haway · 2003. 10. 9. · sodium salts). therefore, under hts conditions, precipitation of any...
TRANSCRIPT
AFCL-12064
1JO
1999 December
2Analytical Chemistry BranchChalk River Laboratories
Chalk River, Ontario, KOJ
HaWay
‘Reactor Chemistry Branch
Lemire’, Nancy B. Tosello’ and James D.
bY
Robert J.
ANTIMONY(III) AND ANTIMONY(V) SOLIDSIN BASIC AQUEOUS SOLUTIONS TO 300°C
AECL
SOLUBILITY BEHAVIOUR OF
ABCL-12064
Dt?cembre 1999
1JO
2 Chimie analytiqueLaboratoires de Chalk River
Chalk River (Ontario) KOJ
reacteursChimie des ’
forme depyrochlore) puisse Ctre moins soluble dans des solutions presque neutres, de faible force ionique.
Sbz05 hydrate (en particulier la l’hypothese que le Bcarter
ajot@. Par consequent,dans les conditions du circuit primaire, la precipitation d’oxydes d’antimoine ou d’oxydes mixtes estpeu probable. On ne peut pas
soit se1 de sodium ne moldm” dans n’importe quelle solution
aqueuse neutre ou basique (en supposant qu’aucun 0,00005 2 totales d’antimoine en solution
utilises dans ces experiences produisent desconcentrations
solides tous les a ce que
solides contenant del’antimoine dans les solutions oxydantes basiques depend fortement des cations et de leurconcentration en phase aqueuse.
On pourrait s’attendre
mat&es m&me si la composition des Sb(OH);), (SbOj ou esp&ces anioniques de la solution
d’antimoine stabilite des
solidesd’antimoine(V) correspondent aux variations de la
mat&es a 250°C. Ces variations de solubilite des sup&ieures decroit aux temperatures 2OOOC etB croit de 25 protone solubilite de cet antimoniate de sodium partiellement
presente une structure depyrochlore. La
Na&H(H20)]2_&b206, qui mat&e solide, concerne la Btaient instables dans les solutions d’hydroxyde de
sodium en ce qui SbzOs.xHzO et l’antimoniate de sodium simple
trouve que le250°C, dans les solutions oxydantes, on a legerement. A decroit 200°C,
puis se stabilise ou Sb203 augmente d’environ deux ordres de grandeur entre 25 et
reduite auminimum, la solubilid du
(Ill) en antimoine (V) est
solides d’antomoine enfonction de la temperature.
Dans les solutions dans lesquelles l’oxydation de l’antimoine
mat&es d&ermination des variations de la solubilite des esp&ces d’antimoine en solution et servent de
guide dans la stabi1it.e en fonction de la temperature des
result&s fournissent des renseignements sur lacharge et la
3OOOC. Les B allant de 25 a des temperatures d’antimoine(IIl) et (V) dans des solutions
basiques solubilite des sels et des oxydes mesure la
reduire au minimum la liberation et la redeposition de cesisotopes, on a
a l’arr& du reacteur.
Dans le cadre d’un programme visant l’entree d’oxygene lors de a associes Cte D ont
l’int&ieur du circuitprimaire d’un reacteur CANDU
a d’activite ‘%Sb dans le transport 122Sb et isotorsLe role et l’importance des
Halliday’
Resume
Tosello’ et James D. Lemire’, Nancy B.
3oO°C
Par
Robert J.
JUSQU’A D’ANTIMOINE(V)
DANS LES SOLUTIONS AQUEUSES BASIQUES
SOLIDESD’ANTIMOINE(III) ET
MAT&ES SOLUBILII% DES
EACL
COMPORTEMENT DE LA
ABCL- 12064
1JO
1999 December
2Analytical Chemistry BranchChalk River Laboratories
Chalk River, Ontario, KOJ
Sb205 (especially the pyrochlore form)might be less soluble in near-neutral, low-ionic-strength solutions.
‘Reactor Chemistry Branch
moldrn-3 in any neutral or basic aqueous solutions (assuming no addedsodium salts). Therefore, under HTS conditions, precipitation of any antimony oxides or mixedoxides is unlikely. It cannot be ruled out that hydrated
2 0.00005
Sb(OH)$, even though the compositions ofantimony-containing solids in basic oxidizing solutions are strongly dependent on the cations andtheir aqueous phase concentrations.
All solids used in the present experiments would be expected to generate total solution antimonyconcentrations
(SbOj or
200°C and decreases at temperatures above250°C. These solubility changes for the antimony (V) solids reflect changes in the stability of theanionic antimony solution species
Na2,[H(H20)]2_2,Sb206, which has a pyrochlore structure. The solubility of this partiallyprotonated sodium antimonate increases from 25 to
Sb2Os.xHzO and simple sodiumantimonate(V) were found to be unstable in sodium hydroxide solutions with respect to the solid,
25O”C, in oxidizing solutions, 2OO”C, and then levels out or
decreases slightly. At Sb203 increases by about two orders of magnitude between 25 and
antimony(lll) to antimony(V) is minimized, the solubility ofln solutions in which oxidation of
antimony(lll) and (V) oxides and salts have been measured in basicsolutions at temperatures from 25 to 300°C. The results provide information on the charge andthe stability as a function of temperature of antimony solution species and, hence, a guide to thetrends in the temperature dependence of the solubilities of antimony solids.
CANDU@reactor primary heat transport system (HTS), have been associated with oxygen ingress duringreactor shutdown. As part of a program to minimize the release and redeposition of theseisotopes, the solubilities of
‘“Sb to activity transport in a 122Sb and
Halliday2
Abstract
The major contributions of the isotopes
ANTIMONY(II1) AND ANTIMONY(V) SOLIDSIN BASIC AQUEOUS SOLUTIONS TO 300°C
Robert J. Len-rim’, Nancy B. Tosello’ and James D.
ABCL
SOLUBILITY BEHAVIOUR OF
RBFBRBNCES . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31
Na2,[H(H20)]2_2aSb206.H20 from 25 to300°C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28
5. CONCLUSIONS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30
6. ACKNOWLEDGMENTS . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31
7.
= 0.75 in Basic Solutions 274.2.3.1 Comparison of the Solubility with Other Solids at 25 and
75°C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 274.2.3.2 Solubility of
Na2,[H(H20)]2_2$b206.H20, a (NaSb(OH)h) in Basic Solutions . . . . . . . . . 24
4.2.3 Solubility of NaSb03.3H20(s)
“NaSb(OH)h(s)” and OtherSodium Antimonates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23
4.2.2 Solubility of
Antimony(IlI) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 194.2 Antimony(V) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23
4.2.1 Rationale for the Measurements Using
300°C ................................3.3.3.1 Preliminary Results .............................................................3.3.3.2 The Solubility of Solid B as a Function of Temperature
and Hydroxide Ion Concentration.......................................
77777
121212131515
17
4. DISCUSSION . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174.1
Antimony(III) .....................................................................................................3.3 Antimony(V) ......................................................................................................
3.3.1 Preparation and Characterization of the Solid Phases ..........................3.3.2 Solubility Experiments for Temperatures Below 100°C.. ....................
3.3.2.1 Preliminary Results .............................................................3.3.2.2 The Solubilities of Solids B and C at 25 and 75°C ............3.3.2.3 Other Experiments ..............................................................
3.3.3 Solubilities for Temperatures from 200 to
2 200°C .......................................3.2
MEASUREMENTS ..............................................3.1 General Procedures for Measurements for T EXPERIMENTAL SOLUBILITY
Sb205 ........................................... 52.3.3 Previous Solubility Measurements for Sodium Antimonate(V) ............... 6
3.
Sb203 ........................................... 32.3.2 Previous Solubility Measurements for
THB CI-IFMISTRY OF ANTIMONY(III) AND ANTIMONY(V) ............................. 22.1 The Aqueous Species ......................................................................................... 22.2 The Solids ........................................................................................................... 22.3 Previous Solubility Measurements ..................................................................... 3
2.3.1 Previous Solubility Measurements for
i
CONTENTS
1. INTRODUCTION . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1
2.
a Calculated from the Results at allTemperatures ....................................................................................................Literature Tabulations of Chemical Thermodynamic Values for AntimonyAqueous Species ...............................................................................................Literature Tabulations of Gibbs Energy of Formation Values forAntimony(III) and Antimony(V) Oxides at 25°C .............................................Literature Tabulations of Enthalpy of Formation Values for Antimony(III)and Antimony(V) Oxides at 25°C ....................................................................Literature Tabulations of Entropy Values for Antimony Oxides at 25°C ........
3737
40
4668
11
11
13
14
15
16
16
1822
26
29
37
37
3838
logloK,(25”C) and
(logioK,) with Values ofa Calculated from the Experimental Results for Each Temperature and fromValues of
Na2u[H(H20)]2_2,Sb206.H20 NaSb(OH)h ...................................................................................................
Activity Products for
NaSb03or
Sb203. .............Values of the Solubility Product for Solids Nominally Hydrated
NaSb(OH)e (initially Solid A) from the PresentStudy .................................................................................................................Total Antimony Concentrations for Solids B and C as Measured for BasicOxidizing Solutions at 25 and 75°C .................................................................Results of Equilibration of Mixed Antimony Solids with Water at 75°C(unless otherwise noted) ...................................................................................Total Antimony Concentrations as Measured over Solid B (initially solid A,but converted to solid B during the experiment) for Basic OxidizingSolutions at 250°C ............................................................................................Total Antimony Concentrations as Measured over solid B (initially solid A,but converted to solid B during the experiment) for Basic OxidizingSolutions after Heating to 250°C and Cooling to Room Temperature.. ...........Total Antimony Concentrations for Solid B as Measured for BasicOxidizing Solutions at 200 to 300°C or after Cooling to Room Temperature.Calculated Thermodynamic Quantities for the Dissolution of
SbzOs/Sb Mixtures ......................Results of neutron activation analyses of solid B .............................................Molar Mass per Mole Sb of Various Antimony(V) Compounds ContainingOxygen, Sodium or Hydroxide Ions or Water ..................................................Experimental Solubilities of
SbzOS .........................................................................Reported Solubilities of Sodium Antimonate in Water ....................................Experimental Solubility Measurements for
Sb203 .........................................................................Reported Solubilities of
4- 1:Table 4-2:
Table 4-3:
Table Al:
Table A2:
Table A3:
Table A4:
LIST OF TABLES
Reported Solubilities of 2- 1:Table 2-2:Table 2-3:Table 3-l:Table 3-2:Table 3-3:
Table 3-4:
Table 3-5:
Table 3-6:
Table 3-7:
Table 3-8:
Table 3-9:
Table
.....................................................................................A.1 Simple Aqueous Ions and Hydrolysis Species of Antimony.. ..........A.2 Antimony(III) and Antimony(V) Oxide Solids ................................A.3 Chemical Thermodynamic Measurements for Mixed Oxides
Containing Antimony .......................................................................
Table
..........
ii
Appendix A: Literature Thermodynamic Data for Aqueous Antimony Species and SelectedOxide Solids
Na~.&IO.~Sb&l&O) . . . . . . . . . . . . . . . . . . .
3839
40
9
1010
1922
23
25
29
Sb203 at temperatures from 90 to 300°C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .Calculated total solution concentrations of Sb(III) as a function oftemperature and hydroxide concentration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .Sodium ion concentration (M) as a function of total Sb(V) concentration forsolubility measurements of sodium antimonates in basic solutions at 25 and75°C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .Solubility measurements for solid B, a mixed oxide of antimony(V)(hydrated pyrochlore-structure sodium salt,
.
Solubility of . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .Sb203 from 15 to 50°C (for details, see text)
- Solid C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .Total concentrations of antimony(III) in aqueous solution in equilibriumwith
NaSb(OH)e18]..........................
X-Ray Diffraction Pattern for Sb205 [ - Solid B; (b) Partially Dehydrated
- Solid A . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .X-Ray Diffraction Pattern for (a) the Pyrochlore Structure SodiumAntimonate
NaSb(OH)e
4- 1:
Figure 4-2:Figure 4-3:
Figure 4-4:
Figure 4-5:
LIST OF FIGURES
X-Ray Diffraction Pattern for
.111
Table A5: Literature Tabulations of Heat Capacity Values for Antimony Oxides at25°C . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Table A6: Temperature-Dependent Heat-Capacity Values for Antimony Oxides . . . . . . . . . . . .Table A7: Temperature-Dependent Heat-Capacity Values for Alkali
Metal/Antimony(V) Mixed Oxides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Figure 3-l:Figure 3-2:
Figure 3-3:Figure
. .
CANada Deuterium Uranium; registered trademark.* CANDU:
300°C. Theresults from these studies were used to draw conclusions about the relative stabilities of theantimony solids, the nature of the aqueous antimony species and, for oxidizing conditions,constraints on the probable total concentrations of antimony species in solution as a function oftemperature.
[7]. A more detailed study was made for antimony(V) solids. Different sodiumantimonate(V) solids were prepared, and characterization attempted. Solubilities of theseantimony(V) solids in basic solutions were measured at temperatures from 25 to
3OO”C, and to confirm the literatureresults at 200°C
Sb203 did not change greatly between 200 and
pH and temperature.
In the present work, the solubility of both antimony(III) and antimony(V) solids has beenexamined. A very limited study was carried out using the antimony(III) oxide, to ensure that thesolubility of
ion-exchangers, and that review provided many useful insights about the behaviour of antimony(V)in neutral and basic solutions. For both antimony solids and aqueous species (see Appendix A),values in several standard tables of chemical thermodynamic data differ substantially. Solubilitystudies are one means of obtaining information about changes in aqueous species as a function of
[8] thoroughly reviewed thequalitative features of the antimony(V) oxide solids related to the preparation of inorganic
[2-71. These span a fairly widerange of temperature (15 to 200°C). Belinskaya and Militsina
Sb203, in neutral to basic solutions have been reported antimony(lII)
oxide,
high-temperature (50 to 300°C) solution properties of antimony, for both reducing and oxidizingconditions, and hence to help determine conditions that minimize antimony release, transport anddeposition.
The nature of antimony species in aqueous solutions, and the solids, stable and metastable, thatcan exist in contact with such solutions, is very complex, especially for oxidizing conditions.There does not appear to be any single paper or document that describes antimony solutionspecies and solids in a coherent, comprehensive manner. This is particularly true if the behaviouras a function of temperature is of interest. Several studies of the solubility of
11. Although it is not certain whetherthe antimony is initially mobilized by a physical or chemical mechanism, it is clear that antimonycan be readily released and transported in solutions under oxidizing conditions. The primaryimpetus for the work described in this report was to enhance our understanding of the
11. Antimony is also presentin some pump seals and bearings at Gentilly-2. Oxygen excursions during shutdown atGentilly-2 have resulted in large increases in out-of-core radiation fields, adversely affectingscheduled maintenance. This has led to the routine use of an oxidizing antimony removalprocess at the start of each annual maintenance shutdown [
123Sb, an element found as a minor (butunmeasured and unspecified) component of some reactor materials [
121Sb and ‘%b. These antimony isotopes are activation products from irradiation of
the naturally occurring isotopes of antimony,‘22Sb and
CANDU* primary heat transport systems(HTS) are
1. INTRODUCTION
Among the major contributors to activity transport in
191 have181. Rumpel et al. [ P-Sb204 [ 181. Further heating to 935°C yields Sb204.35, that has a defect
pyrochlore structure [
Sb205)to between 650 and 850°C leads instead to a partially reduced solid,
[8]. Heating antimonic acid (i.e., hydrated Sb205 by heating in air at 1 bar Sb205) cannot be
dehydrated to
Sb205 (or even on material prepared in research laboratories) to be suspect unlessproper characterization is provided for the solid. Antimonic acid (hydrated
“Sb205)‘, and consider any studies based on commerciallyavailable
[8]. We found similarproblems with currently available
Sb204, and this was also reported by other groups Sb204.4) or
Sb205)’were actually found to be either an amorphous, partially reduced solid (of approximatecomposition
171 noted that samples of commercially available “anhydrous
[3].
In 1970, Stewart and Knop [
131.
2.2 The Solids
Antimony(III) oxide exists in two forms. The commercially available orthorhombic form of theoxide (valentinite) is easier to prepare, and occurs more commonly in nature. However, the cubicform (senarmontite) is reportedly more stable near room temperature
0.5Sb205(aq) (2.2)
This does not appear to be compatible with the spectroscopic study of Jander and Ostmann [
+ SbOj (aq) + (aq) Sb20’: + 0.5 (aq) SbOi
“K3Sb04” resulted in the sequential conversion:
161,based on potentiometric and conductometric experiments, proposed that addition of acid to asolution of
[ -(aq) in solution does not yet appear to be proven, Prasad SbO$
H20 (2.1)
Although, the existence of
SbO(OH):- + * Sb(OH); + OH-
> 12) there is an equilibrium between twomonomolecular anionic species, with the most probable reaction being:
(pH mol.dm-3. The same authors, using absorption spectroscopy,
found that for basic Sb(V) solutions
[13] reportedfinding evidence for polymer formation only for acidic solutions with total antimonyconcentrations greater than 10”
[SbsOi2(OH)& anion. Jander and Ostmann 151 reported the isolation and crystal
structure of a salt containing the . Recently, Nakano et al. [ H12_,(SbO&
141 proposed a series of anionicdodecamers
l-14]), it is not clear what species willform in very dilute solutions of Sb(V). Lefebvre and Maria [
macro-concentrations, Sb(V) forms polymeric species. As for many hydrolytic polymers, once formedthese are slow to depolymerize, even though simpler species may be thermodynamically stable ina particular solution. Therefore, even though there have been several studies of the speciesformed in weakly acidic and neutral solutions (e.g., [ 1
pH values between 1 and 2 at room temperature. However, at HSb(OH)6 is a moderately
strong acid, ionizing at lo], it is implied that [9,
Snecies
At equilibrium in aqueous solution, antimony generally forms antimony(III) species underreducing conditions, and antimony(V) species under oxidizing conditions. Antimony (III)chemistry is moderately well understood, at least near room temperature. This is not true forantimony(V). In some general references
2
2. THE CHEMISTRY OF ANTIMONY@) AND ANTIMONY(V)
2.1 The Aaueous
[2], all these measurements were done usingsolutions at temperatures between 15 and 50°C. In general, agreement between the results of thedifferent studies is excellent-much better than most studies of oxide solubilities.
Schulze [7] and a single experiment by Popova et al. c$ Figure 4-l). Except for the work of[2-71 (Table 2-l; also
Sb203 in water and basic aqueous solutions have been reported previouslyby a large number of authors
Sb203
Solubility values for
[S].
2.3 Previous Solubilitv Measurements
2.3.1 Previous Solubility Measurements for
NaSbmSbT07.The nature of the solids is discussed more thoroughly by Belinskaya and Militsina
[33] reported a mixed-oxidation-state pyrochlore-structure solid,
[32] used different concentrations of alkali metal hydroxide solutions at room temperatureto control the extent of substitution of the metal cations into the oxide pyrochlore structure.Gol’dshtein et al.
Sb205 in acidic media, while Baueret al.
l] demonstrated ion-exchange properties for hydrated [3 [30] and
Abe (NaSb(OH)h) at 180 to 320°C. Baetsle and Huys “SbOsNa” trihydrate
“Sb2Na205(OH)2” bydehydration of
[29] prepared a pyrochlore-type solid
[23]gave no indication of any analysis of his material for hyperstoichiometric water.
Alkali metal antimonates having a pyrochlore structure have been prepared by wet and drymethods. Montmory et al.
[22] to inclusions of some of the motherliquor in the crystals prepared at lower temperatures. It is therefore unfortunate that Asai
[24]) was recovered fromcooler solutions. The difference has been attributed
Karlicek Dravotsky and NaSb03.3.5HzO of 6HzO”
(possibly the same as the “Na2H2Sb207 + H20”) precipitated from hot solution, but that ((‘Na2H2Sb207 + 5
[28], that the hexahydroxy-compound[22] found,
in agreement with Knorre and Olschewsky Na+ ions. Beintema Sb(OH)i octahedra with distortion to accommodate the
[23]. The structurehas
NaCl(aq) “KSb(OH)6” with NaSb(OH)e
as prepared by treatment of a dilute solution of NaSbOs(3-x)H20. A full crystal determination has been done for
NaSb(OH)h and an amorphous solid with aformula
“NaSb(O&” is probably a mixture of KSb03.2.3H20, and
suggested that “KSb(OH)e” corresponded to the compound being
171 reported that analyses ofcommercially available
[27]. Similarly, Stewart and Knop [ K[HSb03(0H)].H20, and other possibilities were
discussed by Balicheva and Roi [26] suggested
MSb(OH)estructure. Lisichkin et al.
[25]. The latter formula is not compatible with a KSb03.2.6H20 [24], whereas the potassium salt is
recovered as MSbO3.xI-IzO (x = 3.5-3.6)
[21]. The lithium and sodium salts, as recovered from concentrated aqueous solutions, have theapparent stoichiometries
LiSb(OH)h are reported to be hexagonal or trigonaltetragonal[21-231, crystals of NaSb(OH)G are
[20].
A wide variety of solids containing antimony(V) and alkali metals have been reported. Thedifferent salts with antimony to alkali metal ratios of 1: 1 apparently either do not have the samestructures, or, at least in some cases, are not simple hexahydroxy salts. Although crystals of
Sb50i20H.H20, for which the crystal structure was reportedby Jansen
Sb204.s(OH)O.4 is closely related to (orpossibly identical to) the compound
900 bar. The
Sb205 byheating (at 300 to 750°C) various antimony oxides in the presence of small amounts of waterwith oxygen gas at pressures from 80 to
Sb204.s(OH)0.4 and two forms of anhydrous Sb204.4(OH)r.2,
3
reported the preparation of
........................................................................
Cont’d.. _ . . . . . . . . . . . . . . . . . _ . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . - . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . [61. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.60x10” r?
[61[41[41[41141141[41[41r41[41[41[41[41WIU-dWI161WI[61[61151[51151[51151151[51c51[31[31[61WIWIWI161WIWIPI
4.29~10-~ r?
4.32~10-~ r?0.0129 r?
2.52~10-~ r?4.87~10~ r?2.56~10~ r?1.3OxlO~ r?9.7ox1o-5 r?6.8ox1o-5 r?4.80~10‘~ r?4.5ox1o-5 r?6.1~10~ r?9.1ox1o-5 r?
9.78~10” r?0.0117 r?0.0142 r?
6.86x1o-3 r?5.06~10-~ r?3.26~10” r?1.63~10-~ r?9.98~10~ r7.58~10~ r3.78~10~ r4.28~10~ r2.38~10-~ r1.48~10~ r9.80~10-~ r5.20~10-~ r4.oOx1o-5 C
9.20~10-~ r
9.43x10” r?0.0117 r?
7.63~10-~ r?5.78~10” r?4.12~10” r?2.57~10” r?1.3ox1o-3 r?5.50x10” r?
1.0097*0.6732*
6.x10”0.01030.04240.09150.4580.7021.99
6.~10~3.49x10a6.~10-~1.58~10-~1.45x1o-72.3572*
1.3482*1.684”2.0209”
1.0112*0.6742*
1.oox1o-70.005050.01010.02020.04040.040.07490.09980.3372”
1.oox1o-71.oox1o-7
1.3493*1.6862”2.0232”2.3597”
1.0122*0.6749*
6.71x10-’0.3375”
SbzO&Reference
151515151515151525252525252525252525252525252525253535353535353535353535353535
T/“C Form of
Sb203
4
Table 2-l: Reported Solubilities of
[lo] concluded that it isunlikely a pure solid phase was present in these experiments, because of the ease with which gels
2-2), Baes and Mesmer
[4]. Their solid was prepared by hydrolysis of the Sb(V) hydrochloride salt and dried by heatingto 90°C. Although the reported total concentration of aqueous antimony species in water incontact with this solid was quite low (Table
Sb205 at 35°C in acid solutions was reported by Tourky and Mousa
Sb205
The solubility of (hydrated)
$ r:orthorhombic form (valentinite); c: cubic form (senarmontite)
2.3.2 Previous Solubility Measurements for
(*). In those cases, the hydroxide concentration was alsodetermined at the end of the experiment, and is the value reported here.
t Initial hydroxide ion concentration, except when marked by anasterisk
171[71
C
[71C
[71C
171C
[71C
[71C
[71C
[71C
[71C
[71C
[71C
[71C
[71C
[21C
II71r?
[71C
[71C
[71C
[71C
[71C
[61C
[61r?
[61r?
WIr?
WIr?
WIr?
WIr?
WIr?
[61r?
8.85~10”0.1 0.01450.1 0.0153
r?
8.04x 10”0.03
5.92~10”0.03
5.87~10”0.01
4.35x10”0.012.23~10-~
3.90x10”2.23~10-~4.4ox1o-32.23~10-~4.90x10”1O-62.23x 4.53x10”2.23~10-~4.36~10”2.23~10-~4.52~10-~2.23~10‘~4.44x1o-32.23~10-~3.40x10”7.28~10-~290x10~5.95x1o-73.3ox1o-45.95x1o-72.70~10”5.95x1o-73.20~10~5.95x1o-74.oOxlO~5.95x1o-73.1OxlO~5.95x1o-7
2.0134* 0.01921.6789* 0.0149
8.65~10‘~1.3432” 0.0119
5.95x1o-31.0077”0.6715*
1.6827* 0.01112.0185” 0.01412.3553” 0.0161
8.66x10”SbzO&
Reference
353535355050505050909090909090100200200200200200200200200200200200200200200
1.3464”mol*ti3
[Sh]T/ Form ofmol*ti3t[QH’yT/“C
2- 1 (Concluded)
5
Table
[371
Notes
avg. Sb, Na analyses
r391100 0.03
[38180 0.0093
[39175 0.030
[39170 0.0084
[38150 0.0060
c39150 0.015
136135 0.0044
~25133.5 0.00412
[38125 0.0033 139125 0.012
[37118 0.00229 136125 0.00299 136125 0.0053
(?) 0.00538[Sb],/mol*ti3 Reference
15 T/“C
pH or hydrogen ion concentrationwas not reported for any of the experiments. Consequently, the nature of the anionic antimony
Table 2-3: Reported Solubilities of Sodium Antimonate in Water
[35-391 (Table 2-3).
One of the problems with the reported solubilities is that the
[341 water
2.3.3 Previous Solubility Measurements for Sodium Antimonate(V)
There have been several reports of solubility measurements for sodium antimonate in water
HClwater
100 0.0212
mol.dm-3HCl
0.050mol.dm-3
HCl0.100
mol.dm-3HCl
0.5 16mol.dm-3
HCl1.064
HCl1.981 moldm”
HCl2.458 moldm”
HCl2.900 moldm”
HCl3.748 moldm”
HCl4.092 moldm”
mol.dm”
141
4.600
[4135 0.00027 1
0.ooo101141
35
14135 0.00007 1
[4135 0.000057
[4135 0.000043
14135 0.000035
[4135 0.000059
14135 0.000125
[4135 0.000287
[4135 0.000372
[Sb]T/mO&i3 Reference Medium35 0.000487
T/“C
Sb205
15-minute equilibration at 100°C.
Table 2-2: Reported Solubilities of
[34] derived a “maximum”solubility value from a
6
are prepared from such solutions. Glixelli and Przyszczypkowski
(3.1)
Three different solids were used in the present solubility studies, and at least one of thesematerials was probably a mixture or solid solution.
Antimonv(V)
3.3.1 Preparation and Characterization of the Solid Phases
3- 1.
3.3
300°Csolubility experiments was essentially identical to the pattern for the initial valentinite. Themeasured solubilities are listed in Table
Sb203(c)
The XRD pattern of the oxide in the mixture recovered from the autoclave after the
+ 1.502(g)
99.9999%), and thismechanical mixture was used as the charge in the static autoclave. Excess oxygen would then betaken up by the metal, the overall reaction being:
Sb(c) +
Sb203 was mixed with metallic antimony shot (Alfa,
Sb203 (Aldrich, 99.999%) was used in these experiments without further purification. The XRDpattern for the oxide showed lines only for the orthorhombic form (valentinite). Because it wasuncertain whether traces of oxygen would generate Sb(V) in our solutions at elevatedtemperatures, the
Antimonv(III)
Ko,i radiation. Antimony concentrations were determined by neutron activation and inductivelycoupled plasma-atomic emission spectroscopy (ICP-AES).
3.2
pm silver filter. Immediately after passing through the hot filter, the sample wascondensed, weighed and acidified. Samples of the final solid(s) were recovered at the end of theexperiment after the autoclave had been cooled to room temperature. Powder X-ray diffraction(XRD) patterns were obtained for the antimony solids using a Siemens Diffractometer with Cu
mL Autoclave Engineering titanium autoclave equipped with a stirrer.Solutions were sampled at temperature by preheating the stainless-steel filter holder andsampling line to the temperature of the autoclave. The sampling-line connection to the autoclavewas then opened, and liquid driven by the hydrostatic pressure in the autoclave was forcedthrough a 0.45
300
2 200°C
The solubility measurements at 200 to 300°C were carried out in aqueous sodium hydroxidesolutions in a
[25], where both the sodium and antimonyconcentrations were measured, is very high, and indeed greater than the solubility of the lithiumsalt as reported in the same paper.
3. EXPERIMENTAL SOLUBILITY MEASUREMENTS
3.1 General Procedures for Measurements for T
Karl&k Dravotsky and [38] are markedly higher. The value based on the %
composition analyses of
[39] are similar,while those reported by Urazov et al.
[36] and Blandamer et al.
7
species is not known. The values reported by Tomula
[21,40], although therewas considerable variation from sample to sample. In some spectra, several of the peaks in thediffraction pattern did not appear (the pattern was then consistent with a face-centred cubicstructure, a = 0.80 nm), while many of the remaining peaks were very strong and sharper thanfound for other samples. It may be that certain samples as prepared for XRD analysis werelayered, causing some peaks to be weak. However, the “cubic” pattern could not be specifically
NaSb(GH)h &i radiation (Figures 3-1,3-2(a) and 3-3). The powder XRD patterns found for some
samples of solid A were consistent with that reported for
25O”C, and was filtered andoven-dried at 105°C.
Solid C was found to form on heating sodium antimonate (solid A) in contact with aqueoussodium hydroxide solution at 75°C for times ranging from several days to several weeks. Thesolid was recovered from the solution by filtration and oven-dried at 105°C.
Powder XRD patterns for the three solids were obtained using a Siemens Diffractometer withCu
moldm”) of aqueous sodiumhydroxide, either at room temperature, or by treatment for 1 d at
lo4 NaOH(aq) solution in a titanium autoclave at 250°C for one week. The
residual solid was washed with a very dilute solution (
cm3 of cold ethanol. The solid was dried in an ovenat 105°C.
Solid B was found to form on heating sodium antimonate (solid A) in contact with0.05 mol dm”
cm3 of ice-cold distilled water and 250
cm3 of distilled water (at 50°C) and then the solution was immediately cooled, first to 25°C(for several hours), then in an ice-water bath. The crystals were filtered and washed with250
NaCl dissolved in10
75”C, and the solution temperature was then adjusted to50°C. Precipitation of sodium antimonate was initiated by the addition of 1 g
cm3) at [38,39]. Potassium antimonate (Aldrich Chemicals, 5 g) was dissolved in cooled boiled-out,distilled, deionized water (100
NaSb(OH)b, (solid A) was prepared as described in the literature
mol*dm -3
0.00670.000420.00750.00620.00220.00120.00490.0016
not measurednot measured
0.0014
Sodium antimonate, nominally
[OH’]e,,d mol*ti3
6 0.0054 0.01043 0.0023 0.0028 11 0.0032 0.01041 0.0016 0.01041 0.002 1 0.0028 12 0.0016 0.0028 12 0.0017 0.01046 0.0023 0.0028 11 0.00064 0.01045 0.00048 0.0028 12 0.00034 0.0028 1
mol*ti3PH’liniti~
Duration/d[Sbhd
16A-3 2517-3 25
Contact
16A-1 30014-3 25
16A-2 25017-2 25015-1 300
T/“C
14-1 20017-1 20014-2 25015-2 250
Sb2O$Sb Mixtures
Run
8
Table 3-l: Experimental Solubility Measurements for
- Solid ANaSb(OH)h Pattcm for
10
Figure 3-l: X-Ray Diffraction
2030405060Xl40
NaOH(aq) at room temperature, and conditioned for a furtherday at 250°C in that medium.
moldm‘3 NaOH(aq). Sample B-9A
was washed with 0.0001 mol+dm-3
NaOH(aq) at roomtemperature, and conditioned for five days at 250°C in 0.01
mol.dm‘3 NaOH(aq) at
250°C for seven days. Sample B-6 was further washed with 0.0001 mol.dm-3
* 13) g per mole antimony (assuming uncertainties of 5% in the neutron-activationanalysis values). Analyses for three separately prepared samples of solid B are shown inTable 3-2. All three samples were prepared by heating solid A in 0.05
0.07), and an apparent molarmass of (265
f
ln some cases, the patterns of thesolid recovered after long equilibration periods at 75°C resembled the initial solid; in other cases,certain peaks were markedly weaker.
Neutron activation analysis was carried out on samples of solids B and C. Solid C was found tobe 8.7 wt.% Na, 46 wt.% Sb, i.e., a Na:Sb atomic ratio of (1.00
“NaSb(OH)i’ solid.
[18] for samples fired in air at 220°C for 50 h (Figure 3.2(b)). However,elemental analyses of our pyrochlore-type solid B (see below) indicate that it is not a simplehydrated oxide, as it contains substantial amounts of sodium. For samples of solid C, the XRDpatterns were related to those of the initial
Sb204.35 is complete). Thereported changes in the peak positions during the heating are small; however, the intensitieschange markedly. The pattern of the material recovered from our experiments (presumably incontinual contact with water during the experiments) very closely resembled the pattern reportedby Stewart et al.
stepwise heating of “antimonic acid” in airfrom room temperature to 735°C (at this temperature, reduction to
[18] reported aseries of XRD patterns showing the transformation on
Sb205.xH20. Stewart et al.
9
identified with any found in the literature. Initially, it seemed that the “cubic” solid could beidentified with solid C discussed below; however, additional experiments showed the system tobe more complicated.
Solid B provided a pattern that was identifiable with the pyrochlore structure, similar to thepattern of partially hydrated antimony(V) oxide,
- Solid CNaSb(GH)h
181
Figure 3-3: X-Ray Diffraction Pattern for
Sb205 [ (b) Partially Dehydrated -
Solid B;
28
Figure 3-2: X-Ray Diffraction Pattern for (a) the Pyrochlore Structure Sodium Antimonate
2820 103050 40002-l 10 0
703060 4060
O-70
alIWO-
zI40
1am-
Iii.8i
i,Km-%
g-
80II(b)(3) Imw
10
100
250°C, and forms apyrochlore solid akin to the hydrated oxide, but incorporating at least some sodium in the solid.However, the material recovered from the autoclave is sufficiently insoluble, under mostconditions, that quantitative elemental analysis by standard methods is difficult. The peaks in thediffraction pattern are quite sharp, whereas the presence of a large percentage of amorphousimpurity in the solid might have been expected to cause substantial broadening or an erraticbaseline. Various compounds based on polymeric hydrated antimony(V) oxide and having the
[28]. Solid B has a fairly high sodium content,considering the similarity of the XRD pattern to that of the pure hydrated oxide. On the basis ofthe diffraction pattern, it was concluded that solid A decomposes in water at
“dihydropyro-antimonate” (but see Ref. 8). The estimated water-to-antimony ratio for solid C is slightlygreater than might be expected from earlier work
NaSb03.3.5HzO 255.8264.8
Comparison of the apparent molar masses (Table 3-3) with those for various possible solidssuggests that solid C is probably a hydrated sodium antimonate or a hydrated
Na2H2Sb207’6H20 255.8Na2H2Sb207.5H20 246.7NaSb(GI& 246.7NasHSb40i2.6H20 214.3NaSbOs.H20 210.7NaSb03 192.7Sb205.H20 170.8Sb2G5 161.8
f 13212+4
solid C 265
f 4) g per mole antimony.
Table 3-3: Molar Mass per Mole Sb of Various Antimony(V) Compounds Containing Oxygen,Sodium or Hydroxide Ions or Water
Formula App. Molar Mass perSb
solid B
0.03), and an apparentmolar mass of (212
+
+ 0.03 212
From the average, assuming 5% uncertainties in the analyses, solid B was found to be8.1 wt.% Na and 57 wt.% Sb, i.e., a Na:Sb atomic ratio of (0.75
f 0.05 21057.3 0.75
+ 0.05 21458 0.71
z!z 0.06 21457 0.74
11
Table 3-2: Results of neutron activation analyses of solid B
Sample wt.% Na
6 8.57 8.0
9A 7.8avg. 8.1
Apparentwt.% Sb Na:Sb Molar Mass
per Sb57 0.79
NaOH(aq) solutions at 25 and 75°C. These results are discussed inSections 4.2.2 and 4.2.3.
moldm-3
25”C, as the solutionswere left to re-equilibrate at the lower temperature for ten days before the solubilities weremeasured. It is also possible that solids A and C are essentially identical, and that theexperimental equilibration periods used were too short.
3.3.2.2 The Solubilities of Solids B and C at 25 and 75°C
Experiments were also done (Table 3-5) to establish the solubilities of solids B and C in0.003 and 0.04
NaOH(aq) solutions.
The apparent increase in the solubility of the solid on cooling to 25°C might suggest that thesolid that formed at 75°C is less stable than solid A at 25°C. However, the kinetics fortransformation of solid C to the original solid A must then be quite slow at
moldrn‘3 NaOH(aq) solution at 75°C. The solubilities were considerably greater at 75°C than at 25°C forboth the 0.01 and 0.1
mol.dm-375”C, as was the 0.01 NaOH(aq), and all solutions were undersaturated both at 25 and moldrn-3
0.5)“C for17 days, sampled, equilibrated at 75°C for four days, sampled and finally re-equilibrated at 25°Cfor ten days before final sampling and examination of the solids. The results are listed inTable 3-4. Unfortunately, insufficient solid was used in the experiments with 0.001
+
(Na2a[H(H20)]2_~Sb206.H20, with a = 0.75).
3.3.2 Solubility Experiments for Temperatures Below 100°C
3.3.2.1 Preliminary Results
The solubility measurements at lower temperatures were carried out in Nalgene high-densitypolyethylene bottles held in a thermostated bath. In a preliminary study of the solubility ofsodium antimonate (solid A), the solutions were initially equilibrated at (25.0
[32]
NaOH(aq) to contact the solution is weak. Of course, the extentof hydration of the solid may well have changed during cooling of the autoclave to roomtemperature. It is also possible that solid B as synthesized in the present work is a mixture of amixed oxide and the hydrated oxide and/or an amorphous sodium-containing solid. However,except for the extent of hydration, the formulation is consistent with that proposed by Baueret al.
B-9A, and the differences in the ratios arewithin the uncertainty limits of the analyses. Thus, any correlation between the value of a andthe concentration of the last
> > B-6 Na+(aq) in the last
solutions contacting the solids was B-7 B-9A, whereas the concentration of > > B-7
[32]. The Na:Sb ratio in the three analyzed samplesdecreased in the order B-6
I 0.87) reported by Bauer et al. 5 a
(Sb205)(Na20)0.75.3.04H20,but this is within the continuous range of solid solutions with the pyrochlore structure(0.20
5 0.67. The analyses of our solid Bare consistent with a gross formula with a somewhat greater ratio,
171 concluded that the pyrochlore structure would he found only forsodium antimonate compounds having a Na:Sb ratio of
NaSbmSbT07.Stewart and Knop [
[33] reported a very similar pattern for a solid characterized as Na2,(H30)2-22,Sb2G6’H2G, and
Gol’dshtein et al. [32] reported patterns for a series of partially substituted solids
[8]. For example, Bauer et al.
12
pyrochlore-type structure generate almost identical XRD patterns
131.
121Sb line-width measurementsindicate that the Sb(V) in aqueous basic solutions is not in a totally symmetrical environment (orthat more than one species is present). This result may be related to whatever phenomenon wasresponsible for the absorption spectroscopic results of Jander and Ostmann [
NaOH(aq) standards. The residual solids wererecovered for XRD analysis. The analysis results are listed in Table 3-6.
Preliminary nuclear magnetic resonance (NMR) results from
pH (by also usingmeasurements with standard acid solutions), and that sodium errors were either negligible orcould be corrected by comparison with the
NaOH(aq) solutions of different known concentrations weremeasured with the same electrodes on the same day, and the unknown hydroxide concentrationwas determined by comparison with these standards. Checks were done to ensure that theresponse of the electrodes was approximately Nemstian over a wide range of
pH with low sodium ion error) against an Accumet 13-620-5 1 calomel reference electrode.The potentials for several standard
pH meter) the potential of an Accumet 13-620-295 glass electrode (forhigh
COz(g). Bottles weresampled after 2 1 to 155 days and solutions were submitted for Sb and Na analysis by ICP-AES.The hydroxide ion concentration of each final solution at room temperature was determined bymeasuring (Accumet 25
cm3 of deionized, distilled water inNalgene high-density polyethylene bottles held in a thermostated bath. Each bottle wascontained in a closed outer bottle to minimize the possible ingress of
25”C),in an attempt to determine the relative stabilities of solids B and C. Samples of the solids(approximately 0.2 g of each) were contacted with 10
75.4”C and are probably solubilities for the metastable solid C (see text).
3.3.2.3 Other Experiments
A further series of experiments was carried out at 75°C (and one additional experiment at
f 0.0001)* Values were obtained from re-equilibration of the solutions held four days
at
f 0.0001)0.1 75.4 (0.0023
f 0.0002)(0.0023 2 (0.0029
i 0.0001)2 (0.0016
f 0.00001)”
75.475.475.4
o.OOOOl)*25.4 (0.00014
f O.OOOOO9)
25.4 (0.00015 f
O.OOOOO9)25.4 (0.000100
+ + 0.00008)”
25.4 (0.000100
f 0.00004)25.4 (0.00132
* 0.00004)25.4 (0.00065
O.OOOl)*25.4 (0.00060
* 2 (0.0016 f 0.0001)
25.42 (0.0016
[Sb]T/mol*dm”0.0010.0010.010.010.010.100.100.100.10
0.0010.010.1
25.4T/“C[NaOHJ/moldni3
NaSb(GH)h (initially Solid A) from the Present Study
Initial
13
Table 3-4: Experimental Solubilities of
NaOH(aq) but noantimony solid.
0.045 -- 46* Shaded results are for experimental “blanks” using vials containing
lob51.8 xOooo6)f(0.092ooooo5 -- 43
0.0438lO&a.6 x0.oooo3)f(O.ooo62O.OOOl) 0.039 c2 43
0.0029ff 0.003) (0.0016 O.OOOl)
0.038 c2 43(0.053
ff0.003) (0.0016
O.OOO2)0.041 c2 21
0.0438 (0.052
ff0.0438 0.003) (0.0027
O.OO27?) 0.026 c2 21(0.047
f f 0.002) (0.0029 210-9 C2 43
0.04380.0438
(0.040 110-5, Oooo5)fOool) (0.0073 f
110-9 C2 430.0029 (0.010
110-5, Oooo5)fOool) (0.0074 foooo34 c2 21
0.0029 (0.011Oooo5)f (Ooo88 Oool)f
210-9 c2 210.0029 (0.012
110-5, O.OOO6)+Oool) (0.0100fooooo3 j 0.036 B-9A 43
0.0029 (0.014f(O.OOO46+ 0.003j
0.oooo3) 0.040 B-9A 430.0438 (0.049
f (Oooo47 f 0.003)O.OOOO3) 0.042 B-9A 40
0.0438 (0.050 f (O.OOO46 + 0.002)
0.oooo3) 0.037 B-9A 210.0438 (0.047
f (Oooo45 f 0.002)(0.044 O.OOOll B-9A 46
0.04380.oooo5)f (Oooo77 O.OOO3)f
O.OOO15 B-9A 270.0029 (0.0056
O.OOOOS)f (O.OOO72 Oooo2)f (Ooo55 O.OOO14 B-7B 25
0.0029ooooo3 jf (O.OOO48 O.ooo2)f (Ooo44
O.OOO40 B-7B 250.0029
0.oooo4)f (Oooo64 O.OOO2)f 7s”c
0.0029 (0.0056
lOa 0.042 -- 83O.ooO) 51.8 xfO-0438 (0.043o.oo17 -- 831(-P<l.& xO.oooo)f(U.oo33
O.OOOO2) 0.039 Cl 83
0.0029+(O.OOO25 Oooo)f(0.044
0.049 Cl 830.0438 O.OOOO2)
f(O.OOO26Oooo)
f 0.oooo2)0.045 Cl 43
0.0438(0.044
f(Oooo30
f0.003)
O.OOOO2) 0.043 Cl 430.0438
(0.049
f (O.OOO27 f 0.003)ooo15 Cl 83
(0.048 O.OOOl)f0.ooo1)
0.0438(0.0018 f(O.oo5 1
ooo15 Cl 830.0029
O.OOOl)+ Oooo3) (0.0018 + (Ooo50 ooool j 0.0018 C2 43
0.0029fOooo3) (0.0018 St(Ooo50
ooo19 c2 430.0029
O.OOOl)f (O.oo51~ 0.003) (0.0018 oooool j 0.049 B-7B 83
0.0029f (O.OOol8 oooojf (0.044
0.oooo1) 0.040 B-7B 830.0438
f(0.ooo170.ooo)+-0.oooo1) 0.040 B-7B 43
0.0438 (0.043+(0.ooo17+ 0.003)(0.044
0.oooo1) 0.040 B-7B 430.0438
z!z(0.ooo17f 0.003)(0.0440.oooo1) 0.0011 B-7B 83
0.0438f(0.ooo180.ooo1)f
O.OOOOl) 0.0010 B-7B 830.0029 (0.0025
f(O.OOO18O.OOOl)fooooo2 j 0.0016 B-9A 46
0.0029 (0.0027f (oooo35 oooo2 jf
0.oooo2) 0.0018 B-9A 460.0029 (0.0034
f (Oooo34 Oooo2)*
/mol-fi325°C
0.0029 (0.0034
Duration/dSolid wnxpt[Sb]&nohlm-3[Na]T/mol-ti3
Final[NaOHj/mol-ti3
75”C*
Initial
14
Table 3-5: Total Antimony Concentrations for Solids B and C as Measured for Basic OxidizingSolutions at 25 and
181 and various mixed antimony oxide solids.
NaOH(aq) concentration are for samples taken from the same experimentalrun at different times.
After the autoclave had been cooled to room temperature, but prior to removal of the autoclavelid, nitrogen gas was used to slightly pressurize the autoclave to force liquid samples through thesampling line. The solution concentrations of antimony and sodium species after cooling theautoclave to room temperature are listed in Table 3-8. This table includes the results from laterexperiments (runs done to synthesize further samples of solid B), in which the solutions were notsampled at 250°C.
As discussed previously in Section 3.3.1, the solids recovered from both preliminary experimentshad powder X-ray diffraction patterns that bore no resemblance to that of the original solid A.Instead, these hard, compacted solids had essentially the same XRD (pyrochlore structure)pattern as a slightly hydrated oxide [
1o-9* Initially solid A or C, but probably all solid A was converted to C by the end of an
experiment.
3.3.3 Solubilities for Temperatures from 200 to 300°C
The solubility measurements at 200 to 300°C were carried out in aqueous sodium hydroxidesolutions in a titanium autoclave equipped with a stirrer, as described in Section 3.1.
3.3.3.1 Preliminary Results
Two preliminary experiments were carried out to determine the approximate solubility ofantimony(V) in basic solutions at 250°C. In both experiments, the initial solid was solid A,which was found to convert to solid B during the experiment. The solubilities measured fromthese samples taken at 250°C are listed in Table 3-7. The solubilities reported for the solutionswith the same initial
210-5,<f 0.45f 0.65 7.721o-9
0.2019 0.1992 c 9.965 1 43 (25°C) 9.792 105,If 0.16f 0.24 2.63
1O-53.72
+ 1.6 6 xf 1.7 16.41o-9
27.82 10-5,5Z!I 0.16f 0.17 2.63
1o-93.31
2 10-5,<f 0.7f 0.9 10.51o-9
14.4;r lo-5,5f 0.7+ 0.9 10.8
1o-915.1
2 10-5,5f 0.18+ 0.44 2.921o-9
5.832 10-5,<f 0.16* 0.44 2.71
1o-95.83
2 10-5,<f 0.45+ 0.65 7.76
= 109.9539.925
Days212525424283
213155
9.35*
0.1932 c0.1989 A0.2003 A0.2017 A0.2024 A0.169 C0.2004 A0.1835 c
9.9559.9419.9509.9369.961
moldm”
0.19220.26270.26370.26370.25970.1870.26280.1965
mol*dmJmoldm”NaSb:OHjd Duration[OH’1
solid BlO?Wrl@?JW~J32O Testg g
15
Table 3-6: Results of Equilibration of Mixed Antimony Solids with Water at 75°C (unlessotherwise noted)
mol.dm-3. The similar antimony concentrations found for the solutions at 250°C in the twoexperiments (Table 3-7) suggested that the final hydroxide concentrations in the two experimentswere probably similar, or that the solubility of solid B at 250°C was almost independent of thehydroxide concentration in the moderately strong basic solutions. The final total concentrationsof antimony in solution at room temperature (after cooling from 250°C) appeared to vary withthe measured sodium ion concentrations, but it was unclear whether the differences might alsohave resulted from variable equilibration times at the lower temperature. Therefore, samples of
moldm‘3 hydroxide). The analytical results for the final sodiumion concentrations (Table 3-8) are markedly greater than the initial concentrations, but lower than0.1
(> 0.2
x)HzO(l) (3.2)
Indeed, the final solution from the second experiment (on cooling to room temperature) wasfound to be strongly basic
- 2NaOH(aq) + (5 Sb205.xH20(s) + + 2NaSb(OI-&(s)
f 0.00003)
after 250°C expt.after 250°C expt.after 250°C expt.
re-equilib. 8 d at 25°Cre-equilib. 25 d at 25°Cre-equilib. 41 d at 25°C
re-equilib. 132 d at 25°Cafter 250°C expt.after 250°C expt.
The hydroxide ion concentrations of the final solutions were expected to be dependent almostentirely on the quantity of sodium hydroxide formed from decomposition of the salt.
0.OOoO2)(0.00026
f 0.OoOO2)
(0.00017 f
0.OOoO2)(0.00017
f + 0.00002)
(0.00019
o.ml)(0.00018
f 0.OoOO2)
(0.00017 f f 0.00002)
(0.00015
0.OOoo3)(0.00023
f
f 0.004)
(0.00044
It 0.005)0.05 (0.075
f 0.005)0.05 (0.077
* 0.005)(0.082
+ 0.005)(0.085
+ 0.005)(0.08 1
+ 0.005)(0.077
f 0.005)0.05 (0.073
f 0.003)0.05 (0.066
(0.0021* 0.0001) after 250°C expt.0.01 (0.03 1
it 0.0003)O.oool (0.0058 [Na]TlmoldmJ
Sampled[NaOHlhuokdm”
[Sb]&uol*dm”
it 0.0002) 7
Table 3-8: Total Antimony Concentrations as Measured over solid B (initially solid A, butconverted to solid B during the experiment) for Basic Oxidizing Solutions afterHeating to 250°C and Cooling to Room Temperature
Initial Final
f 0.0002) 60.05 (0.0029
f 0.0002) 40.05 (0.0027
f 0.0002) 60.05 (0.0034
f 0.0002) 30.01 (0.0027
[NaOH)/mol*dm.30.01 (0.0027
Time/d[Sb]T/mol*dm”
16
Table 3-7: Total Antimony Concentrations as Measured over Solid B (initially solid A, butconverted to solid B during the experiment) for Basic Oxidizing Solutions at 250°C
Initial
[7].
For oxidizing conditions, with the exception of acidic concentrated antimony solutions at roomtemperature, it is not certain what solution species are formed nor what solids are thermo-dynamically stable. There are good indications that one or more anionic species are formed in
50°C, the only previous useful resultsfor aqueous solution species appear to be for antimony(III) in neutral and basic reducingconditions to 200°C
antimony(III) and antimony(V),it appears that the behaviour of antimony solids and solutions is best understood for reducingconditions near room temperature. For temperatures above
spectrometry (ICP-MS). This indicated that the majoridentifiable component in the black solid was antimony, with a similar (atom fraction) amount ofsodium and small amounts of tin and lead (normal impurities in antimony compounds). Thus,the black solid is probably another unidentified sodium antimonate and not a product of areaction involving the titanium of the autoclave.
4. DISCUSSION
From a literature survey and our own preliminary results for both
tet.ra/metaborate fusion method, and analyzed byinductively coupled plasma-mass
Sb205. Asample was dissolved using a lithium
2OO”C, slightly longer equilibration periods were used.
Over the course of several equilibrations, small amounts of a very fine black solid were formedin mixture with the white solid B. An XRD analysis of the (mechanically separated) black solidshowed peaks at 0.372 and 0.413 nm that do not occur in the pattern for hydrated
NaOH(aq) at 200 and 300°C. For the experiments atmoldm-3
3.3.2.3),although in some cases the results showed considerable scatter. The procedure was carried outfor duplicate experiments using three different concentrations of sodium hydroxide. Experimentswere then carried out using 0.01
moldm~3), and re-equilibrated with a fresh sample ofthe aqueous sodium hydroxide. The hydroxide concentrations of the sampled solutions weremeasured by comparison with suitable concentration standards (as discussed in Section
HCl(aq). The total antimony concentrations and the solution sodium ionconcentrations in the samples were determined (Table 3-9). After the autoclave had been cooledto room temperature and the contents sampled, the solid was recovered, washed with very dilutesodium hydroxide solution (usually 0.0001
moldrn-3
mol.dms3) at 250°C for 48 h, sampled, left for an additional 24 h, sampled again, cooledto 25°C overnight, and sampled again. Each solution sample was immediately mixed with anexcess of 2
(0.001,0.005or 0.01
NaOH
cm3 aliquots of the finalsolution at 25°C for 8 to 132 d (using procedures similar to those described in Section 3.3.2.1).The total solution concentrations of antimony and sodium species from this third experiment arealso listed in Table 3-8.
3.3.3.2 The Solubility of Solid B as a Function of Temperature and Hydroxide IonConcentration
Samples of solid B were washed with water and equilibrated with aqueous
17
solid and solution from a third experiment were re-equilibrated with 10
25”C(c)* Experimental problems (loss of water from the autoclave during the run, possible sampling line problems).** Samples marked (a), (b) and (c) are from the same autoclave run; (a), (b )and (b’) are successive samplings on
different days at the same temperature; (c) is the sample taken after the autoclave was brought back to roomtemperature (usually on the same day).
*** Not measured.
f 0.0001) _*** at f 0.0026) (0.0010 3OO”C(b)
0.01” (0.0478 lo-’ at
300”C(b)<
lo-’ at 3OO”C(a)
<
25”C(c)0.0018 at
300”C(b)0.0056 at
3OO”C(a)0.0098 at
25”C(c)0.0065 at
300”C(b)0.0073 at
3OO”C(a)0.0074 at
25”C(c)0.0070 at
300”C(b)0.007 1 at
3OO”C(a)0.0128 at 0.0092 at
25”C(c)250”C(b)
0.0082 at
250”C(a)0.0089 at
25”C(c)0.0099 at
250”C(b)0.0067 at
25O”C(a)0.0112 at
25”C(c).---------------0.0083 at
sampled**
0.0038 at 250”C(b)250”C(a)
0.0039 at
25”C(c)0.0056 at
250”C(b)0.0059 at
25”C(c).------ .---------0.0072 at 250°C (a)0.0073 at
o.ooo2 at 250”C(b)25O”C(a)
0.0015 at
25”C(c)0.0005 at
25O”C(b)0.0009 at
25O”C(a)0.0023 at
o.2)10-5
0.0022 at
f 1O-5 (1.4 zt 1.6) o.2)10-5
(20.0 z!z 1.1)10-5 (0.7 f
It 0.0001)(11.1
f 0.0002) (0.0010 + 0.0001)
(0.0035 + 0.0010) (0.0012
i 0.0001)(0.0152
f 0.0009) (0.0018 + 0.0001)
(0.0137 f 0.0007) (0.0012
f 0.0001)(0.0114
+ 0.0009) (0.0021 + 0.0001)
(0.0147 f 0.0007) (0.0015
f 0.0001)(0.0118
* 0.0008) (0.0018 0.ooo1)
(0.0124 f f 0.0010) (0.0009 f 0.0001)
(0.0158 zt 0.0010) (0.0020
+ 0.0001)(0.0160
f 0.0010) (0.0024
+ 0.0001)
(0.0171
i 0.0008) (0.0025 zt 0.0002)
(0.0125 0.ooo9) (0.0036 f
L+Z 0.0002)(0.0150
+ 0.0009) (0.0036 f 0.0001)
(0.0145 f 0.0007) (0.0015
f 0.0002)(0.0117
f 0.0010) (0.0037 f 0.0002)
(0.0157 + 0.0009) (0.0039
--_______ __(0.0148
f 0.0002~-------------- (_0.0034 OXKIO6)f f 0.0002)
(0.0098 f 0.0007) (0.0040
f 0.0002)(0.0110
f 0.0006) (0.0038 f 0.0001)
(0.0104 It 0.0006) (0.0017
f 0.0002)(0.0106
f 0.0008) (0.0037
(~~~s70~ti~~--
(0.0135
‘_(o~oi._~o.._)__ f 0.0002)f 0.0003) ~0.0037 @0052 f 0.0002)f 0.0003) (0.0041 f 0.0002)
(0.0064 f 0.0003) (0.0040
zt 0.0003)(0.0061
f 0.0004) (0.0046 f 0.0002)
(0.0070 f 0.0014) (0.0042
f 0.0002)(0.0072
f 0.0004) (0.0042
25”C(c)
(0.0075
zk 0.0001) 0.0065 at A 0.0006) (0.0012 2OO”C(b)
(0.0108 +: 0.0003) 0.0077 at + 0.0009) (0.0044
2OO”C(a)(0.0148
f 0.0002) 0.0078 at f 0.0008) (0.0040 25”C(c)
(0.0133 f 0.0001) 0.0068 at f 0.0007) (0.0021
2OO”C(b)(0.0109
f 0.0003) 0.0098 at f 0.0009) (0.0044 2OO”C(a)
(0.0141 f 0.0003) 0.0075 at f 0.0009) (0.0045
[OH7
(0.0141
[Sb]&uol*dm”[Na]+uol*dm”
0.01*
Final
0.01*0.01*
------__-___0.0050.0050.0050.0050.0050.005---------0.010.010.010.010.010.01
300°C0.010.010.010.010.010.010.010.010.01
[NaOH]/mo&dm”200°C
0.010.010.010.010.010.01
250°C0.0010.0010.0010.0010.0010.001
3OO”C, or after Cooling to Room Temperature
Initial
18
Table 3-9: Total Antimony Concentrations for Solid B as Measured for Basic OxidizingSolutions at 200 to
Sb203 was[7], were done using the orthorhombic (valentinite) form of the solid rather than the cubic
(senarmontite) form that is reportedly stable at these temperatures, or else the
Popovaet al.
Sb203from 15 to 50°C (for details, see text).
In general, agreement between the results of the different studies is excellent-much better thanmost studies of oxide solubilities. Almost all the measurements, again except for those of
antimony(II1) in aqueous solution in equilibrium with
mol-dms3
Figure 4-l: Total concentrations of
loo
[OH-] 10” lo* 10”lo4 1o-5 1o‘7 10”
lOA-X
X X
n X Xx X
I I I I I I I I
?aX
X
[73VAS/SHO]% 10”: 50°C ti
0[73VASISHO]E +
m[48TOU/MOU]5 35°C X
-au _m+m
$1
this work[73VAS/SHO][52GAY/GAR]lo-*-
[39SLO]
[73VAS/SHO]
25°C
0 [1883SCH]W
lop
15%
[2], all these measurements were done using solutions attemperatures between 15 and 50°C (Figure 4-l).
Schulze [7] and a
single experiment by Popova et al. [2-71 (cf. Table 2-l). Except for the work of
Sb203 in water and basic aqueous solutions have been reported previouslyby a large number of authors
SbzOs, listed in Table 3-l. Theresults are fairly scattered, and there are no evident trends of solubility with temperature or baseconcentration. The concentrations of base as measured in the equilibrated samples are probablyless reliable than the total solution concentrations of antimony, and the “final” measuredhydroxide ion concentrations were not used in the data analysis.
Solubility values for
Antim0nvUI.I)
The following is an analysis of the solubility data obtained for
25”C,should be regarded as, at best, highly speculative.
19
basic solutions at room temperature, but prior to the present study there has been essentially noinformation as to whether this remains so at higher temperatures. The previous solubilitymeasurements are suspect, because interconversion of several solids appears to be possible withchanges in temperature. Any calculations of the speciation of antimony under oxidizingconditions based on the earlier solubility studies, particularly for temperatures greater than
a3(H20(aq)))(a2(OH-) / a2(Sb(OH),) Kz=
Sb(OH)s(aq). The equilibriumconstant for reaction 4.2 can be written as
HzO(aq) is equal to1.0, as is the activity coefficient of the neutral aqueous species, Sb(OH)s(aq). In the present analysis, we have assumed that the activity for
HzO(aq) and of
4- 1.
The equilibrium constant for reaction 4.1 depends on the activities of
[7]. It is the initial base concentrations that are shown in Figure Popova et
al. [4], and
[S] were stated to be initial concentrations, and it wasassumed that this was also the case for the experiments of Tourky and Mousa
2Sb(OH)i(aq) (4.2)
Many of the reported measurements list only the initial base concentration, not the finalconcentration. Where the only major aqueous antimony(III) species is a neutral species, or if theinitial base concentration is much greater than the measured total concentration of antimony insolution, the data analysis is not affected. However, if the initial base concentration and the finalantimony concentration are of the same order of magnitude, the final base concentration may bereduced from its initial value from the stoichiometry of reaction 4.2. The reported baseconcentrations of Gayer and Garrett
* 20H-(aq) 3H20 + SbzOs(c) +
2Sb(OH)3(aq) (4.1)-T 3H20 Sb203(c) +
Sb203 in neutral and basic solutions have generally been interpreted interms of two equilibria:
SbzOs were to precipitate from solution in a reactor system, the solidwould probably be the less-stable valentinite.
The solubility results for
Popova et al., overestimation of the calculateddifference in the Gibbs energies of formation of the two polymorphs, or systematic errors ineither study. Regardless of the cause, most of the following data analysis has been done withoutdifferentiating between the solid actually used in each study. Based on the diffkulties we haveencountered in attempting to synthesize senarmontite from aqueous solutions, it seems that in theunlikely circumstance that
al’s values are, if anything,greater than those found in the present study. This result could indicate a small amount ofvalentinite impurity in the senarmontite sample of
Popova et [7] at 200°C over senarmontite should be systematically lower than our measured
values over valentinite by a factor of 2.65. However, Popova et al.
antimony(III) concentrations of
k.Lmol-’ differencefound in a reanalysis of available heat capacity, enthalpy and transition thermodynamic data,summarized in Appendix A.
Calculations based on the same assessment suggest that the total
[41-43], or the 5.7 kJ.mol-’
that has been suggested in several standard compilations kJ.mol-’ at 25°C. This is a smaller difference in stability than the 7-9
[3] carried outmeasurements over many months to compare the solubilities of the two forms at roomtemperature, and found a difference of only 0.36 in the logarithm (base 10) of the values; i.e.,approximately 2
20
synthesized under conditions known to yield the orthorhombic form. Bloom
15_3~oc, the average heat capacity ofreaction between 15 and 300°C) were calculated using data from all the studies discussed above;
A&, avs (i.e., A& A,S” and A,G”, 4- 1) of J-K-‘.mol’) and T is the temperature in kelvin. Parameter
values (Table
ln(T/298.15)} (4.4)
where R is the gas constant (8.31451
avg -TA,C, AS”&(T-298.15) - avg (A&, AGO25 + { l/RT) exp -(
A$ravg, for reactions 4.1and 4.2.
Equation 4.3 was combined with the equation
K =
300°C, A& for 15 to AS” at 25°C and an average value of ArGo, Sb203 solubility data to generate thermodynamic quantities
200°C; this is true even if theirvalues in water are compared with our values in basic solutions. This could also indicate that oursolutions were undersaturated, but we cannot be certain without further experimental work.
Attempts were made to use the
[7] at Popova et al. than those measured by 200 to 300°C are less
SbzOs-Hz0 system in which equilibrium wasreached starting from supersaturated solutions.
A comparison of our solubility measurements with other measurements reported for temperaturesbetween 90 and 200°C is shown in Figure 4-2. In general, our measured solubility values from
3OO”C, and these solutionsapparently remained supersaturated when they were cooled to a temperature where the solubilityis less. They had not returned to equilibrium in the 1-to-5-day period before sampling. No otherexperiments seem to have been reported for the
Sb(OH)i anion at low temperatures.
Values at 25°C from the present study are systematically greater than those reported previously.Our measurements represent solutions that were saturated at 200 to
[4,6] used hydroxide solutions with concentrations greater than thoseconsistent with the Davies equation (i.e., IO.2 to 0.3 M). Unfortunately, such solutions alsohave provided the best evidence of formation of the
K; [OH-] (4.3)
Some of the studies
KT + =
PWHMKWOHhl + [Sb]r =
antimony(III) in solution can beexpressed as:
fust approximation, the total concentration of
[45], that do not contain termsspecific to the ion) is adequate.
Therefore, to a
[44] (orextended versions of that equation, such as the Davies equation
Debye-Htickel equation
the
ionic medium in the solution is sufficiently low that the
Q,~(~B)&B-, is equal to 1 .O, provided y is a molar activity
coefficient. The ratio of the activity coefficients, where terms ¬ed [A] are molar concentrations, “a” is an activity and
124f315f 505k4.6 -4 f 33
12.142f54 -134 f 3.3
f 128
24.9
56+204 -130 f 3.0AI 40
12.946f64 -151 + 3.5
f 76
24.9
Ik 120 -23 f 2.0 35 f41
12.9Zk 66 -155 f 3.5 47
/PK-‘*mol-’24.9
15-300°C4c,
/JK’~mol~lMrnol-’4-s”4G”
I 1 .O Mreaction (4.1)reaction (4.2)
all data [OH-] IO.3 Mreaction (4.1)reaction (4.2)
Sb203
all data reaction (4.1)reaction (4.2)
all data [OH-]
4- 1: Calculated Thermodynamic Quantities for the Dissolution of
K2. The calculations were repeated, but in an attempt to minimize activity coefficienteffects on the derived thermodynamic quantities, only measurements for solutions with totalionic strengths less than or equal to (a) 1.0 M and (b) 0.3 M were used.
Table
Sb203 at temperatures from 90 to 300°C. The “best-fit” line for 25°Cbased on the low-temperature data is shown for comparison, as is an arbitrary linewith a slope of 1 .O.
reported “initial” hydroxide ion concentrations were adjusted to “final” values using the “fitted”value of
moledmd
Figure 4-2: Solubility of
,’
[OH-]
,’:
,’.’
slope=1
:..----arbitraryline:thiswork
3ooc
0ltlkwork250% A
thiswork2w”c
[75POPKHO]200°C
[1883SCH]
[75POPlKHO]
100°C 0
0 90°C
22
SbzOs “antimonic acid” would beexpected to have a high solubility in neutral and basic solutions, with the formation of anions
“NaSb(OH)&)” and Other SodiumAntimonates
On the basis of literature values for acidic solutions, (hydrous)
Antimonv(V)
4.2.1 Rationale for the Measurements Using
moLdma
Figure 4-3: Calculated total solution concentrations of Sb(III) as a function of temperature andhydroxide concentration.
4.2
[OH-]
10"lo4lo9lo-'10"10"
IIIIII
lo-'I
L
< 0.3 M.Sb2Os solubility measurements (without regard to the particular polymorph) for hydroxide ionconcentrations
I 0.3 M, the calculated totalconcentrations of antimony species at various temperatures, as shown in Figure 4-3, wereestimated using the constants from an equal weight “least squares” treatment of all available
10s2 M for all basic solutions with [OH-] IO.3 M, more than anorder of magnitude greater than at room temperature.
Based on the parameters for hydroxide ion concentrations
Sb203 is near 300°C the total equilibrium concentration of aqueous antimony species
in equilibrium with
2OO”C, it then remains constant, or begins to decrease only slightly between200” and 300°C. Even at
Sb203in neutral to basic solutions increases by more than an order of magnitude between roomtemperature and
Sb(OH)s(aq) is much better defined, and it is apparent that although the solubility of 200°C than for those near room temperature. The temperature dependence of dissolution to
form 2
Sb(OH)i is probably less important for temperaturesAr(4.2$‘25 are
highly correlated. Indeed, it appears that avg and A~4.&,
ill-defined, especially at higher temperatures. The calculated values of
Sb(OH)i are still
23
It appears that the thermodynamic quantities related to the formation of
(NaSb(OH)6) in Basic Solutions
As Figure 4-4 shows, at 25°C solid C appears to be more soluble (i.e., less stable) than solid A.At 75°C solid A is unstable with respect to solid C, and the latter is much more soluble than at
NaSb03*3H20(s)
the sodium salts used in the present study.
4.2.2 Solubility of
those for 181. Thus, the results would be no easier to interpret than
[8,32]. Further, the extent of its hydrationchanges markedly with temperature [
171, and in contact with sodium hydroxide solutions the simple oxide would beconverted, at least in part, to a sodium salt (or salts)
[8,
Sb205 mightalso have provided useful information, but that solid is not particularly simple to synthesize in apure form
[29,32]. Study of the solubility of (hydrated) OH-). However, solid B is still not well-characterized, although apparently similar material hasbeen synthesized by other methods
Na+ (and/or25”C, except at high solution concentrations of
75”C, and then to solid B at higher temperatures. Thus, neither of the (presumably) simplecompounds A or C is suitable for the study of solubility as a function of temperature.Nevertheless, solubility measurements using the simple sodium antimonates (A and C) provide asatisfactory starting point for a study of Sb(V) in basic solutions at low temperatures.
The use of the more complex, probably non-stoichiometric, solid B is more problematic. Itseems to be reasonably easy to synthesize, and analyses suggest that variation in thestoichiometry is not extensive in basic solutions, and is not affected substantially by alteration inthe washing or ripening procedures. Also, as discussed in a later section, solid B is stable withrespect to solids A and C even at
2.5H20(1) (4.7)
the concentration of antimonate in solution would vary as the inverse square of the concentrationof aqueous sodium hydroxide. In each case, it should be possible to relate measurements atdifferent temperatures to changes in stabilities of the aqueous antimony(V) species as a functionof temperature.
However, the present work has established that in dilute sodium hydroxide solutions the simplesodium antimonate (solid A) is probably converted to solid C at some temperature between 25°Cand
OH(aq) + OSSb205(aq) + Na+(aq) + * NaSb03.3H20(s)
4H20(1) (4.6)
the concentration of antimonate in solution would vary directly as the concentration of aqueoussodium hydroxide. If the reaction was
SbOi(aq) + Na+(aq) + * NaSbOs.3H20(s)20H-(aq) +
3H20(1) (4.5)
has an equilibrium constant such that the concentration of antimonate in solution should varyinversely as the concentration of aqueous sodium hydroxide. If the reaction was
SbOj (aq) + Na+(aq) + * NaSb03.3H20(s)
NaSb03*3H20), is reported to be stable and is only sparingly soluble. Therefore,measurements of the solubility of this solid as a function of aqueous sodium hydroxideconcentration were expected to provide a method for investigating the behaviour of aqueousantimony(V) species in neutral and basic solutions. For example, the dissolution reaction
NaSb(OH)h(s) (probably better writtenas
24
containing Sb(V). The easily purified monosodium salt,
Sb(V) concentration forsolubility measurements of sodium antimonates in basic solutions at 25 and 75°C.Lines from the least-squares fits to each set of values are also shown.
However, as can be seen from Table 3-5, the experimentally measured final hydroxide ionconcentrations were substantially lower than the initial concentrations of the aqueous sodium
moLdme
Figure 4-4: Sodium ion concentration (M) as a function of total
‘\
10”
[Sb],
-\‘\
‘\
A C 75°C‘\
25%0 C
75%
B25’C
B
A25”C
n
v
OH(aq) concentrations.SbOj (aq) (i.e., total Sb) and, where reasonable, the experimentalNa+(aq),
[46]. The total ionic strength was calculated fromthe experimental
A=0.509 at 25°C and A = 0.564 at 75°C
Ic.5) (4.9)
with
I?( 1 + 1.5 = -A logloe
Debye-Htickel equation:
(4.8)
The activity coefficients were estimated for these relatively dilute solutions using an extended
SbO;*a aNa+ =Ks,
- 1 (closer than if it is assumed that A andC are the same solid). We are unable to decide, based on present evidence, whether solids A andC are essentially the same or if they are two distinct phases.
From the experimental data listed in Table 3-5, the solubility (activity) products of solid C at 25and 75°C have been calculated based on reaction 4.5
(NaOH(aq)) concentration. This decrease is approximatelylinear, and the slope of the line in Figure 4-4 is close to
Na+
25
25°C. At both 25°C and 75°C the antimony concentration of solutions in contact with solids Aand C decreases with increasing
25OC are actually more concordant with the literature values than those for solid A. This mightNaSb(OH)b. Our values for solid C at[39] for [36] and Blandamer et al.
[38] are markedly less negative than those from othersources, and it must be presumed that their analyses were affected by the presence of colloidalantimony solids. For 25°C our values for solid A are in fair agreement with, but slightly lowerthan, those of Tomula
log&, values of Urazov et al.
[39]
The
Hz0 -4.159NaSb(O& 80[381
1391present work
1391[381[391[361[391[381[361
1361present workpresent work
fo.105
io.034fo.074
fo.050
fo.146
fo.046
H20 -3.046NaOH(aq) 10 -4.094
75
H2O -4.24175
Hz0 -4.51870
Hz0 -3.64850
H20 -4.77650
H20 -4.83 135
H20 -5.01733.5
H20 -4.55 125
H20 -5.10025
NaOH(aq) 11 -5.07225
NaOH(aq) 4 -5.25625
Hz0 -5.32525
NaSb(OH)a
18
NaSb(OH)6C
NaSb(OH)hNaSb(OH)6NaSb(OH)bNaSb(OH)dNaSb(OH)eNaSb(OH)eNaSb(OH)e
NaSb(OH)eAC
NUlIlber Avg. Sigma 95% Confid. ReferenceMediUmT/“C
NaSb(OH)e
Solid
NaSbOsor
3-5), and comparison values from the literature, are listed in Table 4-2.
Table 4-2: Values of the Solubility Product for Solids Nominally Hydrated
NaOH(aq) were sufficiently great that the final concentrations of these species could be assumedto be equal to the initial concentrations without introducing undue errors.
The average values of the solubility products (based on the measurements reported in bothTables 3-4 and
Na+(aq) and OH-(aq) concentrations were not determined. However, the initial concentrations of
log&,.
The solubility measurements listed in Table 3-4 were also used to determine values of thesolubility products for solids A (at 25°C) and C (25 and 75°C). In these experiments, the final
Na+(aq).The extra uncertainties introduced by this procedure are less than 0.01 in
NaOH(aq), the ionic strength was assumed, for thepurpose of the activity coefficient calculation, to be equal to the final concentration of
(N2 gasflushing prior to closing the vials, and double containment) taken against this happening, or smallamounts of acid may have been leached from the plastic bottles despite preconditioning bysoaking in basic solutions and water. When the measured hydroxide ion concentration wassubstantially lower than that of the initial
CO2 despite the precautions
NaOH(aq) was low. The reasonfor this is not clear, although the antimony solids themselves are unlikely to have been theprimary cause (sample blanks with no solid showed the same tendency (Table 3-5)). Thesamples may have been contaminated with atmospheric
26
hydroxide solvent, especially when the initial concentration of
[4] is very similar to the solubility of solid B in dilute basic solutions.mol.dm-3) reported by
Tourky and Mousa Sb205 in water at 35°C (0.00027 mol.dm-3 at 25°C. The solubility of
NaOH(aq) concentrations of approximately0.08
moldm-3 at 75°C. Solid B would becomeunstable with respect to formation of solid A for
moldm-3 at 25°C and 1 NaOH(aq) concentrations of
approximately 0.2
13]), and furtherassuming that the antimony solution species over all the solids is the same, solid B would beexpected to become unstable with respect to solid C only at
[ 4.2.3.2), assuming no changes in the predominant antimony species in solution with
increasing hydroxide ion concentration (probably an oversimplification
Nai.sH.&b206(s) would be expected to decreaseas the square root of increasing sodium hydroxide concentration (i.e., more slowly than in thecase of the simple sodium antimonate). If this stoichiometry is accepted for solid B (cf.Section
L4,75
(4.10)
(4.11)
(4.12)
Thus, the antimony concentration over hydrated
= ,-_p.25 aNa+aSbO;
0.75
K, is defined by
For a = 0.75, this reduces to
(5-4a)H20(1)
and the equilibrium constant
2SbOj(aq) + 2aNa+(aq) + *(2-2a)OH(aq) Na2a[H(H20)]2_2aSb206’H20(S) +
4-4), and this suggests that solid B is more stablethan solid C under these conditions.
One of the simplest descriptions of the dissolution equilibrium is:
NaOH(aq)(Table 3-5) is less than that of solid C (Figure
mol.dm-3 Na2,[H(H20)]2_2,Sb206.H20 (a = 0.75) at 25 and 75°C in 0.003 to 0.04
= 0.75 in Basic Solutions
4.2.3.1 Comparison of the Solubility with Other Solids at 25 and 75°C
The apparent equilibrium concentration of antimony in aqueous solution over
Na2,[H(H20)]2_&b206.H20, a
Sb(OH)&aq)),and that the speciation is not dependent on the hydroxyl ion concentration in these solutions.
4.2.3 Solubility of
SbOj(aq) or Na+(aq) and a singly charged anionic monoantimonate species (i.e., NaOH(aq) are consistent with the assumptions that these solids dissolve to formmoldm~3
[36,39] and in 0.003 to0.1
[36,39] involved heating toslightly higher temperatures. The values of Blandamer et al. for 70 and 80°C are similar to ours(for solid C) within the uncertainties.
The similar results for the solubility products determined in water
27
be explained if the methods used to synthesize or treat the solids
(KJ is strongly dependent on the value of a. If a isnot independent of temperature, a set of values for the solubility product at different temperatureswill not yield useful information concerning changes in stability of the aqueous antimony speciesas a function of temperature. Therefore, values of the activity product and a were first calculatedfrom the results at each temperature for which measurements were done (Table 4-3).
25O”C, and even more so at 300°C(Figure 4-5).
The calculated value of the activity product
2OO”C, but decreases slightly at
Na+(aq) and OH- concentrations. If anything, the variation was less,especially near 250°C. The results from samples taken on successive days at 250°C (i.e.,samples from the same autoclave run) are fairly consistent; the values from successive runs underthe same conditions are somewhat less so. The total solution concentrations of antimony oncooling the autoclave to 25°C are lower than the high-temperature values, but generally greaterthan from solutions equilibrated for longer periods of time (Table 3-5). If all the results areconsidered together, regardless of the actual base concentration, the solubility of solid B in basicsolutions increases from 25 to
300°C varied onlyslightly with changes in
200 to
Na20r[H(H20)]2_2,Sb206.H20 from 25 to 300°C
As for the results at lower temperatures, the solubility of solid B at
Na+(aq) were consistently greater than the total solution concentrations of antimony, and thefinal measured hydroxide concentrations were low (Table 3-6). This is consistent with simplepartial dissolution of the solid C (or A) from the mixture, with some portion of the sample ofsolid C never coming into contact with the bulk of the solution. At present this is the bestexplanation we have found for the results listed in Table 3-6. However, there then is no apparentreason why the duplicate experiments over 25 days should have given essentially identicalresults, and the same applies to the pair of 42-day experiments. The same difficulty would ariseeven if we were to assume that equilibrium cannot be attained with one of the pure solids withinthese periods of time.
4.2.3.2 Solubility of
XRD patterns of the residual material show that both initial solidswere still present after all of the experiments, although qualitatively the ratio of solid C to solid Bappears to have decreased after extended equilibration times. The final solution concentrationsof
with water would generate a solution in equilibrium with both. The totalantimony, sodium and hydroxide concentrations would then be fixed (for a specific value of a).These concentrations could then be compared with the values calculated from the two solubilityproducts. Alternatively, if the solubility differences were large, one of the solids could becompletely converted to the other. However, contacting water with mixtures of solid B andsodium antimonate (solids A or C) at 75°C even for more than 200 days apparently did not resultin establishment of equilibrium between the solids, probably in part because the samples werenot agitated continuously. The
75”C, B and Aat 25°C) in contact
28
It was hoped that long-term equilibration of mixtures of two solids (B and C at
A&, for reaction 4.10 (assumed to be independent oftemperature) were determined using a non-linear least-squares fit (Table 4-3). Using the fitted
AS and 25”C, and values of logtoK, for
J.K“.mol-‘.
Using the results for all five temperatures, a single (T independent) value for a, a value for
If: 27) -(218 A& =J.K-‘.mol-’ and-I 9.5) A&S = (38.4
-
* The other fitted parameters are & 0.02)- (0.7 1 all 36
024) -3.37f zk 1.01) (0.14 -(1.26 f 0.02) -3.22
300 6f 0.08) (0.76 -(3.45
f 0.44) -3.16250 12
f 1.81) (0.77 -(3.42 z!z 0.02) -3.71
200 4f 0.10) (0.76 -(4.05
z!z 0.11)75 8
-(4.52 f 0.03)f 0.13) (0.62 -(4.10
~ogloKx@[email protected]),
Measurements25 8
&dehh&T/“C Number of
logloKX(25”C) and a Calculated from the Results at all Temperatures*
(logroK,) with Values of aCalculated from the Experimental Results for Each Temperature and from Values of
Nah[H(H20)]2_2aSb206.H20
It 0.03) found by neutron activation analysis of three separatelyprepared samples of solid B.
Table 4-3: Activity Products for
0.0075-O.O098,25O”C 0.0005-0.0112, and 300°C 0.0065-0.0128.
Except for the value calculated from the 300°C solubility measurements, a does not differ greatlyfrom the Na:Sb ratio of (0.75
l-O.O42,2OO”CO.OOlO-O.O49,75”C 0.0001 (mol.dm-3): 25°C Nat.sH&b&I&o). Hydroxide ion
concentration
300
Temperature (“C)
Solubility measurements for solid B, a mixed oxide of antimony(V) (hydratedpyrochlore-structure sodium salt,
250 200150100 500 o.ooo1
29
Figure 4-5:
pH 12 even at 25°C;[lo] would overestimate the solubility of antimony
at
Sb203 in basic oxidizing solutions;the nature of insoluble antimony solids in basic oxidizing solutions is probably stronglydependent on the nature of the solutes, especially simple cations;for modeling total antimony concentrations in oxidizing solutions, the use of antimony(V)species as discussed in Baes and Mesmer
[9,47], it is probable that alkali metal antimonates are stable relativeto potential-pH diagrams
Sb(OH););the temperature-dependence of the solubility of other antimony(V) solids in basic oxidizingsolutions would be expected to change similarly (again based on the reasonable assumptionthat the solubility changes are primarily controlled by changes in the stabilities of theantimony solution species);metastable antimony(V) solids can persist for extended periods of time in contact withoxidizing basic solutions at temperatures at or below 75°C;although Sb(V) solids are generally not shown as having a region of predominance in
(SbOj or 25O”C, probably primarily reflecting changes in the stability
of the anionic antimony solution species
2OO”C, anddecreases at temperatures above
this pyrochlore-structure sodium antimonate increases from 25 to f 0.03) in basic aqueous solutions at 250°C;
the solubility of Na~[H(H20)]2_2,Sb206.H20 (a = 0.75
NaOH(aq);simple sodium antimonate is converted to a hydrated pyrochlore-structure sodium salt,
mol.dm-3
25O”C, and probably to 300°C;comparison of the solubility product of sodium antimonate(V) as determined in basicsolutions with values reported in the literature suggests that the same antimony solutionspecies is predominant in oxidizing solutions at 25 to 75°C from neutral solutions tosolutions containing between 0.01 and 0.1
Sb203 were to precipitate from solution in areactor system (unlikely because the concentrations of antimony solution species are toolow), the solid would probably be the less-stable valentinite;all of the antimony(V) solubility measurements are consistent with formation of amonoanionic antimony species in oxidizing basic solutions for temperatures from 25 to
Sb203 in basic solutions increases from 25 to 200°C and probably decreasesslightly between 200 and 300°C;it seems that in the unlikely circumstance that
5. CONCLUSIONS
Based on the work described above, we draw the following conclusions:
l
l
l
l
l
l
l
l
l
l
l
the solubility of
30
value a = 0.71, none of the calculated values for the antimony concentrations differ from thecorresponding experimental values by more than a factor of 2.1. Considering the possiblesampling problems at high temperatures, the apparently slow approach to equilibrium at lowtemperatures, and the fact that the calculations were done using the simplification of assumingthat the heat capacity of reaction is independent of temperature, this agreement is reasonablygood.
18 305-310 (1973)).Khim. l8,
161-164 (1973). (English translation from Zh. Neorg. Antimony(III) in Alkaline Solutions by a Solubility Method”, Russ. J. Inorg. Chem.,
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Sot. 74,2353-2354 (1952).
25”“, J. Amer. Chem. [S] Gayer, K.H., Garrett, A.B., “Equilibria of Antimonous Oxide (Rhombic) in Dilute
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7. REFERENCES
Totland did the ICP-MSanalyses and, with P. Robinson, the neutron activation analyses. D. Guzonas and C. Stuartprovided useful comments on a draft of the report. We also wish to thank D. Guzonas for manyuseful discussions.
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ACKNOWLEDGMENTS
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2 0.00005
31
all solids used in the present experiments would be expected to generate total antimonyconcentrations of
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36
sb20, -838.9 -829.2 -864.7 -829.2 -829.3 -829.1
* The year of compilation was 1964.
ortho. -615.0 -626.5 -631.8 -624.7 -626.6 -624.7 -626.3Sb203,-634.4 -641.0 -626.8 -632.2 -634.3Sb203, cubic -623.4
WIr491WI[71[4111421*r91/ ref.&“/kJ-mol’l
Compound
A2-A6.
Table A2: Literature Tabulations of Gibbs Energy of Formation Values for Antimony(III) andAntimony(V) Oxides at 25°C
Antimonv(III) and Antimony(V) Oxide Solids
Tabulated values from the literature for simple oxides of antimony (III) and antimony(V) arelisted in Tables
!%O; -274.1
* The year of compilation was 1964.
A.2
SbOi -345.2 -340.19 -342.9 -339.5 -339.74 56.8 55.2Sb(OH)3. -645.1 -644.7 -647.3 -644.7 116.3 125.5 192.9SbO’ -175.7 -177.11 -179.6 -175.8 -175.64 -7.1 22.33
r4911711411[421*[49117114111421*191/ref.So/J-C’moP at 25°C
SpeciesA&?VlcJ-mo~’ at 25°C
[50,51].
Table A 1: Literature Tabulations of Chemical Thermodynamic Values for Antimony AqueousSpecies
[49] were based onthe electrochemical studies of Vasil’ev et al.
SbO as assessed by Past SbO+ and high-
temperature solubility results. Values for [7] were based on their own Popova et al. SbO+, the values for the solution species from
H20(1). Except forH20 have been treated as identical, and reported
chemical thermodynamic values have been adjusted using the values for
[48]. In all cases, proposedspecies differing by integral multiples of
[41,42] is traceable to Thomsen kJmo1“ AfH’(HSb(OH)b) = -1478.6
Simnle Aaueous Ions and Hvdrolvsis Snecies of Antimony
Tabulated values from the literature for simple aqueous ions and hydrolysis species ofantimony are listed in Table Al. A value for one additional species should be mentioned:
Snecies and SelectedOxide Solids
A.1
Thermodvnamic Data for Aaueous Antimonv
37
Appendix A: Literature
Ci(Sb203, orthorhombic) is not given explicitly.[7]. A value for
[54].The latter set of tables was not available to the authors of the present report.Based on the authors’ equation for the temperature dependence of the heat capacity oftransformation
[58], and not in the tables of Glushko et al. Barin et al. [43] and Barin [59]. The problem may
only be in the tables of
[SS].Probably the value for the orthorhombic form from Gorgoraki and Tarasov
Barin et al.
x 104.4 111.85101.38 111.9 x+7.3 101.38 101.38 111.8 101.4
The year of compilation was 1964.As reported by
lll.S$ 104.61431[561[55.l[521rlw[41][541?[421*
Cg/JKbd-’
Sb&. ortho.Sb203, cubic
I ref.
A5: Literature Tabulations of Heat Capacity Values for Antimony Oxides at 25°C
Compound
i 8.4
Table
f 4.2 110.45123.0 123.0 141.0 123.0 134.6
124.9 125.1 124.9 125.1
1431122.2 132.6 132.7
[Sal[SSI[49][521VIS”/JdlllOl-l
[58].Barin et al. t As reported by
Sb2@ 125.1 125.1 125.1
* The year of compilation was 1964.
SbzO3, ortho. 123.0 141.0llO.& 132.6 132.4Sb&, cubic
[41][541T[42]*ref.I compolmd
value for the orthorhombic form.
Table A4: Literature Tabulations of Entropy Values for Antimony Oxides at 25°C
0 Probably the [58].Barin et al. $ As reported by
[57].Knacke Barin and t As reported by
sbzo5 -971.9 -971.9 -1007.5 -1007.5 -971.9 -971.9 -993.7 -1007.5 -971.9
* The year of compilation was 1964.
-708.S5 -708.8 -706.9 -708.8 -701.6 -708.8 -708.6Sbz03, ortho.-708.85 -715.5 -720.4 -709.4 -711.6 -716.1 -720.3Sbz03, cubic -720.3
1431WIWI[4911521r711411Kw[531?[421*/ ref.A&P/kJmol-’
Compound
Antitnony(III) andAntimony(V) Oxides at 25°C
38
Table A3: Literature Tabulations of Enthalpy of Formation Values for
[64],S’(Sb203, cubic, 25°C) and CODATA consistent auxiliary data Ci(Sb203(s), T),[42], is used. From this, the transition enthalpy and the selected expressions for
kJ.mol-‘, selected by the U.S. National Bureau of StandardsAfH’(Sb203,
orthorhombic, 25°C)) = -708.55
25”C), orrecent experimental values for either the orthorhombic or cubic form, the value
A$I“(Sb203(s),
0.092918T
In the absence of a recent, detailed analysis of literature values for
T/K)/J.K-‘.mol“ = 66.296 + Cr(Sb203, cubic,
C,(Sb203, cubic, T) is:[61]), the consistent function for kcal.mol-’ kJ.mol-’ (i.e.,
1.0
J.R’.mol-’ can be selected (the uncertainties are estimates). Fromthese values, and the assumption that the enthalpy of transition at 606°C is 4.184
& 10) S’(Sb203, cubic) = (114 J-K-‘.mol-’ andf 10) Ci(Sb203, cubic) = (94 [60], are accepted, and the values
S”(25”C) for the orthorhombic form, based on the measurements ofAnderson
C;(T) and
J.KW1.mol-’ less for the cubic form than for theorthorhombic form Although there are doubts as to the absolute values for the heat capacitiesand derived entropy values, the differences should be approximately correct. Therefore, thevalues of
S”zoc is 8.6 K-‘.mol-’ near 25°C; also, Ci(Sb203, orthorhombic) by
7.2 J Ci(Sb203, cubic) is systematically less than [59] showed that
[63], the entropy values derived from the measurements are also likely to beincorrect.
Assuming that the compounds were correctly prepared and characterized, Gorgoraki and Tarasov
[59] are incorrect Gorgon&i and Tarasov. If the heat capacities of Gorgoraki and Tarasov
there may have been a systematicproblem in results of
[62] have suggested that Best& 58% greater than those of Anderson. However, based on results for
arsenic oxides, Chang and
C;(Sb203, cubic)throughout the temperature range of the measurements. The values of the latter authors for theorthorhombic solid are
(7&2)% greater than C”,(Sb203, orthorhombic) are [59]. This work
showed that the values of Gorgon&i and Tarasov Sb203(s) as measured between -208 and 27°C by
[41,49,54,56] have preferred the heat capacity values for both formsof 21.2”C. Some compilers
[60] measured the heat capacity of the orthorhombic form from -2 13.4 toSb203(s) have been carried out using the orthorhombic form of the
solid. Anderson
[61].Although the cubic form is more stable at low temperatures, the orthorhombic form is metastableover a wide temperature range, and most of the earlier chemical thermodynamic measurementsreported in the literature for
Sb203, two solids exist-the orthorhombic valentinite and the cubic senarmontite
[60]; probably the sample used for the measurement wasprimarily the orthorhombic form.
For
From the same data t
1521[43,521 t[43115611411 t
Reference
7:9 71.575.31 97.4992.048 66.107
114.01 8.318 -13.435141.33 -3.732 -20.112
lO%/(T/I#b C
+lO%(T/K) + @JK’mor’ = a
SbzosSb& ortho.
SbzO3. cubic
39
Table A6: Temperature-Dependent Heat-Capacity Values for Antimony Oxides
Compound
alkali metal antimonates,NajSbO4, and
chemical thermodynamic values for a wide variety of other anhydrous SbO$(aq)) was not established. The entropy of
kJ.mol-‘. However, the nature of the aqueous antimony species so formed(assumed by the authors to be
f 0.4) [68] as being
(18.2 Na3Sb04 in water was also reported
1691
The enthalpy of solution of
& 5.4)f 1.5) (76.8 f 13) (20.8 [691
(183 f 2.2)f 0.4) (35.6 f 8) (7.0
I681(132
if391f 5.8)
Reference
-(73.1 f 0.8)f 17) (9.9 f 0.4)
(209 -(4.8 (156rt9)f 4)
10%!/(T/K)2A B C
(65
+lO%(T/K) + = A C~JK’*moK’
Na$bO,NaSbOJ
MSb03solids. Their values are summarized in Table A7.
Table A7: Temperature-Dependent Heat-Capacity Values for Alkali Metal/Antimony(V)Mixed Oxides
Compound
(298.15-4OO”C) of the same compound and of a series of [68,69]
determined the heat capacity
[42,67]. This presumably was done because the reviewers decided that the value was unreliable.Certainly, Mixter’s reaction product was not well characterized. Kasenova et al.
[66], but was dropped from subsequent editions of these tableskJ.mol-‘), was recalculated for the U.S. National
Bureau of Standards Circular 500 kcalmol’ (-1473 AfH”(NasSb04,25”C) = -352
NasSb04 based on a determination of the heat ofreaction of “pulverized antimony and sodium peroxide”. A value based on this heat of reaction,
[65]reported a value for the enthalpy of formation of
Sb205.
A.3 Chemical Thermodynamic Measurements for Mixed Oxides Containing Antimony
Calorimetric data are extremely limited for the Sb(V) salts and/or mixed oxides. Mixter
25’C areextrapolated. There are substantial questions with respect to all of the primary experimentalchemical thermodynamic data for
Sb205 for temperatures greater than [60], even if the compiled heat capacity is quite
different. All chemical thermodynamic values for S”(Sb205,25’C) based on Anderson’s work
[60] for aslightly hydrated oxide. Most compilers have proposed approximately the same value for
[54,56] (original sourcenot known to the present reviewers) are greater than that measured by Anderson
Sb205 in at least two compilations Ci(Sb205) are derived from these measurements.
The values for the heat capacity of J.K-‘.mol-’ for
J.K’.mol-‘) for 17°C.All tabulated values near 118
cal.K-l.mol‘l (217.7 (Sb20ss2.224H20) = 52.03 ) and J~R’~rnol’ cal.K“.mol-’
(130.1 17H20) = 3 1.10 Ci(Sb205.0.3
Sb204). After allowingfor the lower oxide, Anderson reported
Sb205 mixed with a lower oxide (presumed to be [60] did his low-temperature heat-capacity measurements using
samples of hydrated
181, and the products of the combustion reactions were not wellcharacterized. Anderson
[ Sb205 yielded
a partially reduced solid [65] to prepare anhydrous
kJ.mof’.
It is probable that the synthesis method used by Mixter
kJ.mol-* and the difference in the Gibbs energies of formation between the two formsis 5.7
AG”(Sb20s, orthorhombic, 25°C) is-626.39
kJ.mol-‘. The value of AG’(Sb203, cubic, 25°C) = -632.08
40
[65] discussed above. The entropies were estimatedfrom values for arsenates and phosphates. Thus, there appear to be no reliable directlydetermined enthalpy of formation or entropy values for any of these solids, nor for the Sb(V)aqueous species.
[71] unavailable to thepresent reviewers. However, it is probable that this key value was also based on a (different)recalculation of Mixter’s measurements
kJmol-’ from a set of standard tables Afw(NasSb04) = -1485.3 [70]. The enthalpy values are relative to
41
were recently estimated by Kasenov et al.
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