All Gale crater analog water‐producing phases evolved high‐temperature water attributed to smectite dehydroxylation, structural water, or inclusions. Peak water release temperatures were 841, 495, and 790 °C for SapCa‐1, NAu‐1, and Stapafell basaltic glass, respectively (Figure 5 ). Negligible amounts of O 2 and HCl were evolved from Stapafell basaltic glass and NAu‐1 (Figure 5 ). SapCa‐1 did not evolve O 2 but did evolve a small HCl release with a peak ~841 °C (Figure 5 ), which may be caused by minor (<0.00 wt.%; Table 4 ) chloride contamination because it was naturally sourced. This HCl evolution was taken into account when interpreting the results of the perchlorate and chloride mixtures with SapCa‐1. The EGA results from the individual water‐producing phases demonstrated that they yield water that should be available for HCl production during the SAM‐EGA analog analyses described below.

In the case of the NaClO 4 /nontronite mixture, HCl production was catalyzed by the addition of reactants (NaCl and H 2 O (g) ; reaction (7)) as well as the consumption of products (Na 2 O; reactions (7) and (8)). The rapid increase in NaCl (from perchlorate decomposition) and water vapor (from nontronite dehydroxylation) as well as the presence of Si and Al oxides in the sample shifted reaction (7) to the right and increased HCl production. Si and Al oxides formed as solid products of the thermal dehydroxylation of nontronite (Table 2 ). XRD and Near Infrared Spectroscopy analyses of nontronite demonstrated that it loses its crystalline structure ~475 °C but Al‐OH bonds remain >975 °C (Gavin & Chevrier, 2010 ). There was no additional HCl release at the peak melting temperature of NaCl because the NaCl was likely completely converted to HCl, Na 2 O, and phases represented as (B) via reactions (7) and (8) or because there was insufficient H 2 O (g) present at the NaCl melting temperature to react with any remaining NaCl in the sample.

Reaction (7) can be catalyzed by the addition of phases represented as (A) in reaction (8), including SiOand Al, which react with NaO to produce phases represented by (B). The consumption of NaO shifts reaction (7) to the right and increases HCl production even below the melting point of NaCl (Uchida & Kamo,). Laboratory mixtures of pure SiOand Alwith NaClproduced more HCl when heated compared to NaCl by itself (Uchida & Kamo,).

NaClOmixed with nontronite evolved HCl at a lower temperature than the NaClO/saponite and NaClO/basaltic glass mixtures because nontronite evolved water due to dehydroxylation at a relatively lower temperature (497 °C; Figure 5 ). Na perchlorate thermally decomposed to solid NaCl, which reacted with water vapor to form HCl and is shown by equilibrium reaction (7) (Uchida & Kamo,).

Thermal data showed that the endotherm peak temperature associated with the melting of NaCl occurred before the high‐temperature HCl release peak in the NaClO 4 /saponite and NaClO 4 /basaltic glass mixtures (Figures 8 and S8). Additionally, the magnitude of the HCl release was affected by the amount of evolved H 2 O. The HCl evolution from the NaClO 4 /basaltic glass mixture was lower in magnitude than the NaClO 4 /saponite mixture because the basaltic glass evolved less water than saponite (Figure 5 ).

Evolved water reacting with melting chlorides resulted in evolved HCl peak temperatures at 795, 815, and 497 °C, for the NaClOmixed with SapCa‐1, Stapafell basaltic glass, and NAu‐1, respectively. Na perchlorate by itself evolved Owith a peak at 578 °C but did not evolve HCl because it decomposes to a chloride rather than an oxide (Figure 8 ; reaction (5)). When Na perchlorate was mixed with water‐producing phases, the chloride then melted and reacted with evolved water to produce high‐temperature HCl (reaction (6); Figure 8 , S8, and S9; e.g., Fraisslera et al.,).

The formation of chlorides was confirmed in the saponite/Mg (ClO 4 ) 2 mixtures using ion chromatography. Results demonstrated that chlorides were present in the heated saponite/Mg (ClO 4 ) 2 mixture but not the heated saponite (Table 4 ). The specific type of chloride could not be determined using this method because it only analyzes dissolved ions, and the weight percentage of chloride was too low to be detected by X‐ray diffraction.

The high‐temperature HCl release in the Mg (ClO 4 ) 2 /saponite mixture occurred due to evolved water reacting with melting chlorides including CaCl 2 and NaCl, although their melting endotherms were obscured by the saponite dehydroxylation endotherm (Figures 6 , S4, and S5). Additionally, the high‐temperature HCl release was not observed when Mg (ClO 4 ) 2 was mixed with phases that did not contain Na or Ca, including nontronite and Stapafell basaltic glass (Figure 6 ).

Mg (ClOmixed with saponite evolved an additional high‐temperature HCl release with a peak at 824 °C and was caused by melting chlorides that formed and reacted with water from saponite dehydroxylation (Figure 6 ). Chlorides likely formed on particle surfaces by a reaction between evolved HCl (from Mg (ClOdecomposition) and exchangeable cations (e.g., Na and Ca) in the saponite interlayer (Table 2 ; reactions (3) and (4)).

An additional high‐temperature HCl release was only observed in the Mg (ClO 4 ) 2 /saponite mixture, and not in mixtures of Mg (ClO 4 ) 2 with nontronite or Stapafell basaltic glass (Figure 6 ). Mg (ClO 4 ) 2 mixed with nontronite or Stapafell basaltic glass did not evolve additional high‐temperature HCl releases because there were no chlorides present in the mixtures to react with the high‐temperature water. This is because Mg (ClO 4 ) 2 decomposes to an oxide rather than a chloride, as shown in reactions (1) and (2).

HCl release peaks from Mg (ClOand Mg (ClOmixed with SapCa‐1, NAu‐1, and Stapafell basaltic glass occurred between 485 and 545 °C due to perchlorate decomposition and at 825 °C in only the mixture with SapCa‐1 due to evolved water reacting with melting chlorides (Figure 6 ). Previous studies demonstrated that Mg (ClOthermally decomposes to MgO and evolves Oand Clgas (reaction (1); Figure S4; Marvin & Woolaver,; Acheson & Jacobs,). Water from the hydrated perchlorate then reacts with the Clto produce HCl, with a peak at approximately 542 °C (reaction (2); Figures 6 7 , and S4; Marvin & Woolaver,). Results from Mg (ClOanalyzed by itself agreed with previous studies.

Laboratory detections of evolved O 2 and HCl from NaCl without water‐producing phases and NaCl mixed with 10 mg of SapCa‐1 saponite, 20 mg NAu‐1 nontronite, or Stapafell basaltic glass (SBG). Green vertical dashed line indicates the melting endotherm peak temperature of NaCl (Figure S8). Complete evolved gas and differential scanning calorimetry/thermal gravimetry data are provided in Figures S17–S20. Ion current unit is amperes.

NaCl mixed with SapCa‐1, Stapafell basaltic glass, and NAu‐1, evolved HCl release peaks at 796 °C, 795 °C, and 522/770°, respectively, due to evolved water reacting with the melting chloride (Figure 10 ). NaCl by itself did not evolve O 2 or HCl within the tested temperature range, as expected. Mixtures of NaCl with saponite or basaltic glass evolved HCl after the melting point of NaCl and at similar temperatures as H 2 O evolutions (Figures 10 , S8 , S10 , and S11 ). The HCl release from the NaCl/nontronite mixture had peaks at 522 and 770 °C, which were similar to the temperatures of nontronite dehydroxylation and the NaCl melting endotherm peak, respectively (Figures 5 and S20). The first HCl release (peak 522 °C) was caused by a similar mechanism as in the NaClO 4 /nontronite mixture, as described in section 3.2 and shown by reactions (7) and (8). An additional HCl release (peak 770 °C) occurred because NaCl, which was not completely consumed via reaction (8), melted and was more reactive with water present in the oven from lower temperature reactions.

MgCl 2 mixed with saponite evolved an additional HCl release with a peak at 809 °C due to chlorides that formed, melted, and reacted with dehydroxylation water (Figure S14 ). Chlorides (e.g., NaCl and CaCl 2 ) were produced when HCl from MgCl 2 decomposition reacted with interlayer exchange cations (e.g., Na and Ca) in the saponite, similar to what happened in the Mg (ClO 4 ) 2 /saponite mixtures (reactions (3) and (4)). Evolved HCl from the MgCl 2 /nontronite mixture was broadly similar to MgCl 2 by itself (Figure 9 ). MgCl 2 mixed with Stapafell basaltic glass evolved a mid‐temperature HCl release and an additional small high‐temperature HCl release with a peak at approximately 805 °C, possibly caused by residual Cl 2 from chloride decomposition reacting with evolved water (Figures 9 and S15 ). The low temperature (<300 °C) HCl release in the MgCl 2 /basaltic glass mixture was caused by MgCl 2 dehydration (Huang et al., 2011 ). The MgCl 2 /basaltic glass mixture may have been exposed to atmospheric moisture for longer than the other mixtures, leading to a higher amount of low‐temperature evolved H 2 O (Figure S15 ).

MgCl 2 by itself produced more HCl than the MgCl 2 /water‐producing phase mixtures because it evolved more low‐temperature H 2 O, likely because the sample had a higher surface area exposed to atmospheric moisture or because it was exposed to atmospheric moisture for a longer duration.

Laboratory detections of evolved O 2 and HCl from MgCl 2 and MgCl 2 mixed with 10 mg of SapCa‐1 saponite, 20 mg NAu‐1 nontronite, or Stapafell basaltic glass (SBG). Blue and green vertical dashed lines indicate melting endotherm peak temperatures for CaCl 2 and NaCl, respectively (Figure S8). Complete evolved gas and differential scanning calorimetry/thermal gravimetry data are shown in Figures S13–S16. Ion current unit is amperes.

MgClby itself and mixed with SapCa‐1, NAu‐1, and Stapafell basaltic glass evolved HCl release peaks between 435 and 535 °C due to chloride thermal decomposition. MgClby itself, as expected, did not evolve any Obut did evolve HCl with a peaks at approximately 473 and 532 °C (Figures 9 and S13 ). MgClevolved smaller HCl releases with peaks at ~160 and ~215 °C as products of dehydration, leading to an intermediate phase of MgOHCl (Figure S13 ; Huang et al.,). MgOHCl then completely decomposed to MgO and HCl (Figures 6 and 9 ; reaction (9); Huang et al.,). The condensed thermal decomposition reaction is shown in reaction (9) and a complete step‐by‐step thermal decomposition mechanism is presented in Huang et al. ().

3.4 Comparison of Laboratory Mixtures to SAM Samples

3.4.1 Samples That Coevolved O 2 and HCl Below 600 °C The evolved O 2 and HCl between approximately 200 and 600 °C in the Cumberland (CB) and Big Sky (BS) mudstone samples were consistent with being sourced from Mg perchlorate or chlorate (Figures 2 and 6). Mg chlorate thermally decomposes similarly to Mg perchlorate and cannot be ruled out as a source of O 2 and mid‐temperature HCl in CB (Hogancamp et al., 2018). BS evolved two midtemperature HCl releases with peaks at 289 and 424 °C, both of which could have been caused by the decomposition of Mg perchlorate (Figure 2). The HCl release peak at ~424 °C in BS could not be caused by Mg chlorate because Mg chlorate decomposes at a lower temperature (~380 °C; Hogancamp et al., 2018). Fe oxychlorines also evolve mid‐temperature HCl; however, their O 2 release peaks occur at lower temperatures than observed in CB and BS (Glavin et al., 2013). CB and BS also evolved high‐temperature HCl releases (>600 °C), which could be caused by phyllosilicate dehydroxylation water reacting with original chlorides and/or chlorides from oxychlorine decomposition. Although CH and TP evolved HCl with peaks <600 °C, they did not coevolve with O 2 and were therefore inconsistent with being sourced from Mg oxychlorines.

3.4.2 Samples With O 2 and High‐Temperature (>600 °C) HCl Evolved O 2 and HCl from the JK, CH, WJ, MJ, TP, BK, GH, RH, and all eolian samples (RN, GB1, and GB2) were consistent with being sourced from oxychlorines other than Mg‐ phases (Figures 3 and S21–S23). These samples evolved O 2 between 200 and 600 °C but evolved HCl at higher temperatures (peaks >600 °C), which was consistent with the thermal decomposition of NaClO 4 mixtures (e.g., Figure 8). K‐ and Ca‐ oxychlorine phases, while not examined in this work, could also have contributed to the evolved O 2 and higher temp HCl releases. This is because K‐ and Ca‐ oxychlorine phases thermally decompose similarly to Na‐oxychlorine phases yielding O 2 and chlorides that would be available for HCl production at higher temperatures. 2 and peaked below the melting points of common chlorides (Figures S22 and 23). Chloride melting temperatures could be reduced due to impurities in the sample, which are known to decrease and broaden their melting temperatures (e.g., Pavia et al., 1982 2(g) and O 2(g) on the reaction between chlorides and H 2 O (g) . Similarly to Si and Al oxides, SO 2(g) and O 2(g) have been shown to catalyze the reaction between chlorides and H 2 O (g) through a conjugate reaction, increasing the rate of HCl production (reaction (10); Henriksson & Warnqvist, 1979 1983 (10) Several samples (GB1, BK, TP, WJ, MJ, and CH) evolved high‐temperature HCl that did not coevolve with water or Oand peaked below the melting points of common chlorides (Figures S22 and 23). Chloride melting temperatures could be reduced due to impurities in the sample, which are known to decrease and broaden their melting temperatures (e.g., Pavia et al.,). Additionally, the HCl releases observed in these samples may be caused by the catalyzing effect of SOand Oon the reaction between chlorides and H. Similarly to Si and Al oxides, SOand Ohave been shown to catalyze the reaction between chlorides and Hthrough a conjugate reaction, increasing the rate of HCl production (reaction (10); Henriksson & Warnqvist,; Uchida & Kamo,). WJ, MJ, TP, BK, CH, and GB1 evolved HCl after the evolution of O 2 from perchlorate decomposition and at similar temperatures as SO 2 evolutions (Figure 11). SO 2 and O 2 evolutions in these samples may have contributed to the HCl releases that peaked below the melting point of common chlorides. Figure 11 Open in figure viewer PowerPoint Evolved SO 2 , O 2 , and HCl detected by SAM‐EGA of Windjana (WJ), Mojave (MJ), Telegraph Peak (TP), Buckskin (BK), and Gobabeb 1 (GB1). cps=counts per second. All samples use use oven 2‐model 1 temperature logs.

3.4.3 Samples With High‐Temperature (600 °C) HCl and No O 2 The high‐temperature HCl releases without evolved O 2 in DL, QL, MB, OU, HF, and ST support the presence of chlorine‐bearing phases, such as chlorides, without the presence of oxychlorines (Figures 9, 10, S24, and S25). The only samples in this group that evolved mid‐temperature HCl release peaks characteristic of MgCl 2 were ST and DL, although they also peaked at a higher temperature (Figures S24 and S25). QL, MB, OU, and HF likely contained chlorides other than MgCl 2 (e.g., NaCl), which reacted with water in the system from phyllosilicate dehydroxylation or from another water‐producing phase and produced high‐temperature HCl.

3.4.4 Samples With Mid‐temperature H 2 O Without Coevolving HCl Mid‐temperature H 2 O without coevolving HCl was observed in several SAM samples (BS, RH, QL, GB1, GB2, GH, CH, BK, TP, MJ, DL, QL, and HF) and may be due to evolved water in the SAM samples contacting phases other than chlorides. The laboratory mixtures of NaClO 4 or NaCl with nontronite produced coevolving mid‐temperature HCl and H 2 O; however, the water from nontronite dehydroxlation only had one other phase to react with. SAM samples contained multiple phases (e.g., feldspars, olivines, and pyroxenes) other than perchlorates, chlorides, and phyllosilicates, which were all in contact with the evolved water. The decreased direct contact between chlorides and H 2 O (g) in the SAM samples may have caused negligible mid‐temperature HCl production from chloride/H 2 O (g) interaction.

3.4.5 New and/or Unexpected Results From This Work These analog experiments provided the first laboratory evidence that high‐temperature HCl releases detected by the SAM‐EGA in Gale crater samples were caused by water‐producing phases reacting with chlorides (original or from oxychlorine decomposition). These experiments also demonstrated that the HCl release temperature was dependent on several factors including the type of oxychlorine or chloride, the temperature at which water evolved, and the presence of catalyzing phases such as Si or Al oxides. Mid‐temperature (<600 °C) HCl releases were expected in Mg perchlorate or chloride mixtures with water‐producing phases but were also observed in mixtures of Na perchlorate and chloride with nontronite. High‐temperature (>600 °C) HCl releases in mixtures of Na perchlorate or chloride with water‐producing phases were expected after the chloride melting point, when Cl is more volatile and reactive with water vapor. However, mixtures of Na perchlorate and chloride with nontronite produced mid‐temperature HCl releases below the melting point of NaCl. These mid‐temperature HCl releases were caused by two factors: (1) the rapid input of H 2 O (g) from nontronite dehydroxylation and (2) the increased reactivity of NaCl (s) with H 2 O (g) due to the presence of catalyzing phases, including Al and Si oxides. Further SAM analog laboratory experiments will directly test the effect of catalyzing phases (SiO 2(s) , Al 2 O 3(s) , and SO 2(g) ) on HCl production. Although concurrent mid‐temperature H 2 O and HCl releases has not been observed in SAM data, they may be observed in future samples analyzed by SAM, in particular those that contain dioctahedral smectites. The mixture of Mg perchlorate with saponite produced mid‐ and high‐temperature HCl releases characteristic of the presence of chlorides but was not expected because Mg perchlorate decomposes to an oxide rather than a chloride. The high‐temperature HCl release was caused by the reaction between chlorides that formed and then reacted with evolved water from saponite dehydroxylation. Chlorides formed by the reaction between HCl from Mg perchlorate decomposition and cations in the interlayer exchange sites in the saponite. The presence of mid‐ and high‐temperature HCl releases in samples analyzed by SAM may be caused by water reacting with original chlorides, chlorides from oxychlorine decomposition, or chlorides that formed through a reaction between HCl and other phases in the sample (e.g., saponite).