The powder XRD patterns of all oxides of antimony show only sharp peaks of one phase (see supporting electronic information). The refined lattice parameters are summarized in Table 2.

Special attention was paid to the hydration state of Sb2O5⋅nH2O as the amount of H2O strongly influences the calorimetric results. The color of the initial product changed after the TG analysis from yellow to white, indicating H2O loss and possible partial reduction. The total weight loss measured for our Sb2O5⋅nH2O sample was 9.83 wt %. According to , Sb2O5⋅nH2O gradually loses water and becomes H2O-free at around 843 K. Additional weight loss was attributed to partial reduction of antimony and release of O2 gas. Using the same interpretation of our TG data, a weight loss of 9.28 wt % is attributable to H2O, and the remaining weight loss of 0.55 wt % is due to the partial reduction of the Sb2O5.

The H2O content determined for Sb2O5⋅nH2O was taken into account for the reduction of calorimetric data. For the calorimetric experiments, the sample Sb2O5⋅nH2O with 9.28 wt % H2O was used. The presence of H2O in the sample was corrected with the assumption that this water is loosely bound and its heat content is equal to the heat content of free (unbound) H2O (H1073K-H298.15K=73 kJ mol-1). Therefore, Tables 3 and 4 contain reactions and reaction enthalpies that refer to Sb2O5 and not Sb2O5⋅nH2O.

The chemical reactions considered in the calorimetric experiments are summarized in Table 3. The drop solution enthalpies of antimony oxides in lead borate and sodium molybdate are listed in Table 4. The final oxidation state of the antimony in the lead borate melt under oxygen flushing and bubbling is assumed to be 5+. There are no experimental data to confirm this assumption, but the experience with similar systems e.g., arsenic in oxide melts, and the magnitudes of the measured enthalpies support the assumption. No difficulties or irregularities were encountered in the experiments with lead borate. The dissolution of the antimony oxides in sodium molybdate results in more scattered data than in lead borate due to baseline shifts. A possible explanation is a reaction of the antimony with the sodium molybdate solvent and the formation of an unknown refractory compound. Nevertheless, consistent drop solution enthalpy data can be obtained if the number of drops is increased. For sodium molybdate, the results for Sb2O5 are more consistent (smaller baseline shifts) than for the Sb3+-containing oxides.

The drop solution enthalpies of brandholzite were measured in both oxide melt solvents, and the enthalpies of formation from oxides and from elements were calculated via a thermochemical cycle, using reactions in Table 3 and their respective enthalpies (Table 4). The cycle can be expressed by the equation 22ΔfHo(brandholzite)=ΔH7+ΔH3-ΔH4+12ΔH11+ΔH17+ΔH13+12ΔH21.

The enthalpies of formation calculated from the measurements in the two solvents (Table 5) are consistent within their uncertainties.

Low-temperature heat capacity data (Fig. 2a) for brandholzite show no anomalies, as expected for a phase with the elements Mg, Sb, O, and H. Integration with polynomials gave the standard entropy at T=298.15 K of So=571.0±4.0 J mol-1 K-1. DSC data were measured between 280 and 300 K and agree well with the PPMS data. The DSC data are ≈ 0.5 % higher than the PPMS data.

The relatively large uncertainties (in kJ mol-1) of the ΔfHo values from the calorimetric work are related to the high molecular weight of the samples and the large magnitude of heat effects of the primary high-temperature oxide melt solution calorimetric data. The usual uncertainty on these data is about 1 % of the measured signal. In the case of brandholzite, the signal is large because of the appreciable amount of H2O in this phase and the heat effect caused by this H2O. Hence, the uncertainties are not a sign of any problems in calorimetry. Despite these uncertainties, the data are very useful in evaluating the equilibria below.

The drop solution enthalpy of byströmite was measured in the two molten oxide solvents, and the enthalpy of formation from oxides and elements was calculated via a thermochemical cycle, using reactions in Table 3 and their respective enthalpies (Table 4). The cycle can be expressed by the equation 23ΔfHo(byströmite)=ΔH7+ΔH3-ΔH5+ΔH17+ΔH13.

The standard entropy, calculated from the low-temperature heat capacity (Fig. 2b), is 139.3±1.0 J mol-1 K-1. DSC data were measured between 280 and 560 K. In the region of overlap, the DSC data are 1.3 % lower than the PPMS data. The DSC data were shifted to match the PPMS data and then fitted to a Maier–Kelley polynomial. The results are listed in Table 6.

A full-profile refinement of the powder XRD data (Fig. 3a) quantified the fractions of minerals in the sample. This analysis gave 85.2 % chapmanite, 9.9 % quartz, and 4.9 % calcite. Small crystals of pyrite, visible during the separation of the fine fraction, were not captured by this analysis. Inspection of the sample in a scanning electron microscope did not reveal the presence of any other crystalline phases or grains which could be suspected to be amorphous. It has to be noted, however, that the carbonates were also not seen and are suspected to form microcrystalline coatings on the sheet silicates. Chapmanite forms book-like aggregates of platy crystals (Fig. 3b). Tiny (less than a few micrometers) hexahedral crystals used to be pyrite. In the EDX analyses, they gave strong Fe and S but also O signal, suggesting that they are pseudomorphs of iron oxides with sulfate after pyrite. This observation agrees with the fact that pyrite was not captured by PXRD. It is not clear if pyrite weathered naturally or during the initial sample handling. Quartz occurred as larger grains with typical conchoidal fracture.

This sample was subjected to electron microprobe analyses. The analyses were done on flakes of the chapmanite sample obtained by the first fine-fraction separation as polishing of this sample turned out to be impossible. These flakes provide a relatively, but not perfectly, flat surface and are porous, thus diminishing the quality of the analyses. Hence, these results can be considered only as semiquantitative. They are scattered – for example CaO averages at 1.3 wt % with a standard deviation of 0.5 wt % (n=30), and MgO is 1.0 ± 0.6 wt %. These results agree well with the fraction of carbonate refined from the PXRD data. The SO3 signal is ubiquitous, giving 3.7 ± 4.3 wt %. It agrees with the supposition that pyrite in the sample weathered to X-ray amorphous iron oxides (perhaps schwertmannite) with elevated sulfur content. The P2O5 concentration of 2.0 ± 0.4 wt % is attributable to the previous treatment with sodium hexametaphosphate. The only impurity that can be assigned to chapmanite is Al2O3, with 1.4 ± 1.3 wt %.

Obviously, calorimetric results from such a sample would be difficult to interpret. Therefore, the sample underwent several cycles of cleaning. The powder was re-dispersed three times in 5 % HCl overnight and then rinsed, filtered, and dried. Furthermore, a finer fraction of the sample was obtained by re-dispersing the sample, using the same procedure as described in the methods section but allowing it to sediment for 48 h, thus eliminating the larger quartz grains. After this procedure, the sample was rinsed extensively in an attempt to remove the phosphate.

Quartz and carbonates were removed from the sample, as evidenced by the PXRD data after the treatment. We recognized that not all phosphate was removed, and this impurity is not easy to correct for. There remained 0.4 wt % P2O5 and 0.2 wt % SO3 in the sample (recalculated ICP-MS data after the total digestion). Experience with high-temperature oxide melt calorimetry shows, however, that impurities under 1 % can generally be neglected if there is no way to correct for them. This is true if the contribution of such impurities to the measured heat effect is likely to be within the experimental error of the measurement and as long as their heat effects are comparable to that of the major phase. This is the case for P2O5 and SO3. Pyrite, on the other hand, would be an example of an impurity with much different, in this case much higher, signal. Our analyses showed that pyrite was not present, either in the initial or purified sample. The molar Al / (Al + Fe) ratio of 0.06 was considered in the chemical formula used for calorimetric calculations. The molar Sb / (Al + Fe) ratio was 0.98, which is slightly less than 1 but within the uncertainty of the analysis. A slight Sb deficiency in chapmanite was also detected by . They assumed that some of the vacant octahedra were occupied by Fe although there was no evidence of Fe2+ in the sample that could support such an assumption. The method selected (ICP-MS) is considered to be more accurate and precise than the electron microprobe analysis, but the downside is that Si is partially volatilized during the digestion as SiF4. Hence, Si cannot be determined, and the complete stoichiometry cannot be fixed. The formula constructed for the reduction of the calorimetric data is (Fe1.88Al0.12)Sb(Si2O5)O3(OH), with molecular mass of 431.1537 g mol-1. All thermodynamic data presented in this work refer to this formula and molecular mass.

The sample contained adsorbed water that manifested itself strongly during the DSC measurements. During the dynamic DSC measurement, the endothermic signal from the adsorbed H2O (at its peak at T≈620 K) exceeded the intrinsic heat capacity of the sample by ≈ 70 %. To remove the adsorbed H2O prior to the calorimetric experiments, the sample was treated statically (i.e., not under continuously increasing temperature) at 470 K for 3 min and cooled. The loss of structurally bound H2O begins at 620 K and continues to 900 K, when it is interrupted in the data by a mass gain. This mass gain is probably related to oxidation of Sb. The nature of high-temperature products was not investigated. The PXRD data indicated no structural changes in chapmanite after the thermal treatment at 470 K.

In addition, the purification led to preferential enrichment of the sample in smaller particles. Small particle size could also influence thermodynamic properties. It would be desirable, in the future, to repeat the calorimetric experiments on a synthetic, pure sample of chapmanite. However, synthesis protocols for such a phase are unknown at this time. Despite these complexities, we believe that our results are an accurate representation of the thermodynamic properties of this phase.

From the two solvents commonly used for high-temperature oxide melt solution calorimetry see, lead borate is the only option for the calorimetry on chapmanite. This restriction is caused by its ability to dissolve silicates, in contrast to sodium molybdate. For this reason, we have cross-checked the two solvents and established that lead borate is a suitable solvent for calorimetry of antimonous and antimonic compounds.

The drop solution enthalpy of chapmanite was measured in lead borate, and the enthalpy of formation from elements and oxides was calculated via a thermochemical cycle, using reactions in Table 3 and their respective enthalpies (Table 4). The cycle can be expressed by the equation 24ΔfHo(chapmanite)=0.94ΔH8+0.06ΔH9+0.5ΔH1+2ΔH10+0.5ΔH11-ΔH6+0.94ΔH18+0.06ΔH19+0.5ΔH12+2ΔH20+0.5ΔH21.

The enthalpy of formation from oxides, for the reaction 250.94Fe2O3+0.06Al2O3+0.5Sb2O3+2SiO2+0.5H2O→(Fe1.88Al0.12)Sb(Si2O5)O3(OH), is ΔH25o= -22.6±3.7 kJ mol-1.

Heat capacity of chapmanite was measured at sub- and superambient conditions. The low-temperature Cp data show a pronounced but broad anomaly centered at T=10.7 K (Fig. 4). Integration of the low-temperature Cp data gave So=305.1±2.1 J mol-1 K-1. Combination of the enthalpy of formation and entropy gave ΔfGo=-2973.6±4.7 kJ mol-1 and ΔG25o=-34.6±5.6 kJ mol-1 with the corresponding log K25=6.06. The DSC data were shifted to match the PPMS data and fitted by a Maier–Kelley polynomial (Table 6).

Combining Eq. (25) with dissolution reactions for hematite, corundum, valentinite, and quartz defines the dissolution reaction for chapmanite: 26(Fe1.88Al0.12)Sb(Si2O5)O3(OH)+6H+→1.88Fe3++0.12Al3++2SiO20+Sb(OH)30+2H2O, with log K26=-17.10. This value can be used for geochemical modeling with software packages such as PHREEQC or Geochemist's Workbench.

An initial check of the calorimetric data is the comparison of the enthalpies for a reaction that relates all three antimony oxides: 270.5Sb2O3+0.5Sb2O5→Sb2O4.

A cursory examination of the values in Table 7 shows that the enthalpies for Eq. (27) are scattered and inconsistent. Early data from dissociation pressure measurements agree with the newer calorimetric work of . Our data form another cluster of values with the reaction calorimetry of and the critical selection of . The enthalpies for Eq. (27) span a range of 110 kJ mol-1, which makes critical comparison of our and earlier data complicated.

The data agreement is even worse for the reaction 280.5Sb2O3+O2→Sb2O5 as shown in Fig. 5. The values calculated from reported data scatter almost over 150 kJ mol-1 and document the difficulties of measurements on antimony oxides. These compounds are sensitive to heat treatment and oxygen fugacity and are hard to control. This trouble was noted early on, for example by J. J. Berzelius who wrote in 1812 that “I have never worked with a material with which it was so extremely difficult to obtain constant results”, referring to his work on antimony oxides . Obviously, much of this trouble has persisted until today.

When considering the enthalpies of formation of the antimony oxides, the most striking is the scatter for Sb2O5. gave -962.3±8 kJ mol-1, and p. 151 reported -1007.51 kJ mol-1, but presented a much different value of -1128±48 kJ mol-1, to name a few. Recently, determined the enthalpy of the drop solution for metallic Sb. They also derived the enthalpy of formation of Sb2O5 to be -957.0±3.1 kJ mol-1, but we improved this value by using a better characterized Sb2O5 sample. Combining our calorimetric data for Sb2O5 and the data of for Sb, we obtain -953.0±2.4 kJ mol-1. With this ΔfHo value for Sb2O5, our data for the Mg-Sb phases (brandholzite and byströmite) are reasonably close to what would be expected from solubility data (see Table 8).

The data for Sb2O3 are much less scattered, even though not every publication specified if the data relate to valentinite or senarmontite. gave -702.5±8.7 kJ mol-1, p. 151 gave -709.44 kJ mol-1, selected -708.6±2.9 kJ mol-1. A recent EMF study gave -709.74 kJ mol-1 and showed that most available values in the literature cluster tightly between -708 and -710 kJ mol-1.

The available data show extreme discrepancies for Sb2O5. It could be suspected that one of the reasons for the inconsistency is the poorly defined oxidation state in the “Sb2O5” samples. Why otherwise would efforts to produce pure sample involve syntheses under high O2(g) pressure ? It was noted that “antimonic acid cannot be dehydrated by heating in air to give products of constant and reproducible weight without simultaneous reduction of some of the SbV to SbIII” . The nature of the Sb2O5 used in various studies is therefore questionable, but this issue is beyond the scope of this study.

The situation regarding valentinite is much more favorable, and this phase can serve as a good reference compound for the calorimetry on Sb3+ phases. For further work, we adopt the datum from that lies approximately in the middle of the tight cluster of values.

The agreement between our calorimetric data and published solubility data is less than optimal, although differences of similar magnitude are encountered in many systems. The values are summarized for easier comparison in Table 8. Yet, we can circumvent the need for reference compounds in the solution calorimetry and make a direct comparison. Consider the reaction 29MgSb2O6(cr)+12H2O(l)→Mg(Sb(OH)6)2⋅6H2O(cr).

The ΔG29o value, calculated from the solubility data , is 41.4±5.1 kJ mol-1. The calorimetric data from this study give ΔG29o=42.7±27.6 kJ mol-1. The large error relates mostly to the calorimetric datum of brandholzite (Table 4, ±22.0 kJ mol-1) that was already discussed. The agreement between the two ΔG values is excellent. This agreement reinforces the suspicion that the primary source of the differences is the data for Sb2O5, even though the latest ΔfHo value for Sb2O5 our data and yields fairly good results. Hence, a compound like MgSb2O6 would possibly be a better reference compound for solution calorimetry of Sb5+ phases.

Using the solubility data of and , combined with the entropies determined in this work, we propose a set of consistent thermodynamic functions for byströmite and brandholzite. They are listed in Table 9.

In this and previous studies e.g.,, chapmanite was consistently found to be associated with carbonates, mostly calcite or dolomite. Sulfides in the association of chapmanite were noted, for example small pyrite crystals (this work) or marcasite Bräunsdorf, Germany,. In the Keeley mine (Ontario, Canada), chapmanite is intergrown with native silver . In Pezinok (Slovakia), chapmanite was found to be associated with secondary siderite, cryptocrystalline SiO2 (hyalite), and allophane-like material . In both Pezinok and La Bessade (France) , chapmanite was identified in the deeper parts of the veins, partially forming pseudomorphs after stibnite. At La Bessade, chapmanite was intergrown with opaline silica, forming compact aggregates with conchoidal fracture. The association with quartz or chalcedony was described as a characteristic feature by .

These assemblages bear information about the formation conditions of chapmanite. The common association with carbonates suggests that chapmanite cannot be related to acidic solutions related to oxidative weathering of pyrite. The presence of pyrite, marcasite, or siderite indicates reducing conditions. The carbonates, for example calcite or dolomite detected in this work or by , could form later than chapmanite, thus invalidating the assumption that acidic conditions are excluded. On the other hand, siderite found with chapmanite at Pezinok seems to be co-genetic and indicates indeed neutral or alkaline conditions. Chapmanite from Argent mine South Africa, was reported to occur on siderite, but no details about this association were provided. Opaline silica, reported from several localities, suggests that the fluids were transporting greater amounts of silicic acid. Hence, chapmanite could be a product of supergene processes, as suggested by , but in deeper, reduced portions of the ore bodies.

The thermodynamic data presented here allow one to model the stability of chapmanite quantitatively. There are two insoluble, competing sinks for antimony in low-temperature settings: schafarzikite and tripuhyite . The question is therefore the relationship between these phases. Figure 6 presents the results of a forward simulation with PHREEQC. In this model, a Ca-Mg-SO4 solution (for its composition, see supporting electronic information) with pH = 5.7 was equilibrated with chapmanite under different redox conditions. They were imposed on the solution by simply varying the input value of pϵ. The starting solution had no Fe, Sb, Al, or Si. After equilibration, pH increased slightly but not more than 0.2 log units. The model shows that under low-pϵ conditions the saturation index of schafarzikite is almost always higher than 0, and this phase could precipitate. Under oxidizing conditions, tripuhyite is predicted to precipitate in a similar manner. There is not even a small window of pϵ conditions under which both competing phases would be undersaturated. It must be pointed out that tripuhyite, and likely also schafarzikite, possess dense structures whose formation from low-temperature fluids is kinetically inhibited . Therefore, rapid transformation of chapmanite to schafarzikite or tripuhyite should not be expected.

What are then the conditions conducive for the formation of chapmanite? Its association with opaline silica suggests elevated activity of SiO2(aq), which would stabilize and favor chapmanite over schafarzikite or tripuhyite, at least under some redox conditions.

It is interesting to examine mine drainage solutions in terms of the saturation state with respect to chapmanite. From a database of mine drainage solutions that we amassed over years, we selected the analyses where Eh was analytically determined (by oxidation–reduction potential measurement in the field or Fe2+ / Fe3+ determination in the laboratory). The results (Fig. 7) show that most of the examined solutions are supersaturated with respect to chapmanite. This observation is in conflict with the reports that chapmanite is such a rare mineral. Yet, unless specifically looked for, chapmanite would escape attention because it would fall into the group of 1 : 1 sheet silicates (like kaolinite) that are commonly found but not often investigated in detail.

Precipitation of chapmanite is likely also kinetically inhibited, as for many sheet silicates at low temperatures. Formation of Fe-rich sheet silicates at low temperatures is possible through aging of hydroxide-silica amorphous precursors e.g.,, adding to the group of minerals that form from gel-like precursors see. This notion fits very well with the observation of amorphous silica in association with chapmanite.