Densities and heat capacities of aqueous arsenious and arsenic acid solutions to 350 °C and 300 bar, and revised thermodynamic properties of As(OH) 3 ∘ ( aq ) , AsO(OH) 3 ∘ ( aq ) and iron sulfarsenide minerals

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Densities and heat capacities of aqueous arsenious and arsenic acid solutions to 350 °C and 300 bar, and revised thermodynamic properties of As(OH) 3 ∘ ( aq ) , AsO(OH) 3 ∘ ( aq ) and iron sulfarsenide minerals

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  Densities and heat capacities of aqueous arsenious andarsenic acid solutions to 350   C and 300 bar, and revisedthermodynamic properties of As ð OH Þ 3  ð aq Þ ,AsO ð OH Þ 3  ð aq Þ  and iron sulfarsenide minerals Erwan Perfetti  a,1 , Gleb S. Pokrovski  b,* , Karine Ballerat-Busserolles  a ,Vladimir Majer  a , Franc¸ois Gibert  c a Laboratoire de Thermodynamique des Solutions et des Polyme` res, UMR 6003, Universite´  Blaise Pascal Clermont-Ferrand, CNRS,Avenue des Landais, F-63177 Aubie` re, France b Laboratoire des Me´ canismes et Transferts en Ge´ ologie, UMR 5563, LMTG—Universite´  de Toulouse—CNRS—IRD—OMP, 14 Avenue Edouard Belin, F-31400 Toulouse, France c Laboratoire Magmas et Volcans, UMR 6524, Universite´  Blaise Pascal, Rue Kessler, F-63000 Clermont-Ferrand, France Received 23 May 2007; accepted in revised form 13 November 2007; available online 26 December 2007 Abstract Densities and heat capacities of aqueous arsenious and arsenic acid solutions of 0.1–0.6 mol/kg were measured using theflow vibrating tube densitometry and differential calorimetry at temperatures to 350   C and pressures to   310 bar. The stan-dard partial molal volumes  V   and heat capacities  C    p   of the neutral aqueous As III and As V (oxy)hydroxide species, As(OH) 3 and AsO(OH) 3 , were obtained from these data, via corrections for partial dissociation and extrapolation to infinite dilution.The generated  V   and  C    p   values, together with the existing data on As III oxide and sulfide minerals solubilities and low-tem-perature As III  –As V aqueous solution equilibria, were used to refine the thermodynamic properties of As hydroxide complexesover a wide temperature–pressure range, in the framework of the revised HKF equation of state and using correlation algo-rithms recently proposed for aqueous neutral species. These revised properties were combined with solubility data for arse-nopyrite (FeAsS) and direct calorimetric heat capacity and enthalpy measurements reported in the literature for arsenopyrite,loellingite (FeAs 2 ), and westerveldite (FeAs), to generate a consistent set of thermodynamic parameters for these iron sulfar-senides. The new Gibbs free energy values of arsenopyrite and loellingite resulting from these properties imply lower solubil-ities of iron sulfarsenides in aquatic environments than have been assumed. The thermodynamic properties of arsenic aqueousspecies and solid phases obtained in this study provide quantitative constraints on As-bearing mineral stabilities and arsenictransport by geological fluids.   2007 Elsevier Ltd. All rights reserved. 1. INTRODUCTION Arsenic is one of the most popular trace elements in geo-chemistry because of its ubiquity in a variety of low- andhigh-temperature natural and industrial environments suchas hydrothermal ore deposits, geothermal fields, mine andmineral processing wastes, and many sedimentary rocksand aquifers (e.g., Vaughan, 2006). Accurate knowledgeof the thermodynamic properties (i.e., stabilities and solu- 0016-7037/$ - see front matter    2007 Elsevier Ltd. All rights reserved.doi:10.1016/j.gca.2007.11.017 * Corresponding author. Fax: +33 0 5 61 33 25 60. E-mail address:  pokrovsk@lmtg.obs-mip.fr (G.S. Pokrovski). 1 Present address: Institut Franc¸ais du Pe´trole, Division Ge´olo-gie-Ge´ochimie-Ge´ophysique, 1-4 avenue de Bois-Pre´au, 92852Rueil-Malmaison Cedex, France. www.elsevier.com/locate/gca  Available online at www.sciencedirect.com Geochimica et Cosmochimica Acta 72 (2008) 713–731  bilities) of As-bearing minerals and aqueous species is a pri-mary step to resolving many issues of the hydrothermal andenvironmental geochemistry of arsenic. The aim of thisstudy is to provide such data for the main As aqueous com-plexes, arsenious and arsenic acids, on the basis of in situcalorimetric and volumetric measurements over a wide tem-perature–pressure ( T   –  P  ) range, and to use these results togenerate a revised set of thermodynamic properties for ironsulfarsenide minerals.Arsenic is an important constituent of hydrothermal-volcanic ore deposits hosting a range of sulfide and arsenideminerals (e.g., orpiment As 2 S 3 , realgar AsS), sulfosalts (e.g.,enargite Cu 3 AsS 4 , proustite Ag 3 AsS 3 ), and Fe–Co–Ni(sulf)arsenides (e.g., arsenian pyrite Fe(As,S) 2 , arsenopyriteFeAsS, cobaltite CoAsS, gersdorffite NiAsS, loellingiteFeAs 2 , westerveldite FeAs, nickeline NiAs, saffloriteCoAs 2 , and their solid solutions). Among these solids, arse-nopyrite and loellingite are common As-bearing end-mem-bers over a wide range of conditions from high temperaturemagmatic–hydrothermal porphyry-style and skarn Sn–Wand Cu(±Au) systems to mesothermal polymetallic Cu– Pb–Zn–Ni–Co–Ag and Au deposits (e.g., Scott, 1983;Sharp et al., 1985; Heinrich and Eadington, 1986; Kovalen-ker et al., 1995; Kerr et al., 1999; Hem et al., 2001; and ref-erences therein). Of particular interest is the association of As with platinum group elements (PGE) like Pt, Ir, Os,which are capable of both entering the major (Fe–Co–Ni)arsenides and forming their own arsenide minerals (e.g.,sperrylite, PtAs 2 ) in Co–Ni–Fe–Cr–PGE magmatic–hydro-thermal deposits (e.g., Gervilla et al., 1998; Gervilla andKojonen, 2002; Barkov and Fleet, 2004; Barkov et al.,2004). The physical–chemical mechanisms responsible forthe close association of As with precious metals like Au,Ag, and PGE in their ores are, however, poorly quantifiedpartly because of the lack of accurate data on the formationconditions and solubilities of sulfarsenides minerals in theore-forming fluids. As a first step, thermodynamic proper-ties of sulfarsenides end-members are required to predictthe stability and energetics of their numerous solid solu-tions and to model trace element incorporation into theseminerals. The understanding of the oxidation and weather-ing kinetics of As-bearing primary minerals in mine tailingsand during ore treatment also requires knowledge of FeAsSand FeAs 2  stabilities at ambient temperatures (Craw et al.,2003). Thermodynamic information on arsenopyrite andother iron arsenides (FeAs, FeAs 2 ) is, however, very scarce.It is largely based on phase equilibria measurements in thedry system Fe–As–S at temperatures above 500   C (Clark,1960; Morimoto and Clark, 1961; Barton, 1969; Kretsch-mar and Scott, 1976; Scott, 1983). The standard thermody-namic properties of Fe sulfarsenides derived there frommay exhibit large uncertainties because of the slow equili-bration kinetics in a dry sulfide–arsenide system, and theinaccuracy of calculated or measured fugacities of sulfurand arsenic gaseous species. Moreover, the extrapolationof these high-temperature properties to 25   C may be af-fected by significant uncertainties. As a result, the thermo-dynamic properties of arsenopyrite and loellingite reportedin the literature are inconsistent and controversial. Forexample, the widely used values of standard enthalpy andentropy of FeAs 2  at 25   C and 1 bar derived from high-tem-perature phase equilibria (Barton, 1969),  D f  H   =   43.5kJ/mol and  S   = 127.2 J/(mol K), are significantly differentfrom more recent calorimetric heat-capacity and heat-of-formation measurements (Stolyarova, 1977; Pashinkinet al., 1991),  D f  H   =   85.8 kJ/mol and  S   = 80.1 J/(mol K).Similar discrepancies exist for arsenopyrite whose Gibbsfree energy of formation at 25   C and 1 bar proposed byBarton (1969),  D f  G   =   109.6 kJ/mol, is   30 kJ/mol morepositive than the corresponding value derived recently fromFeAsS solubility in aqueous solution at temperatures 300– 450   C by Pokrovski et al. (2002a),  D f  G   =   141.7 kJ/mol.The latter value is based on the thermodynamic proper-ties of the As III hydroxide species, arsenious acid As(OH) 3 ,which is the main As-bearing complex in most acid-to-neu-tral hydrothermal fluids (Pokrovski et al., 1996, 2002a,b;and references therein). These properties stem from the sol-ubilities of As 2 O 3  and As 2 S 3  and are limited to 300   C(Pokrovski et al., 1996). Consequently, the interpretationof As–mineral solubilities and stabilities at higher tempera-tures may be subject to substantial uncertainties. Very lim-ited experimental data (to 200   C/ P  sat , Pokrovski, 1996;and to 300   C/ P  sat , Zakaznova-Herzog et al., 2006) andtheoretical predictions (Smith et al., 1986; Shock et al.,1997) are available for the de-protonated form of arseniousacid, AsO ð OH Þ 2  . Although the discrepancies betweenthese sources attain up to 1.5 log unit for the dissociationconstant at 300   C, all data indicate that AsO ð OH Þ 2  may appear in substantial amounts only in alkaline solu-tions (pH P  7–9). In surface waters saturated with atmo-spheric oxygen, As V hydroxide species, arsenic acid,H 3 AsO 4  or AsO(OH) 3 , 2 and its deprotonated counterparts(AsO 2 ð OH Þ 2  , AsO 3 (OH) 2  , and AsO 43  ) are more stablethan their As III analogs. Although their dissociation con-stants are well known to   100   C (e.g., Baes and Mesmer,1976; Smith et al., 1998), the thermodynamic data for aque-ous arsenic acid itself are poorly constrained even at ambi-ent conditions, and its super ambient properties are basedexclusively on predictions from ambient-temperature Gibbsfree energy and enthalpy (e.g., Shock et al., 1997). Despitethe common belief that pentavalent As species are not sta-ble at elevated temperatures in natural systems (e.g., Ball-antine and Moore, 1988), recent discoveries of As V incorporated in serpentinites, formed at depths of   100 km and temperatures of    600   C in subduction zones(Hattori et al., 2005), likely imply a far larger stability of arsenate species in high  T   –  P   fluids than generally assumed.Thus, there is a need for accurate thermodynamic proper-ties for both As III and As V aqueous complexes at elevatedtemperatures. The distribution of As III and As V hydroxidespecies as a function of pH for a typical range of hydro-thermal temperatures calculated using the available litera-ture data is shown in Fig. 1. 2 In this study, we adopt structural formulas, As(OH) 3  andAsO(OH) 3 , for arsenious and arsenic acids, respectively, and theirdeprotonated forms; they are based on XAFS and Ramanspectroscopic data showing that As binds to OH and O groupsin these species.714 E. Perfetti et al. / Geochimica et Cosmochimica Acta 72 (2008) 713–731  The experimental determination of the derivative prop-erties such as partial molal volumes ( V  ) and heat capacities( C   p ) of soluble aqueous species is a convenient way forobtaining the Gibbs free energies at high  T   –  P   by integra-tion using the data at 25   C and 1 bar as integration con-stants. Recent advances in designing vibrating-tubedensimeters and flow-through calorimeters allow acquisi-tion of   C   p  and  V   data for a variety of relatively soluble(>0.2 m/kg H 2 O) aqueous species to temperatures of 450   C and pressures of 300 bar (for representative exam-ples see Majer et al., 1991, 1999, 2000; Majer and Wood,1994; Hnedkovsky et al., 1995, 1996; Obsil et al., 1996;Hnedkovsky and Wood, 1997; Obsil, 1997; Tremaineet al., 1997; Xiao et al., 1997; Degrange, 1998; Clarke andTremaine, 1999; Clarke et al., 2000; Xie et al., 2004; Bule-mela and Tremaine, 2005; Censky et al., 2005a,b; Ball-erat-Busserolles et al., 2007; Slavik et al., 2007). Thesenew data greatly improved the existing physical chemicalmodels for predicting thermodynamic properties of solutesat high  T   –  P   (e.g., Akinfiev, 1997; Tremaine et al., 1997;Clarke et al., 2000; Plyasunov et al., 2000a,b; Sedlbaueret al., 2000, 2002; Sedlbauer and Majer, 2000; Yezdimeret al., 2000; Plyasunov and Shock, 2001a; Majer et al.,2004; Censky et al., 2007). Up to now, there have beenno such measurements for As-bearing species at superambi-ent conditions.The present study was initiated to better characterize thethermodynamic properties of aqueous arsenious and ar-senic acids and iron sulfarsenide minerals at hydrothermalconditions. For this purpose, the standard partial molalvolumes  V   and heat capacities  C    p   of As(OH) 3  andAsO(OH) 3  were determined from calorimetric and volumet-ric measurements up to 350   C and 310 bar. The experimen-tal values were used to obtain, by extrapolation to infinitedilution and correction for partial ionization, the standardmolal volumes and heat capacities from the apparent molalproperties  Y  exp U  defined as Y  exp U  ¼ ð Y     55 : 5   Y  w Þ = m Y   ¼  V    ; C   p   ð 1 Þ where  Y   relates to a 1 kg H 2 O solution,  m  is molality of thespecies, and  Y  w  is a molar property (volume or heat capac-ity) of pure water. These standard properties determined asa function of temperature and pressure and combined withambient-temperature thermodynamic values allow an im-proved prediction of the thermodynamic functions of aque-ous arsenic and arsenious acids within the framework of therevised HKF equation of state (Helgeson et al., 1981; Tan-ger and Helgeson, 1988; Shock et al., 1989; Plyasunov andShock, 2001a). These functions, together with recent solu-bility measurements of FeAsS (Pokrovski et al., 2002a)and literature heat capacity and enthalpy data for FeAsS,FeAs 2  and FeAs (Stolyarova, 1977; Gonzalez-Alvarezet al., 1989; Pashinkin et al., 1989, 1991), allow generationof a consistent set of thermodynamic parameters for ironsulfarsenides. These new data, combined with availablegeological information on mineral parageneses of As-richore deposits and As contents measured in fluid inclusions,are used to model equilibria among As-bearing minerals pH 020406080100    %   o   f   t  o   t  a   l   A  s   (   I   I   I   )  As(OH) 3 ° AsO(OH) 2- 350°C, 300 bar  21012 pH 020406080100    %   o   f   t  o   t  a   l   A  s   (   I   I   I   )  As(OH) 3 ° AsO(OH) 2- 25°C, 1 bar  21012 pH 020406080100    %   o   f   t  o   t  a   l   A  s   (   V   )  AsO(OH) 3 ° AsO 2 (OH) 2 - AsO 3 (OH) 2-  AsO 43- 25°C, 1 bar  pH 020406080100    %   o   f   t  o   t  a   l   A  s   (   V   )  AsO(OH) 3 ° AsO 2 (OH) 2 - AsO 3 (OH) 2-  AsO 43- 350°C, 300 bar  4682101246821012468468 Fig. 1. Distribution of As III (left) and As V (right) hydroxide species as a function of pH at 25   C, 1 bar and 350   C, 300 bar and an ionicstrength of 0.01. The curves were calculated using arsenious and arsenic acids dissociation constants according to Shock et al. (1997). Thevertical dashed lines stand for the pH of the neutrality point of water at given temperature and pressure. The pH ranges of the volumetric andcalorimetric measurements of this study are 4.5–5.5 for As III at all temperatures and pressures; 1.5–2.0 and 2.5–3.0 for As V at the lowest(25   C) and highest (350   C) temperature, respectively.Thermodynamic properties of As hydroxide species and Fe sulfarsenides 715  and the transport and precipitation of arsenic by mag-matic–hydrothermal fluids. 2. MATERIALS AND METHODS2.1. Sample preparation and analyses Arsenious acid aqueous solutions for calorimetric andvolumetric measurements were prepared by dissolution of arsenic trioxide powder (As 2 O 3 , 99.95%, Aldrich) in doublydeionized and degassed water (MilliQ, 18 M X ) under argonatmosphere at 80   C. The solid was dried at 60   C overnightand carefully weighed (±0.1 mg) prior to dissolution to ob-tain about 500 mL of    0.3 m As undersaturated solution(As 2 O 3  solubility at 80   C is   0.6 m As, Pokrovski et al.,1996). Complete dissolution is achieved within 2 days.The obtained stock solution was filtered through a0.45 l m Millipore filter, diluted to obtain a series of con-centrations (from 0.1 to 0.3 m As), and stored under argonin tight Pyrex bottles at 25 ± 2   C. All 0.1–0.3 m solutionswere found to be stable within at least 1–2 months. TotalAs and As III concentrations were checked by the colorimet-ric molybdate blue method (Pokrovski, 1996), flame atomicabsorption, AAS (Pokrovski et al., 2002a) and titrationwith iodine following standard procedures (Charlot, 1966;Perfetti, 2003). Neither colorimetry nor titration detectedAs V in experimental solutions within ±1%. Arsenic concen-trations analyzed by the three methods were always thesame and corresponded, within ±2% of the value, to thosederived by weight of As 2 O 3 . Because the As concentrationscalculated from weights of pure stoichiometric As 2 O 3  andwater are more accurate (uncertainty is less than 0.2%) thanthe wet chemical analyses, we used the weight-derived val-ues in the  C   p  and  V   derivation from measured bulk solutionheat capacities and densities.Arsenic acid aqueous solutions were prepared by disso-lution of hydrated arsenic penta-oxide (As 2 O 5 Æ n H 2 O, n    2–5, Aldrich) in degassed MilliQ water at ambient tem-perature. The solubility of hydrated As 2 O 5  is extremelyhigh (>3 mol/kg As, Linke, 1958), and complete dissolutionwas achieved within 10 min for a   0.6 m solution. The ob-tained solution was filtered through a 0.45 l m Millipore fil-ter, diluted to prepare a series of concentrations (0.1–0.5 mAs) and stored as described above for As III . Because of theuncertain stoichiometry of the solid, As concentrationcould not be calculated reliably from weights. Conse-quently, total As and As V concentrations were determined,respectively, by AAS and iodometric titration (Charlot,1966; Perfetti, 2003). The As V titrimetric analysis is basedon the reduction of iodide (added as excess of KI) to iodine( I  2 ) via the reaction with arsenic acid in acid media (15 wt%HCl) followed by titration of the generated I 2  with a stan-dard sodium thiosulfate (Na 2 S 2 O 3 ) solution. The reproduc-ibility of the analyses is about 1% of the concentrationvalue. Atomic absorption spectroscopy and titrimetry anal-yses of the same solutions were identical within errors (±1– 2%) indicating that no As III was present in experimentalsolutions. Thus, As V concentrations used in  C   p  and  V   mea-surements exhibit uncertainties of 1–2% of the value for thecorresponding molal property  Y  exp U  . 2.2. Density and heat-capacity measurements and datatreatment  2.2.1. Volumetric data Values of apparent molal volumes  V    exp /  were calculatedfrom the experimental densities as: V    exp /  ¼  M  s q w  þ D q   D q m ð q w  þ D q Þ q w ;  ð 2 Þ where  M  s  is the molar mass of the solute and  m  is molalityof the solution;  q w  and  D q  are the density of water and themeasured density difference between the experimental solu-tion and pure water, respectively. The density differences, D q , were obtained at seven temperatures from 25 to350   C and at one to three pressures (ranging from values10–20 bar above the saturated water vapor pressure ( P  sat )to 300–310 bars), using a vibrating-tube densitometer de-scribed in detail by Hynek et al. (1997). Briefly, the instru-ment operates with a thin platinum–rhodium tubeoscillating in a field of two permanent magnets at a reso-nance frequency near 133 Hz. The temperature was mea-sured with a customized 500 X  secondary standardresistance thermometer from Burns Engineering with anaccuracy of 0.05   C. The stability of the temperature is en-sured by a system of superposed thermally regulated jacketsand lids surrounding the block housing the vibrating tube.The density measurements were performed in a flow regimeat a constant flow rate near   0.01 cm 3 /s. The pressure wascontrolled using a Circle Seal back-pressure regulator at theend of the flow system and measured by a DPI 260 Druckelectronic pressure gauge with an accuracy of 0.3 bar. A six-port valve connected with a sample loop allowed alternativeintroduction of water and pressurized solution samples inthe flow system.The densities were also measured with a commercial So-dev–Picker vibrating tube densitometer type 03D (Pickeret al., 1974) at 25   C and atmospheric pressure where thetemperature was measured with an accuracy of 0.01   C.The results were consistent with those obtained on thehigh-temperature instrument and were treated together.The differences  D q  between densities of an experimentalsolution  q  and pure water  q w  were calculated as D q  ¼  q   q w  ¼  K  ð s 2   s 2w Þ ;  ð 3 Þ comparing the periods of oscillation  s  and  s w  of the vibrat-ing tube filled with the experimental solution and purewater, respectively. The calibration constant  K   was deter-mined at each experimental temperature and pressure bymeasurements with a 3 molal NaCl solution up to 150   Cand with nitrogen gas at higher temperatures where theNaCl data are not of sufficient accuracy. Water densitywas calculated from the equation of  Hill (1990), and thedensity of the NaCl solution was obtained from the recom-mended values of  Archer (1992) that were determined onthe basis of evaluating the major literature sources. The rel-ative error in  K   for the high-temperature densitometer wasestimated as 0.2–0.4% of the value based on the reproduc-ibility of the calibration experiments and accuracy of thedensities of the fluids used in calibration. The Sodev–Pickerdensitometer was calibrated by comparing vibration peri- 716 E. Perfetti et al. / Geochimica et Cosmochimica Acta 72 (2008) 713–731  ods corresponding to water and vacuum in the tube and thecalibration constant is expected to be known with an accu-racy of 0.1%.Multiple volumetric measurements for several concen-trations (0.1–0.3 m for arsenious acid) and (0.1–0.5 m forarsenic acid) led to the determination of 81 and 77 new datapoints for  V    exp /  of As(OH) 3  and AsO(OH) 3 , respectively.For all  T   –  P   conditions,  V    exp /  is typically slightly changingwith concentration without any clear systematic trend (seebelow). The error in  V    exp /  can be calculated as propagationof the uncertainty in the density difference  D q  (Eq. (3)) thatwas obtained as a statistical estimate taking into accountthe fluctuations in  s  and  s w  and the systematic errors in  K  and the sample concentration. The raw experimental den-sity differences and the corresponding apparent molal vol-umes are available integrally in electronic annex.  2.2.2. Heat-capacity data Values of apparent molal heat capacities  C  exp  p  ; /  were cal-culated from the experimental data as: C  exp  p  ; /  ¼  M  s c  p   þ ð c  p     c  p  ; w Þ = m ¼  c  p  ; w fð  M  s  þ  1 = m Þ c  p  = ð c  p  ; w    1 = m Þg ;  ð 4 Þ where  c  p  and  c  p ,w  are the specific heat capacities (per unit of mass) of solution and water. The specific heat-capacityratio  c  p / c  p ,w  was determined using a high-temperaturehigh-pressure Picker-type differential flow calorimeter attemperatures from 50 to 350   C and pressures to 310 bars.The instrument and measurement procedure have beendescribed in detail by Hnedkovsky et al. (2002), and onlythe most salient features are given here.The core of the calorimeter is a temperature-controlledmassive block housing two identical cells consisting of a2-mm outer-diameter platinum–iridium tubing equippedwith a wire heater wound around the tube and with resis-tance temperature detectors indicating the temperature riseof the heated fluid compared to the block. The ratio of spe-cific heat capacities  c  p ,s / c  p ,w  is related to the heating powersin the measuring cell as follows (Hnedkovsky et al., 2002) c  p  = c  p  ; w  ¼ ð q w = q Þ SL f 1 þ  f  ð  P     P  w Þ =  P  w g ;  ð 5 Þ where ( q w / q ) SL  is the ratio of water and solution densities atthe temperature of the sample loop (  25   C) and experi-mental pressure, and  f   is the correction factor for heatlosses. The electric powers  P   and  P  w  keep the identical tem-perature rise  D T  exp   2   C when either the sample solutionor water are flowing through the measuring cell. A secondcell connected in series with the measuring cell is used tofacilitate equilibration of the temperature bridge and tocompensate for the flow rate and temperature fluctuationsin the calorimeter. The heat loss correction factor  f   wasdetermined by changing the water flow rate  F  w  in the sam-ple cell, to mimic a change in heat capacity, and by measur-ing the corresponding change in the heating input. Theresulting values of   f   = D FP  w /( D PF  w ) typically range from1.02 (at 50   C) to 1.10 (at 350   C).Considering the small temperature change during the  C   p measurements ( D T  exp   2   C, see above), the temperaturereported for calorimetric measurements is adopted as theaverage value before and after the cell heating( T   =  T  block  + 0.5 Æ D T  exp ). The stability of the block temper-ature is ensured by a system of heated concentric jacketsand lids that also serve for pre-heating fluids entering theblock. A 500 X  secondary standard resistance thermometerof Burns, calibrated by manufacturer against the interna-tional primary temperature standard IPTS 90, was usedfor measuring the block temperature. Considering the pos-sibility of thermal gradients at elevated temperatures( P 200   C), the uncertainty in temperature determinationis estimated to be 0.2   C near the upper temperature limitof the instrument (  400   C).A sample loop with an internal volume of approximately20 cm 3 is placed in a water bath thermostat, filled with theexperimental solution, and pressurized to the experimentalpressure. The sample is then injected into the calorimeterby connecting the loaded sample loop into the stream of water, between the reference and measuring cells. A high-pressure pump working at a constant volumetric flow ratewas used to supply the base flow ( F  w    1.8 cm 3 /min) tothe calorimeter. The pressure was maintained constant bya back-pressure regulator placed at the end of the streamsimilarly as for the densitometer. The pressure in the systemwas measured with a Bourdon-type pressure gauge with anaccuracy of 0.1 MPa.The apparent molal heat capacities of As(OH) 3  andAsO(OH) 3  solutions were determined for 6 target tempera-tures between 50 and 350   C and at one or two pressures,one slightly above the saturated water vapor pressure andanother one near 29 MPa (  290 bar). For each temperatureand pressure, measurements were performed at least atthree different concentrations (0.1–0.3 m) and (0.1–0.6 m)leading to 74 and 72  C  exp  p  ; /  data points for arsenious and ar-senic acid, respectively. The error in the apparent molalheat capacity can be calculated as propagation of theuncertainty in the heat-capacity ratio, that is obtained asa statistical estimate taking into account instabilities of   P  and  P  w  as well as uncertainties in ( q ) SL  and in the  f   factor(see Eq. (5)). In addition, a possible systematic error dueto the uncertainty in the sample concentration must alsobe taken in consideration. A detailed description of theerror estimation can be found in Slavik et al. (2007). Theexperimental heat-capacity ratios, and the resulting appar-ent molal heat capacities are reported integrally inelectronic annex.  2.2.3. Data treatment The obtained apparent partial molal heat-capacity andvolume values (Eqs. (2) and (4)) were corrected for ioniza-tion and relaxation effects (Ballerat-Busserolles et al., 2007)which are related to changes of   V   or  C   p  and enthalpy of theionization reaction of arsenic acid in experimentalsolutions. While arsenious acid can be considered as a non-electrolyte with a negligible degree of ionization in acid-to-neutral aqueous solution at all conditions of ourexperiments (p K  dis   9.3 and   9.0 at  P  sat  and 25 and350   C, respectively, Shock et al., 1997), the arsenic acidis a weak electrolyte undergoing partial ionization in aque-ous solution that diminishes with increasing temperature(p K  dis1   2.3 and   5.2 at  P  sat  and 25 and 350   C, respectively, Thermodynamic properties of As hydroxide species and Fe sulfarsenides 717
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