rm S} \rightarrow 2{ \rm H}_2 { \rm O} + { \rm S} \tag{5}\end{align*} \end{document} \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}2{ \rm H}_2{ \rm O}_2 + { \rm S} \rightarrow 2{ \rm H}_2{ \rm O} + { \rm SO}_2 \tag{6}\end{align*} \end{document}

To examine the ability of H₂O₂ to effect the H₂S→SO₂ conversion in unirradiated ices, we prepared frozen H₂O₂+H₂S and H₂O+H₂O₂+H₂S mixtures at 50 K and slowly warmed them to sublimation. We saw no evidence for H₂S decomposition and SO₂ formation, and the H₂O₂ bands remained relatively constant as the temperature increased. We conclude that no thermally induced redox reactions occurred between H₂O₂ and H₂S at even the highest temperature studied (∼200 K), or at least there was no formation of detectable products.

Turning to ammonia (NH₃), we can envision two types of reactions with H₂O₂. Our previous work (Moore et al., 2007a) showed that no reactions are seen on warming H₂O+NH₃ ices, which we verified in the present study (see Fig. 4). For the present experiments, an oxidation sequence for NH₃ such as \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm NH}_3 \rightarrow { \rm NH}_2{ \rm OH} \rightarrow { \rm NO}_2^{ -} \rightarrow { \rm NO}_3^{ -} \tag{7}\end{align*} \end{document}

FIG. 4.

Infrared spectra of an H₂O+NH₃ (10:1) mixture deposited at 50 K and then warmed at 1 K min⁻¹. Spectra have been offset for clarity.

can be envisioned, giving both \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_2^{ -}$$ \end{document} (nitrite) and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_3^{ -}$$ \end{document} (nitrate) ions. However, we also recognized that H₂O₂ is on the order of ten thousand times more acidic than H₂O (Evans and Uri, 1949), so an acid-base reaction of the type \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm H}_2{ \rm O}_2 + { \rm NH}_3 \rightarrow { \rm NH}_4^{ +} + { \rm O}_2{ \rm H}^{ -} \tag{8}\end{align*} \end{document}

might be possible. To test these ideas, we conducted warming experiments with and without H₂O-ice present. Figure 5 shows the IR spectrum of an anhydrous H₂O₂+NH₃ mixture deposited at 50 K and then warmed to sublimation. The deposited sample possessed absorptions that do not belong to either H₂O₂ or NH₃, indicating that a reaction occurred. The strong, sharp IR feature of NH₂OH near 1190 cm⁻¹ (Nightingale and Wagner, 1954) is not seen, arguing for this molecule's absence. However, there are multiple absorptions (see Discussion) that can be assigned to \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} (1467 cm⁻¹) and NH₄O₂H (2650, 1099, and 832 cm⁻¹).

FIG. 5.

Infrared spectra of an H₂O₂+NH₃ (1:1) mixture deposited at 50 K and then warmed at 1 K min⁻¹. Spectra have been offset for clarity.

All these observations can be interpreted to mean that Reaction 8 took place. Interestingly, a subsequent warming of this same sample had little effect on its spectrum until ∼180 K, where the sample appeared to crystallize. In addition, lowering the deposition temperature to ∼10 K did little to alter the initial spectrum, suggesting that the energy released on deposition was enough to cause Reaction 8 to occur. No IR features were found for the oxidized nitrogen-containing species of (7).

From this anhydrous binary system, we then examined a more complex, three-component sample. Figure 6 shows IR spectra of a three-component H₂O+H₂O₂+NH₃ ice (10:1:1) after deposition at 50 K and warming to sublimation. Initially, the only IR absorptions seen were for those of the three reactants. However, on heating the sample, the H₂O₂ bands began to decrease at ∼100 K and reached the noise level by 130 K. Accompanying this decrease, we observed an increase in the \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} feature near 1467 cm⁻¹. To highlight the differences in Figs. 5 and 6, in Fig. 7 we show spectra of ices containing H₂O₂ and NH₃ with H₂O present and absent, along with reference spectra of H₂O+NH₃ and H₂O+H₂O₂ ices.

FIG. 6.

Infrared spectra of an H₂O+H₂O₂+NH₃ (10:1:1) mixture deposited at 50 K and warmed at 1 K min⁻¹. Spectra have been offset for clarity.

FIG. 7.

Infrared spectra of four ice mixtures each at 50 K: (a) H₂O₂+NH₃ (1:1), (b) H₂O+NH₃ (10:1), (c) H₂O+H₂O₂ (10:1), and (d) H₂O+H₂O₂+NH₃ (10:1:1).

The fourth inorganic molecule we examined was CO. Ices made of H₂O₂+CO and H₂O+H₂O₂+CO were warmed from 10 to sublimation at 1 K min⁻¹ with an eye toward CO₂ formation. However, Fig. 8 shows that no CO₂ formation was observed above the estimated background level in our experiments, with the only major spectral changes corresponding to H₂O₂ crystallization near 160 K. We observed that much of the CO initially present sublimed from the ice as the temperature increased. We suspect that any residual CO present above ∼100 K was likely too small to produce a detectable amount of CO₂, with the large intrinsic strength of the latter's IR features aiding their detection.

FIG. 8.

Infrared spectra of a H₂O₂+¹³CO (1:1) mixture deposited at 10 K and warmed at 1 K min⁻¹. Spectra have been offset for clarity. The inset uses an expanded scale to illustrate the detection limits in our measurements. The ¹³CO₂ band shown at 10 K is an impurity on the level of 0.002% in our ¹³CO source.

3.2. Organic compounds

Each of the four inorganic molecules we combined with H₂O₂ presented a particular appeal, but the present work initially was motivated by the possibility of observing thermal reactions of organics below 200 K. Table 1 lists the four organic compounds we studied. Methane (CH₄) was selected for its simplicity, its presence in a variety of astronomical environments, and its potential presence in many astrobiological settings, both as a possible metabolic product and as a nutrient. Propyne (C₃H₄) was chosen as a typical unsaturated molecule and as one that is potentially more reactive than CH₄. Methanol (CH₃OH) and acetonitrile (CH₃CN) were taken as simple representatives of the alcohols and nitriles, respectively. Oxidation products can be written easily for each of these organic molecules, with those of C₃H₄ perhaps being the most varied. Suffice it to say, however, that no oxidation products were seen within the detection limits in any of our experiments with H₂O+H₂O₂+X ices (X=CH₄, C₃H₄, CH₃OH, and CH₃CN) or in the simpler H₂O₂+X ices. This suggests that these four molecules are not easily destroyed by thermal reactions with H₂O₂ in either H₂O-rich ices at 50–170 K or the simpler two-component anhydrous mixtures with frozen H₂O₂.

4. Discussion

4.1. H₂O₂ oxidation of SO₂ in ices

Thermal reactions between SO₂ and H₂O₂ have been investigated extensively within the atmospheric science community due to their importance in removing SO₂ from Earth's atmosphere (e.g., Clegg and Abbatt, 2001). Based on those studies, (3) and (4) are believed to be the primary reactions between H₂O₂ and SO₂ in the presence of H₂O. Although Earth's clouds are generally much warmer than surfaces of icy satellites, the SO₄ ^2- we have seen in our experiments is likely from similar processes. Critical tests of our analysis of the chemistry are shown in Fig. 3. As required by our interpretation, the decrease in the IR features of the limiting reagents SO₂ and H₂O₂, near 60 K in our experiments, is accompanied by a simultaneous increase in the IR absorbance of sulfur oxyanions, primarily \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} . More significantly, and as expected, the locations of the inflection points for the deceases in reactants and the increase in products are essentially the same (∼130 K).

The subsurface depth on Europa to which these H₂O₂ observations, and others in this paper, apply will depend on the thickness and other properties of that satellite's ice shell. Numerous studies have addressed the ice shell's thickness and have yielded values between about 1 and 40 km (Billings and Kattenhorn, 2005). Taking 100 K as the temperature of surface ices and 270 K for the bottom of the ice shell, as is typically done (e.g., Bray et al., 2014), and making a simple assumption of linearity, we estimate a temperature gradient of 17 K km⁻¹ for an intermediate shell thickness of 10 km. Thus, the highest temperature reported in this paper (∼190 K) corresponds to a depth of about 5 km below Europa's surface.

4.2. H₂O₂ reactions with NH₃

Besides SO₂, the only other molecule in Table 1 that reacted with H₂O₂ was NH₃. However, unlike SO₂, where we observed that H₂O-ice was needed to initiate the reaction, NH₃ combined with H₂O₂ in the absence of H₂O. The first study to show that NH₃ and H₂O₂ react to form a stable compound was published nearly 100 years ago by Maass and Hatcher (1922). The white crystalline powder that formed in their experiments melted near 25°C, leading Knop and Giguere (1959) to propose that the bonding in the reaction product might be ionic and not the weaker bonding found in hydrates of NH₃ and H₂O₂. Knop and Giguere's spectroscopic work, along with a previous study (Simon and Kriegsmann, 1955), suggested that the new ionic species was ammonium hydroperoxide (NH₄O₂H), consisting of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} and O₂H^- ions, a suggestion later confirmed by X-ray diffraction measurements (Tegenfel and Olovsson, 1966; Churakov et al., 2010). That \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} is present in our H₂O₂+NH₃ ices is indicated clearly by an IR feature near 1467 cm⁻¹, suggesting that the protonation reaction (8) occurs more readily than the redox reaction (7). The likely counter-ion is O₂H^-, to which broad IR absorptions at 1100 and 836 cm⁻¹ have been assigned (Knop and Giguere, 1959). While only the lower wavenumber band is unobstructed at 50 K, Fig. 6 clearly shows both absorptions (∼1099, ∼832 cm⁻¹) after the sample was warmed and allowed to crystallize. Interestingly, the strong IR band seen near 2560 cm⁻¹ in Fig. 7a appears to be outside the spectral range of either of the two previous studies of H₂O₂+NH₃ ices. It is reasonable to assign at least part of this broad IR feature to an O-H or N-H stretching vibration in NH₄O₂H (e.g., Churakov et al., 2010).

A comparison of Figs. 5 and 6 shows that the addition of excess H₂O-ice to the H₂O₂+NH₃ system significantly changed the IR spectra and the associated chemistry. Warming the 50 K sample produced \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} at the expense of H₂O₂ as shown by the significant sharpening near 1467 cm⁻¹ and the dramatic decrease at 2850 cm⁻¹, respectively, with both changes being seen even at the highest temperature studied. One interpretation is that solvent-assisted protonation gave the compound (NH₄)₂O₂ or one of its hydrates in our ices. Although our spectra do not allow a firm assignment to be made, hydrate formation agrees with the report of (NH₄)₂O₂·2H₂O in the H₂O+H₂O₂+NH₃ ice system, based on thermal analyses (Giguere and Chin, 1959), and is consistent with observations that NH₄O₂H readily decomposes in the presence of H₂O (Maass and Hatcher, 1922; Knop and Giguere, 1959), perhaps to give OH^- at the expense of O₂H^-. In any case, the formation and persistence of \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} is readily seen in Fig. 6 for an H₂O+H₂O₂+NH₃ ice.

4.3. H₂O₂ reactions and astrobiology

Redox reactions with liquid- and gas-phase H₂O₂ readily occur at room temperature and above but not necessarily at temperatures expected in the outer Solar System. In the present work, we were particularly interested in determining how the eight molecules we examined (Table 1), which included several with biological connections, might respond to frozen H₂O₂ at low temperatures. The concentrations of H₂O₂ used in our experiments are much higher than what has been detected (Carlson et al., 1999a) and are expected in extraterrestrial environments, so negative results in our work (i.e., no reaction seen) suggest no reaction under less-harsh, more-realistic conditions.

For our experiments with solid H₂O₂, we selected as reactants the four organic and four inorganic compounds of Table 1. Of the inorganics, no thermally induced reactions were observed between H₂O₂ and either CO or H₂S. Reaction of the former could have made CO₂, while the latter could have produced SO₂, as given in (5) and (6). No such thermal conversion was seen. Extraterrestrial thermal oxidation of H₂S at low temperatures appears to require a more complex set of conditions than those we employed, perhaps even biological conditions akin to those relevant to sulfur-oxidizing bacteria. In contrast to the reluctance of solid H₂S to react with H₂O₂, here and elsewhere we have documented the relative ease of the SO₂→ \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} conversion (Loeffler and Hudson, 2013). The activation energy derived from an Arrhenius plot in our earlier work led to half-life predictions for SO₂ of about 1 year at 100 K and about 0.3 h at 120 K. The reported subsurface ocean of Europa is presumably at temperatures higher than those of the surface (∼86–130 K), so any subsurface H₂O₂ present, such as from downward transport from the surface itself (Greenberg, 2010), will readily convert SO₂ into \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} , a potential energy source for certain extremophiles such as sulfur-reducing bacteria.

Ammonia (NH₃) was the fourth inorganic compound we examined. Our initial thoughts were of a possible thermal nitrification of NH₃ in H₂O₂-containing ices to make the oxyanions \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_2^{ -}$$ \end{document} and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_3^{ -}$$ \end{document} , as in (7). Chemical reactions indeed were observed on warming ices that contained H₂O₂ and NH₃, but neither \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_2^{ -}$$ \end{document} nor \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NO}_3^{ -}$$ \end{document} was seen. The best that can be said is that under the conditions of our experiments, NH₃ was consumed in H₂O₂-containing ices at ∼100 K and higher. Although the observed reaction products differed in the presence and in the absence of H₂O (see Figs. 5 and 6), \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} was formed in all cases and so would be available as a nutrient for any extraterrestrial counterparts of terrestrial nitrifying bacteria.

Turning to the organic compounds we examined, solid methane (CH₄) and solid methanol (CH₃OH) conceivably could have undergone multiple thermal oxidations in the presence of H₂O₂, such as through the sequence \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document} \begin{align*}{ \rm CH}_4 \rightarrow { \rm CH}_3 { \rm OH} \rightarrow { \rm H}_2 { \rm CO} \rightarrow { \rm CO} \rightarrow { \rm CO}_2 \tag{9}\end{align*} \end{document}

to make CO₂, but we observed no CO₂ on warming ice mixtures containing either CH₄ and H₂O₂ or CH₃OH and H₂O₂. Each product molecule in (9) has at least one fairly strong IR spectral feature that would easily have been detected, but none were seen in the 10–150 K region examined. At room temperature and above, propyne (C₃H₄) can undergo oxidative cleavage to make CO₂ and CH₃COOH (acetic acid), but no evidence for either was found in our low-temperature experiments, and both CO₂ and CH₃COOH would have been relatively easy to detect with our IR methods. Finally, from earlier work we already knew that low-temperature radiation-induced oxidation of acetonitrile (CH₃CN) first makes OCN^- and then CO₂ (Hudson and Moore, 2004), but neither of these products was seen in our warmed, unirradiated CH₃CN-containing ices.

This leads to perhaps the most significant astrobiological implication of our experiments with organics, namely, the stability of these four typical, relatively simple organic molecules in the presence of solid H₂O₂. One can envision a situation in which, for example, vertical transport of material on Europa brings surface oxidants, such as H₂O₂, down to subsurface regions and into contact with organics trapped in ice (Greenberg, 2010). Our experiments showed that such encounters do not necessarily destroy those same organic compounds. In the particular case of CH₄, whether it is primordial, endogenous, or exogenous, it would remain available for methanotrophs.

Finally, this study emphasizes that understanding and predicting Europa's low-temperature solid-phase chemistry, and its contribution to astrobiology, will require substantial laboratory efforts. For example, with standard tables of E ^θ values (reduction potentials) for half-cell reactions, one can predict that all of the eight compounds we combined with H₂O₂ should be oxidized by that molecule, which does not agree with our observations. However, such reference data are nearly always for a standard set of conditions near 25°C, 1 bar, and 1 molar concentration, conditions markedly different from those of Europa ices. Furthermore, extrapolating from calculations on systems with liquid phases to those, like ours, that are entirely formed from solids is not straightforward. Since the influence and interplay of both thermodynamic and kinetic contributions to ice chemistry cannot yet be predicted a priori, we anticipate a significant role for experimental work for the foreseeable future.

5. Summary and Conclusions

Laboratory results have been presented here on the reactivity of H₂O₂ with selected organic and inorganic molecules in ice mixtures at temperatures of the outer Solar System. Using IR spectroscopy, we observed that two of the inorganic compounds studied, SO₂ and NH₃, were consumed in H₂O₂-containing ices by thermally induced reactions, yielding \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm NH}_4^{ +}$$ \end{document} and \documentclass{aastex}\usepackage{amsbsy}\usepackage{amsfonts}\usepackage{amssymb}\usepackage{bm}\usepackage{mathrsfs}\usepackage{pifont}\usepackage{stmaryrd}\usepackage{textcomp}\usepackage{portland, xspace}\usepackage{amsmath, amsxtra}\pagestyle{empty}\DeclareMathSizes{10}{9}{7}{6}\begin{document}$${ \rm SO}_4^{2 -}$$ \end{document} , respectively. These ions are predicted to exist in both surface and subsurface europan ices, although in the latter case they will be hidden from remote observations.

In contrast, at the timescales of our experiments we found that none of the four organic compounds examined underwent low-temperature thermally induced reactions with solid H₂O₂ in either the presence or absence of H₂O-ice. We believe that these observations of a nonreaction speak positively for future attempts to find organics on Europa. On a frozen world where H₂O₂ could move downward to meet trapped subsurface organics, it could descend without modification by them, leaving those organics unaltered, so we can anticipate their discovery through future explorations. Put another way, the formation of surface radiolytically generated H₂O₂ on Europa, and that molecule's downward movement with H₂O-ice, does not necessarily mean that all organics encountered in icy subsurface regions will be destroyed by H₂O₂ oxidation.

Footnotes

Acknowledgments

The support of NASA's Planetary Geology and Geophysics and Outer Planets Research programs is gratefully acknowledged. The authors particularly acknowledge support from the NASA Astrobiology Institute through a grant to the Goddard Center for Astrobiology. Perry Gerakines is thanked for assistance in day-to-day operations of the equipment in our laboratory.

Author Disclosure Statement

No competing financial interests exist.

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Descent without Modification? The Thermal Chemistry of H 2 O 2 on Europa and Other Icy Worlds

Abstract

1. Introduction

2. Experimental Methods

3. Results

3.1. Inorganic compounds

3.2. Organic compounds

4. Discussion

4.1. H2O2 oxidation of SO2 in ices

4.2. H2O2 reactions with NH3

4.3. H2O2 reactions and astrobiology

5. Summary and Conclusions

Footnotes

Acknowledgments

Author Disclosure Statement

References

4.1. H₂O₂ oxidation of SO₂ in ices

4.2. H₂O₂ reactions with NH₃

4.3. H₂O₂ reactions and astrobiology