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Conditional solubility versus pO2− of cerium(III) oxide in molten equimolar NaCl+KCl at 727°C
J. Electroanal. Chem., 91 (1978) 125--131 © Elsevier Sequoia S.A., Lausanne -- Printed in the Netherlands
125
CONDITIONAL SOLUBILITY VERSUS pO 2- OF CERIUM(III) OXIDE IN M O L T E N E Q U I M O L A R NaC1 + KC1 A T 7 2 7 ° C *
RICHARD COMBES **, MARIE-NOl~LLE LEVELUT and BERNARD TRl~MILLON Laboratoire d'Electrochimie Analytique et Appliqude (LA au CNRS no. 216), Ecole Nationale Supdrieure de Chimie de Paris, Universitd Pierre et Marie Curie, Paris (France)
(Received 17th October 1977; in revised form 29th November 1977)
ABSTRACT Reactions between cerium trichloride and oxide ions were studied in NaC1 + KC1 (1/1) at 1000°K, by potentiometry with a calcia-stabilized zirconia membrane electrode. Titration curves clearly demonstrated the existence of soluble cerium oxychloride (CeO +) and precipitated cerium oxide ( C e 2 0 3 ) , with respective dissociation constants 10 - 1 1 a n d 10 - 3 0 (molality scale). The corresponding conditional solubility diagram {log S (Ce III)= f(pO2--)} is presented and discussed.
(I) INTRODUCTION Like m o s t m e t a l ions, c e r i u m ( I I I ) ions in m o l t e n alkali chlorides f o r m stable c o m p o u n d s with t h e o x i d e ion 0 2 - . I n a d d i t i o n t o t h e v e r y scarcely soluble o x i d e C e 2 0 3 , t h e e x i s t e n c e o f soluble o x o c o m p l e x e s , such as the o x y c h l o r i d e CeOC1 ( x - l ) - (which will b e s y m b o l i z e d b y CeO + disregarding t h e s o l v e n t ions C 1 - c o m b i n e d in these species), or C e O ~ , m a y be c o n s i d e r e d if c h e c k e d b y an e x p e r i m e n t a l investigation. F o r m a t i o n or d i s s o c i a t i o n r e a c t i o n s o f these c o m p o u n d s b e l o n g t o t h e n o w familiar set o f o x o - a c i d o b a s i c s y s t e m s [ 1 - - 1 0 ] :
Ce3+ + 0 2 - # CeO+ 2 CeO+ + 0 2 - # Ce203 (s) ... If the reactant is an oxobase stronger than the corresponding cerium oxo com: plex, these equilibria will be displaced to the right side, as for example with OHor CO2-. On the contrary, a strong oxoacid, such as HCI (2 HCI + 0 2 - ~ H20 + 2 CI-), will displace equilibria to the left side. Consequently the solubility of cerium oxide in molten alkali chlorides (total concentration of dissolved cerium(III) species when saturation by metal oxide Ce203 is realized) will depend on the oxoacidity of the melt, i.e. will vary in * Research performed with the aid of the Centre National de la Recherche Scientifique, in the framework of the A.T.P. ~'Epargne d'Encrgie et Op/:rations Chimiques Industrielles". ** To whom correspondence should be addressed: Laboratoire d'Electrochimie Analytique et Appliqude. E.N.S.C.P., 11 rue Pierre et Marie Curie, 75231-Paris Cedex 05, France.
126 function of pO 2-. Here, the analytical signification has been given to pO2-: pO 2- = --log[O 2- ] where [ 0 2 - ] represents the concentration (in mol kg -1) of " f r e e " 0 2 - ions in the melt. In a previous study, Reinhard and Naumann [11] have performed some determinations of the solubility of Ce203 in NaC1 + KC1 (1/1) in the range 700-1000°C, using an isotopic dilution method with radioactive 141Ce. Corresponding values of the solubility product Ks = [Ce 3+] 2 [ 0 2 - ] 3 have been deduced (neglecting oxo complex species in the soluble part). In other respects, the same authors have determined the conditional solubility variations of Ce203 in molten NaC1 at 900°C, in function of the increasing concentration of added oxide ions, observing t h a t a decrease in solubility first occurred, followed by a slight increase attributed to the formation of C e 0 2 . The decreasing part was in good agreement with the hypothesis of the oxo complex CeO + formation. In our work, we attempted to check these results and to determine precisely the values of the equilibrium constants, using the potentiometric titration method by oxo-acidobasic reaction that we developed in several previous studies [ 12--15 ]. The solvent we considered was the NaC1--KC1 equimolar mixture; the temperature was fixed at 727°C (1000 K), for which we had a great number of experimental data in this melt. (II) EXPERIMENTS AND ANALYSIS OF THE RESULTS Titrations were performed by progressive additions of known amounts of sodium carbonate to a solution of CeC13 (~-0.3 mol kg -1) in the NaC1 + KC1 melt. The reactions occurring are assumed to be the following ones (symbolizing the dissolved cerium(III) chloride by Ce3+): Ce 3+ + CO 2--~ CeO + + CO2 (g) 2 CeO + + COa2- # Ce203 (s) + CO2 (g) Interpretation of measurements was made easier by working at an imposed value of carbon dioxide partial pressure (10 - 2 atm in Ar 0.99 atm), in order to fix the activity of this c o m p o u n d in equilibrium with the melt. Instead of Na2CO3, sodium hydroxide could have been used (by fixing the water vapor pressure), but the former was found much easier. Sodium oxide has been avoided because it is difficult to handle at a given stated purity and its harmful effect on the indicator electrode [13], being a much stronger oxide ion donor than the above mentioned oxobases. The pO 2 - indicator electrode is the one with calcia-stabilized zirconia membrane, which was previously described and tested by Combes et al. [12,13,16]. Moreover, it has been used for various determinations of equilibrium constants in molten alkaline and alkaline-earth chlorides by Combes et al. [13--19], and recently in molten alkali nitrates by Stern and Flinn [20]. With the aid of this electrode, coupled with a Ag/AgC1 reference, potentiometric titration curves, pO 2- = f (amount of added carbonate), were obtained. Standardization of the electrode (correspondence between measured potential and pO 2-) was realized
127 by means of the terminal part of the curves corresponding to the buffer system CO2-/CO2, whose equilibrium constant is known [17] and allows to write: pO 2 - = 4.8 + log Pc o: -- l o g [ C O l - ]
(1)
TECHNICAL PROCEDURE
(a) Measuring cell and electrodes The cell was composed of a pure alumina * crucible placed in a vacuum tight silica vessel water-cooled in the uppe r part. The reference electrode was constituted by a silver wire in equilibrium with a silver chloride solution (0.75 mol kg - 1 ) in NaC1 + KC1; the j unct i on with the melt was realized by means of a porous alumina m e m br a ne (tube containing the reference electrode). The pO 2 indicator electrode, which has previously been described [13], is made of a calcia-stabilized zirconia tube filled with a mixture of nickel oxide powders. The potential difference between these two electrodes was measured by means of a high impedance vol t m et er (Tacussel Aries 10.000). The carbon dioxide pressure was imposed by bubbling in the melt an Ar + CO2 mixture whose partial pressures were pr op ort i onal to their respective flows passing through a constant t e m p e r a t u r e device made by " O x h y d r i q u e Franqaise". The t e m p e r a t u r e of the melt was maintained constant within one degree at 727°C, by means of a furnace and a regulation device already described [13].
(b ) Materials Reagent grade sodium and potassium chlorides were used and handled as suggested elsewhere [ 21]. Cerium trichloride is hepta-hydrated at r o o m temperature. In order to avoid any trace of oxide or oxychloride, its d e h y d r a t i o n was c o n d u c t e d under an atmosphere o f h y d r o g e n chloride by a progressive heating inferred f r o m the thermogravimetric analysis available in the literature [ 22,23]. The mixture with NaCI + KC1 was heated under vacuum up to 400°C and melted under a h y d r o g e n chloride atmosphere, which was removed, just before the titration, by the CO2 + Ar mixture bubbling.
(c) Titration The p o t e n t i o m e t r i c curves were obtained by measuring the potential difference between the indicator and the reference electrode, after every addition of weighted pre-dried solid sodium carbonate to the melt containing 0.3 mol kg - 1 of Ce(III). The whole titration was p e r f o r m e d under a partial carbon dioxide pressure maintained at 10 - 2 atm. The resulting small change of Na/K ratio in the solvent composition, was assumed to be negligible. * As already mentioned in a previous paper [13] and as checked again in this work, pure alumina was not attacked, even at the highest available pO 2- values [ 16 ].
128 P° 2_ E/V
10i
~
w
-0.5 -1.0
I 0.5
I 1.0
[Na2CO3/~ECe~3
I 1.5
I 2.0
Fig. 1. P o t e n t i o m e t r i c t i t r a t i o n curve o f a 0.3 mol kg - 1 Ce 3+ s o l u t i o n in NaCl + KC1 ( 1 / 1 ) b y Na2CO3, u n d e r P c o 2 = 1 0 ~ 2 a t m . T = 727 ° C. ( , ) E x p e r i m e n t a l points. ( ) Calculated curve w i t h p K 1 = 10.5 a n d p K 2 = 7.8 ( m o l a l i t y scale).
RESULTS AND ANALYSIS
Three titration curves were performed and were similar to the one presented here. The corresponding potential measurements versus ~ (~ is the ratio of the added carbonate concentration to the original cerium(III) concentration) are represented in Fig. 1. The two expected equivalence points can be noticed; the first, for ~ = 1, corresponds to the formation of CeO ÷ and the second, for a = 1.5, is related to the formation of Ce2Oa from CeO ÷. The theoretical expression of the titration curve, based upon the assumed reactions, is correctly verified by the experimental points ranging respectively from A to B and C to D, as shown in Fig. 1. The difference between the experimental points and the theoretical curve (continuous curve calculated with the constant values given in Fig. 1), observed at the beginning of the titration, can be ascribed to the original presence of a low fraction of Ce(III) in the form of CeO ÷. The difference also observed just after the first equivalence point may be due either to an evolution of the Ce203 precipitate at the beginning of its formation, or to the intermediate formation of mixed condensed compounds, such as CexOx+lClx_2 (x > 2). TABLE 1 Dissociation c o n s t a n t s o f CeO + a n d C e 2 0 3 ( s ) in m o l t e n e q u i m o l a r NaC1 + KCI at 727°C ( m o l a l i t y scale)
pK 1 pK 2 pK s
Exp. 1
Exp. 2
Exp. 3
Mean value
10.5 7.8 29
11.5 7.7 31
11.0 8.2 30
11.0 (-+0.5) 7.9 (-+0.4) 30 (-+1)
129 Part EF was used to establish the correspondence between the potentiometric scale and the pO 2- values according to expression (1). In Table 1 are collected values of the constants pK1, pK2 and pKs, determined by analysis of the three experimental curves obtained. These constants correspond respectively to the equilibria: CeO +
# Ce 3+ + 0 2 -
K 1 = [Ce 3+] [ 0 2 - ][CeO+] -1
Ce203(s) # 2 CeO + + 0 2 -
K 2 = [CeO+]2[O 2-]
Ce203(s) ~ 2 Ce 3+ + 3 0 2 -
K s = [Ce3+]2[O2-]3 = K 2 K 2
(III) DISCUSSION AND CONCLUSION First, the present results can be compared either to those from other works or to thermochemical data. Fairly good agreement can be observed between our pKs values and that obtained by interpolation of Reinhard and Naumann's results [11] (pKs = 30.5 at 1000 K). The comparison of these two pK s values to that calculable from Glassner's [24] thermochemical data for pure substances (pKs = 51.9), leads us to ascribe an activity coefficient equal to 10 -3.7 for Ce 3+ ion -- taking into account the previously determined ones for oxide [ 13] and sodium [25] ions in this m e l t - close to the one available from Flengas and Ingraham's [26] works for Cr 3+ (i.e. 10 -3-5) and not too far from the ones determined potentiometrically by Mellors et al. in NaC1 and KC1 at 800°C [27] (10 -2.6 and 10 -~, respectively). In the case of the oxychloride formation constant, the thermodynamic value (pK~ - 19.6) calculated using the available literature data [28,29] for the reaction: CeC13(1) + Na20(1) # CeOCl(1) + 2 NaCl(1) is consistent with our experimental result taking into account the activity coefficients of the other substances as mentioned above. Unfortunately, no direct comparison with Reinhard and Naumann's work [11] is possible in this case. These authors obtained pK2 = 13.4 in NaC1 at 900°C, which would therefore correspond to a pK 1 value much lower than ours. Thus, it might appear that the oxychloride (CeO +) stability would decrease with the rise in temperature. Next, the discussion will deal with the conditional solubility variations of cerium(III) versus pO 2-. Considering t h a t cerium(III) is soluble only under the forms Ce 3÷ and CeO ÷ -- no appreciable dissolution of cerium(III) oxide (as CeO2) has been observed even at pO 2- ~ 1 -- and that these species are in equilibrium with Ce203 precipitated, one will obtain the following expression for the logarithm of the conditional solubility S (in the molality scale): S = (Kls/2/[02-]3/2)(1 + [02-]/K1)
Whence: log S = --15.0 + 1.5 pO 2- + log(1 + 1011102-]) The corresponding variation curve of log S versus pO 2- is represented in Fig. 2. It is limited to one line with a slope 0.5 if we suppose that solutions are limited to solute concentrations below 1 mol kg -1.
130
0~
5 II
~
:, D
i/
15
I
I
pO 2-
I
,:?:
/,:0 i
10
i ii
I? i
Q
i
IX1 LD /
/
--5
-,
o,,
"
~,
I" ~3 o _
i
6
l.b ]3-
,.
I
Fig. 2. Diagram of cerium(III) oxide conditional solubility versus pO 2at 727°C. Some practical pO 2- buffering mixtures are represented.
in NaC1 + KC1 (1/1)
With the aid of this diagram, the following predictions can be made. (a) Large amounts of cerium(III) oxide can be dissolved (at equilibrium) under the form Ce 3÷ (cerium trichloride) by action of gaseous hydrogen chloride (1 atm), even if it is very wet, because this mixture can raise the solution pO 2 up to 14 and above [13]. (b) Dissolved cerium(III) will essentially be in the form of CeO ÷ (cerium oxychloride) in presence of water vapor, even when slightly charged with HC1 (~ 1%). This explains the hydrolysis of cerium(III) species when dissolved in molten chlorides at high temperature. (c) For practically complete precipitation of Ce203 from molten chlorides, sodium hydroxide or carbonate are the most convenient reactants. But, in the case of carbonate, a high pressure of CO 2 will prevent the precipitation, as can be calculated from the corresponding constants [17]. This set of results is in good agreement with literature data and gives some quantitative explanations of the behaviour of cerium in molten chlorides, pointing out the important part played by the oxychloride species. REFERENCES
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
H. L u x , Z. E l e k t r o c h e m . , 4 5 ( 1 9 3 9 ) 3 0 3 ; N a t u r w i s s . , 2 6 ( 1 9 4 0 ) 9 2 . H. F l o o d a n d T. F o r l a n d , A c t a C h e m . S c a n d . , 1 ( 1 9 4 7 ) 5 9 2 . R. L i t t l e w o o d , J. E l e c t r o c h e m . S o c . , 1 0 9 ( 1 9 6 2 ) 5 2 5 . M. I n g r a i n a n d G. J a n z , E l e c t r o c h i m . A c t a , 1 0 ( 1 9 6 5 ) 7 8 3 . J. G o r e t a n d B. T r e m i l l o n , Bull. S o c . C h i m . F t . , ( 1 9 6 6 ) 6 7 , 2 8 7 2 . N. W r e n c h a n d D. I n m a n , Brit. C o r r o s . J., 1 ( 1 9 6 6 ) 2 4 6 . P. Z a m b o n i n a n d J. J o r d a n , J. A m e r . C h e m . S o c . , 89 ( 1 9 6 7 ) 6 3 6 5 . A. R a h m e l , E l e c t r o c h i m . A c t a , 1 3 ( 1 9 6 8 ) 4 9 5 . B. T r e m i l l o n , P u r e A p p l . C h e m . , 2 5 ( 1 9 7 1 ) 3 9 5 . K. F u n g a n d G. M a m a n t o v in A d v a n c e s in M o l t e n Salts C h e m i s t r y , Vol. 2, P l e n u m Press, N e w Y o r k , 1973. G. R e i n h a r d a n d D. N a u m a n n , Z. P h y s . C h e m . , 2 3 8 ( 1 9 6 8 ) 3 0 9 . R . Combes~ J. V e d e l a n d B. T r e m i l l o n , A n a l . L e t t . , 3 ( 1 9 7 0 ) 5 2 3 . R. C o m b e s , J. V e d e l a n d B. T r e m i l l o n , E l e c t r o c h i m . A c t a , 2 0 ( 1 9 7 5 ) 1 9 1 . F. de A n d r a d e , R. C o m b e s a n d B. T r e m i l l o n , J. E l e c t r o a n a l . C h e m . , 83 ( 1 9 7 7 ) 2 9 7 . R, C o m b e s , F. de A n d r a d e a n d L. C a r v a l h o , C . R . A c a d . Sci. (Paris), 2 6 5 ( 1 9 7 7 ) 1 3 7 . R. C o m b e s , Thesis, Paris, 1 9 7 3 .
131 17 18 19 20 21 22 23 24 25 26 27 28 29
R. C o m b e s , R . F e y s a n d B. T r e m i l l o n , J. E l e c t r o a n a l . C h e m . , 8 3 ( 1 9 7 7 ) 3 8 3 . R. C o m b e s a n d R. L y s y , J. E l e c t r o a n a l . C h e m . , 8 3 ( 1 9 7 7 ) 2 8 7 . R. C o m b e s , R. F e y s a n d B. T r e m i l l o n , t o b e p u b l i s h e d . K. S t e r n a n d D. F l i n n , J. E l e c t r o a n a l . C h e m . , 6 3 ( 1 9 7 5 ) 39. G. J a n z , M o l t e n Salts H a n d b o o k , A c a d e m i c Press, N e w Y o r k , 1 9 6 7 . O. H S n i g s c h m i d a n d H. H o l z , A n o r g . C h e m . , 1 7 7 ( 1 9 2 8 ) 91. T. K o j i m a , J. E l e c t r o c h e m . Soc. J a p . , 1 1 9 ( 1 9 5 3 ) 23. A. G l a s s n e r , T h e T h e r m o c h e m i c a l P r o p e r t i e s o f t h e O x i d e s , F l u o r i d e s a n d C h l o r i d e s t o 2 5 0 0 K , U.S. A t o m i c E n e r g y C o m m i s s i o n R e p t . , A N L 5 1 0 7 . D. T o p o r a n d L. T o p o r , J. C h i m . P h y s . , 6 8 ( 1 9 7 1 ) 1 4 6 4 . S. F l e n g a s a n d T. I n g r a h a m , J. E l e c t r o c h e m . S o c . , 1 0 6 ( 1 9 5 9 ) 7 1 4 . S. S e n d e r o f f , G. Mellors a n d R. Bretz, A n n . N . Y . A c a d . Sci., 7 9 ( 1 9 6 0 ) 8 7 8 ; G. MeUors, P r o c . 1st Austr. Conf. Eleetrochem., Sydney, 1963. C. K o c h , A. BroYdo a n d B. C u n n i n g h a m , J. A m e r . C h e m . S o c . , 7 4 ( 1 9 5 2 ) 2 3 4 9 ; 7 5 ( 1 9 5 3 ) 7 9 6 ; 76 (1954) 1471. H. Baev a n d G. N o v i k o v , R u s s . J. I n o r g . C h e m . , 1 0 ( 1 9 6 5 ) 1 3 3 7 .
125
CONDITIONAL SOLUBILITY VERSUS pO 2- OF CERIUM(III) OXIDE IN M O L T E N E Q U I M O L A R NaC1 + KC1 A T 7 2 7 ° C *
RICHARD COMBES **, MARIE-NOl~LLE LEVELUT and BERNARD TRl~MILLON Laboratoire d'Electrochimie Analytique et Appliqude (LA au CNRS no. 216), Ecole Nationale Supdrieure de Chimie de Paris, Universitd Pierre et Marie Curie, Paris (France)
(Received 17th October 1977; in revised form 29th November 1977)
ABSTRACT Reactions between cerium trichloride and oxide ions were studied in NaC1 + KC1 (1/1) at 1000°K, by potentiometry with a calcia-stabilized zirconia membrane electrode. Titration curves clearly demonstrated the existence of soluble cerium oxychloride (CeO +) and precipitated cerium oxide ( C e 2 0 3 ) , with respective dissociation constants 10 - 1 1 a n d 10 - 3 0 (molality scale). The corresponding conditional solubility diagram {log S (Ce III)= f(pO2--)} is presented and discussed.
(I) INTRODUCTION Like m o s t m e t a l ions, c e r i u m ( I I I ) ions in m o l t e n alkali chlorides f o r m stable c o m p o u n d s with t h e o x i d e ion 0 2 - . I n a d d i t i o n t o t h e v e r y scarcely soluble o x i d e C e 2 0 3 , t h e e x i s t e n c e o f soluble o x o c o m p l e x e s , such as the o x y c h l o r i d e CeOC1 ( x - l ) - (which will b e s y m b o l i z e d b y CeO + disregarding t h e s o l v e n t ions C 1 - c o m b i n e d in these species), or C e O ~ , m a y be c o n s i d e r e d if c h e c k e d b y an e x p e r i m e n t a l investigation. F o r m a t i o n or d i s s o c i a t i o n r e a c t i o n s o f these c o m p o u n d s b e l o n g t o t h e n o w familiar set o f o x o - a c i d o b a s i c s y s t e m s [ 1 - - 1 0 ] :
Ce3+ + 0 2 - # CeO+ 2 CeO+ + 0 2 - # Ce203 (s) ... If the reactant is an oxobase stronger than the corresponding cerium oxo com: plex, these equilibria will be displaced to the right side, as for example with OHor CO2-. On the contrary, a strong oxoacid, such as HCI (2 HCI + 0 2 - ~ H20 + 2 CI-), will displace equilibria to the left side. Consequently the solubility of cerium oxide in molten alkali chlorides (total concentration of dissolved cerium(III) species when saturation by metal oxide Ce203 is realized) will depend on the oxoacidity of the melt, i.e. will vary in * Research performed with the aid of the Centre National de la Recherche Scientifique, in the framework of the A.T.P. ~'Epargne d'Encrgie et Op/:rations Chimiques Industrielles". ** To whom correspondence should be addressed: Laboratoire d'Electrochimie Analytique et Appliqude. E.N.S.C.P., 11 rue Pierre et Marie Curie, 75231-Paris Cedex 05, France.
126 function of pO 2-. Here, the analytical signification has been given to pO2-: pO 2- = --log[O 2- ] where [ 0 2 - ] represents the concentration (in mol kg -1) of " f r e e " 0 2 - ions in the melt. In a previous study, Reinhard and Naumann [11] have performed some determinations of the solubility of Ce203 in NaC1 + KC1 (1/1) in the range 700-1000°C, using an isotopic dilution method with radioactive 141Ce. Corresponding values of the solubility product Ks = [Ce 3+] 2 [ 0 2 - ] 3 have been deduced (neglecting oxo complex species in the soluble part). In other respects, the same authors have determined the conditional solubility variations of Ce203 in molten NaC1 at 900°C, in function of the increasing concentration of added oxide ions, observing t h a t a decrease in solubility first occurred, followed by a slight increase attributed to the formation of C e 0 2 . The decreasing part was in good agreement with the hypothesis of the oxo complex CeO + formation. In our work, we attempted to check these results and to determine precisely the values of the equilibrium constants, using the potentiometric titration method by oxo-acidobasic reaction that we developed in several previous studies [ 12--15 ]. The solvent we considered was the NaC1--KC1 equimolar mixture; the temperature was fixed at 727°C (1000 K), for which we had a great number of experimental data in this melt. (II) EXPERIMENTS AND ANALYSIS OF THE RESULTS Titrations were performed by progressive additions of known amounts of sodium carbonate to a solution of CeC13 (~-0.3 mol kg -1) in the NaC1 + KC1 melt. The reactions occurring are assumed to be the following ones (symbolizing the dissolved cerium(III) chloride by Ce3+): Ce 3+ + CO 2--~ CeO + + CO2 (g) 2 CeO + + COa2- # Ce203 (s) + CO2 (g) Interpretation of measurements was made easier by working at an imposed value of carbon dioxide partial pressure (10 - 2 atm in Ar 0.99 atm), in order to fix the activity of this c o m p o u n d in equilibrium with the melt. Instead of Na2CO3, sodium hydroxide could have been used (by fixing the water vapor pressure), but the former was found much easier. Sodium oxide has been avoided because it is difficult to handle at a given stated purity and its harmful effect on the indicator electrode [13], being a much stronger oxide ion donor than the above mentioned oxobases. The pO 2 - indicator electrode is the one with calcia-stabilized zirconia membrane, which was previously described and tested by Combes et al. [12,13,16]. Moreover, it has been used for various determinations of equilibrium constants in molten alkaline and alkaline-earth chlorides by Combes et al. [13--19], and recently in molten alkali nitrates by Stern and Flinn [20]. With the aid of this electrode, coupled with a Ag/AgC1 reference, potentiometric titration curves, pO 2- = f (amount of added carbonate), were obtained. Standardization of the electrode (correspondence between measured potential and pO 2-) was realized
127 by means of the terminal part of the curves corresponding to the buffer system CO2-/CO2, whose equilibrium constant is known [17] and allows to write: pO 2 - = 4.8 + log Pc o: -- l o g [ C O l - ]
(1)
TECHNICAL PROCEDURE
(a) Measuring cell and electrodes The cell was composed of a pure alumina * crucible placed in a vacuum tight silica vessel water-cooled in the uppe r part. The reference electrode was constituted by a silver wire in equilibrium with a silver chloride solution (0.75 mol kg - 1 ) in NaC1 + KC1; the j unct i on with the melt was realized by means of a porous alumina m e m br a ne (tube containing the reference electrode). The pO 2 indicator electrode, which has previously been described [13], is made of a calcia-stabilized zirconia tube filled with a mixture of nickel oxide powders. The potential difference between these two electrodes was measured by means of a high impedance vol t m et er (Tacussel Aries 10.000). The carbon dioxide pressure was imposed by bubbling in the melt an Ar + CO2 mixture whose partial pressures were pr op ort i onal to their respective flows passing through a constant t e m p e r a t u r e device made by " O x h y d r i q u e Franqaise". The t e m p e r a t u r e of the melt was maintained constant within one degree at 727°C, by means of a furnace and a regulation device already described [13].
(b ) Materials Reagent grade sodium and potassium chlorides were used and handled as suggested elsewhere [ 21]. Cerium trichloride is hepta-hydrated at r o o m temperature. In order to avoid any trace of oxide or oxychloride, its d e h y d r a t i o n was c o n d u c t e d under an atmosphere o f h y d r o g e n chloride by a progressive heating inferred f r o m the thermogravimetric analysis available in the literature [ 22,23]. The mixture with NaCI + KC1 was heated under vacuum up to 400°C and melted under a h y d r o g e n chloride atmosphere, which was removed, just before the titration, by the CO2 + Ar mixture bubbling.
(c) Titration The p o t e n t i o m e t r i c curves were obtained by measuring the potential difference between the indicator and the reference electrode, after every addition of weighted pre-dried solid sodium carbonate to the melt containing 0.3 mol kg - 1 of Ce(III). The whole titration was p e r f o r m e d under a partial carbon dioxide pressure maintained at 10 - 2 atm. The resulting small change of Na/K ratio in the solvent composition, was assumed to be negligible. * As already mentioned in a previous paper [13] and as checked again in this work, pure alumina was not attacked, even at the highest available pO 2- values [ 16 ].
128 P° 2_ E/V
10i
~
w
-0.5 -1.0
I 0.5
I 1.0
[Na2CO3/~ECe~3
I 1.5
I 2.0
Fig. 1. P o t e n t i o m e t r i c t i t r a t i o n curve o f a 0.3 mol kg - 1 Ce 3+ s o l u t i o n in NaCl + KC1 ( 1 / 1 ) b y Na2CO3, u n d e r P c o 2 = 1 0 ~ 2 a t m . T = 727 ° C. ( , ) E x p e r i m e n t a l points. ( ) Calculated curve w i t h p K 1 = 10.5 a n d p K 2 = 7.8 ( m o l a l i t y scale).
RESULTS AND ANALYSIS
Three titration curves were performed and were similar to the one presented here. The corresponding potential measurements versus ~ (~ is the ratio of the added carbonate concentration to the original cerium(III) concentration) are represented in Fig. 1. The two expected equivalence points can be noticed; the first, for ~ = 1, corresponds to the formation of CeO ÷ and the second, for a = 1.5, is related to the formation of Ce2Oa from CeO ÷. The theoretical expression of the titration curve, based upon the assumed reactions, is correctly verified by the experimental points ranging respectively from A to B and C to D, as shown in Fig. 1. The difference between the experimental points and the theoretical curve (continuous curve calculated with the constant values given in Fig. 1), observed at the beginning of the titration, can be ascribed to the original presence of a low fraction of Ce(III) in the form of CeO ÷. The difference also observed just after the first equivalence point may be due either to an evolution of the Ce203 precipitate at the beginning of its formation, or to the intermediate formation of mixed condensed compounds, such as CexOx+lClx_2 (x > 2). TABLE 1 Dissociation c o n s t a n t s o f CeO + a n d C e 2 0 3 ( s ) in m o l t e n e q u i m o l a r NaC1 + KCI at 727°C ( m o l a l i t y scale)
pK 1 pK 2 pK s
Exp. 1
Exp. 2
Exp. 3
Mean value
10.5 7.8 29
11.5 7.7 31
11.0 8.2 30
11.0 (-+0.5) 7.9 (-+0.4) 30 (-+1)
129 Part EF was used to establish the correspondence between the potentiometric scale and the pO 2- values according to expression (1). In Table 1 are collected values of the constants pK1, pK2 and pKs, determined by analysis of the three experimental curves obtained. These constants correspond respectively to the equilibria: CeO +
# Ce 3+ + 0 2 -
K 1 = [Ce 3+] [ 0 2 - ][CeO+] -1
Ce203(s) # 2 CeO + + 0 2 -
K 2 = [CeO+]2[O 2-]
Ce203(s) ~ 2 Ce 3+ + 3 0 2 -
K s = [Ce3+]2[O2-]3 = K 2 K 2
(III) DISCUSSION AND CONCLUSION First, the present results can be compared either to those from other works or to thermochemical data. Fairly good agreement can be observed between our pKs values and that obtained by interpolation of Reinhard and Naumann's results [11] (pKs = 30.5 at 1000 K). The comparison of these two pK s values to that calculable from Glassner's [24] thermochemical data for pure substances (pKs = 51.9), leads us to ascribe an activity coefficient equal to 10 -3.7 for Ce 3+ ion -- taking into account the previously determined ones for oxide [ 13] and sodium [25] ions in this m e l t - close to the one available from Flengas and Ingraham's [26] works for Cr 3+ (i.e. 10 -3-5) and not too far from the ones determined potentiometrically by Mellors et al. in NaC1 and KC1 at 800°C [27] (10 -2.6 and 10 -~, respectively). In the case of the oxychloride formation constant, the thermodynamic value (pK~ - 19.6) calculated using the available literature data [28,29] for the reaction: CeC13(1) + Na20(1) # CeOCl(1) + 2 NaCl(1) is consistent with our experimental result taking into account the activity coefficients of the other substances as mentioned above. Unfortunately, no direct comparison with Reinhard and Naumann's work [11] is possible in this case. These authors obtained pK2 = 13.4 in NaC1 at 900°C, which would therefore correspond to a pK 1 value much lower than ours. Thus, it might appear that the oxychloride (CeO +) stability would decrease with the rise in temperature. Next, the discussion will deal with the conditional solubility variations of cerium(III) versus pO 2-. Considering t h a t cerium(III) is soluble only under the forms Ce 3÷ and CeO ÷ -- no appreciable dissolution of cerium(III) oxide (as CeO2) has been observed even at pO 2- ~ 1 -- and that these species are in equilibrium with Ce203 precipitated, one will obtain the following expression for the logarithm of the conditional solubility S (in the molality scale): S = (Kls/2/[02-]3/2)(1 + [02-]/K1)
Whence: log S = --15.0 + 1.5 pO 2- + log(1 + 1011102-]) The corresponding variation curve of log S versus pO 2- is represented in Fig. 2. It is limited to one line with a slope 0.5 if we suppose that solutions are limited to solute concentrations below 1 mol kg -1.
130
0~
5 II
~
:, D
i/
15
I
I
pO 2-
I
,:?:
/,:0 i
10
i ii
I? i
Q
i
IX1 LD /
/
--5
-,
o,,
"
~,
I" ~3 o _
i
6
l.b ]3-
,.
I
Fig. 2. Diagram of cerium(III) oxide conditional solubility versus pO 2at 727°C. Some practical pO 2- buffering mixtures are represented.
in NaC1 + KC1 (1/1)
With the aid of this diagram, the following predictions can be made. (a) Large amounts of cerium(III) oxide can be dissolved (at equilibrium) under the form Ce 3÷ (cerium trichloride) by action of gaseous hydrogen chloride (1 atm), even if it is very wet, because this mixture can raise the solution pO 2 up to 14 and above [13]. (b) Dissolved cerium(III) will essentially be in the form of CeO ÷ (cerium oxychloride) in presence of water vapor, even when slightly charged with HC1 (~ 1%). This explains the hydrolysis of cerium(III) species when dissolved in molten chlorides at high temperature. (c) For practically complete precipitation of Ce203 from molten chlorides, sodium hydroxide or carbonate are the most convenient reactants. But, in the case of carbonate, a high pressure of CO 2 will prevent the precipitation, as can be calculated from the corresponding constants [17]. This set of results is in good agreement with literature data and gives some quantitative explanations of the behaviour of cerium in molten chlorides, pointing out the important part played by the oxychloride species. REFERENCES
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
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