Solubility product of silver oxide in molten alkali nitrates

ELECTROANALYTICALCHEMISTRYAND INTERFACIALELECTROCHEMISTRY Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands

137

S O L U B I L I T Y PRODUCT OF S I L V E R O X I D E IN MOLTEN A L K A L I NITRATES

A. M. SHAMS EL DIN, T. GOUDA AND A. A. EL HOSARY Laboratory of Electrochemistry and Corrosion, National Research Centre, Dokki, Cairo (U.A.R.)

(Received September 8th, 1967)

Precipitation reactions ill molten salts have not, so far, beenwidelyinvestigated and attention has been mainly focussed on silver halides 1-1~, cyanide 1,4,5,8,~2 and chromate 1. LYALIKOV et al.14 titrated Pb 2+ and Cd 2+ in fused KN03 by adding weighed quantities of KOH to the melts, following the reaction potentiometrically with a glass electrode. BOMBI AND FIORAN115 reported briefly on the titration of fourteen cations (not including Ag +) in fused NaNO3-KN03. The titrant, 0 2-, was electrogenerated in situ by reducing the nitrate melt16, and the neutralization process was followed with an oxygen indicator electrode. In the work of LYALIKOV14 and BOMBI~5 the general precipitation reaction: M~+n/2 0 2- -+ MO~/2

(I)

was regarded as satisfying L u x ' s definition of aciditylL viz., Acid + 0 3- =Base, and the acid character of the different cations were assessed. The applicability of a number of M/MO electrodes for 0 ~- determination in LiC1-KC1 eutectic was established by LAITIICEN AND BHATIA 18, while DELARUE studied, qualitatively, the precipitation of some metal oxides and sulphides in the same melt ~9. As part of an investigation on the electrodic behaviour of silver in fused salts, we studied the reaction: 2 Ag+ + 0 ~ - -+ Ag20

(2)

in KN03 in the temperature range 350-425 °, and in equimolar NaNO~-KNO3 in the range 250-35 o°. Known amounts of Ag+ in these melts were titrated by adding weighed quantities of Na202 as oxide-ion donor, and the course of precipitation was followed potentiometrically with a silver electrode. In some experiments, the potential of an oxygen electrode was measured simultaneously. EXPERIMENTAL The nitrate supporting electrolytes were prepared and dried as already described 2°. AnalaR AgN03 was used without further purification. Na202 was used as oxide-ion donor. A recent study 21 showed it to possess the same basicity and to react in the same manner as NaOH or electrolytically-generated 0 2-. Also it has the advantage over the other two reagents of not introducing water or nitrite into the melt. Interference from the possible formation of silver peroxide, Ag202, was unlikely as this compound is known to decompose thermally at temperatures much lower than those of the present investigation 2~. The titration experiments were carried out by adding small weighed amounts of j . Electroanal. Chem., 17 (1968) 137-143

138

A. M. SHAMS EL DIN, T. GOUDA, A. A. EL HOSARY

Na~02 to 5o.ooo-g nitrate melts containing varying quantities of AgN03. The potential of a silver wire electrode (99.9 + %) was measured relative to a Ag/Ag(I), melt/glass reference electrode 28. The concentration of Ag(I) in the reference electrode was o.12o mole kg -1. In experiments with KNO~ the titration was followed also with an oxygen (Pt) electrode relative to the same reference. Before the equivalence points, steadystate potentials with both the silver and oxygen electrodes established rather quickly, 2-4 min after each Na202 addition, and were highly reproducible. For reasons to be mentioned later, the potentials of the indicator electrode(s) after the equivalence point were measured 3 rain following each Na202 addition. These potentials were satisfactorily reproducible. The titration curves obtained with the oxygen electrode had practically the same features as those of the silver electrode, but were slightly displaced towards more negative potentials. The temperature of the melt was controlled to + 2 °. At each temperature, a series of experiments was conducted with at least four different concentrations of AgN0~. The Jnflexion points of the titration curves corresponded to within _ I % to those calculated theoretically for reaction (2). RESULTS AND DISCUSSION

The validity of the Nernst equation for the Ag/Ag(I) couple in molten nitrates has been repeatedly established 1,~,6,9,10. Owing, however, to differences in the working temperatures and/or reference electrodes, the results of previous investigations cannot be directly applied to our case. Therefore, we carried out similar measurements in KNO3 and in NaNO3-KN03 melts at the temperatures of the titrations, and with the same reference electrode; the steady-state potential of the silver electrode was measured following the addition of successive amounts of AgNOa to the melts. The results of these measurements are shown graphically in Fig. I, A for KNOa and in Fig. I,B for NaNOa-KN03. The points of these curves were experimentally determined, the lines were drawn with a theoretical slope of 2.303 R T / F at the corresponding temper425 * 0

v

-40

..j

325"

.400 ° ,375"

• 425 //400

350" o

/

//375

////,o

Oi '300 ° 0.2?5 °

-46

60 o

~. 12o

I2C

t60

16(

200

200

g, 24o

325 300 275

..

o///i.o

/

2401

-3

i

t

-2

-I

-2

-I

0

Fig. I. N e r n s t relation for t h e Ag/Ag+ electrode in (A), KNO3; (B), e q u i m o l a r N a N O 3 - K N O 3 melts at various temps. J . ElectroanM. C h e m . , 17 (1968) 137-143

Ag20 IN MOLTEN NITRATES

SOLUBILITY OF

139

ature. Agreement between the two sets of figures is satisfactory. In KNOa melts, o.12o molal with respect to AgNO~ (equal to that in the reference half-cell), the potential of the silver electrode is - 2 o , - 1 6 , - 1 4 and - 6 mV at 350, 375, 4 °0 and 425 °, respectively. These potentials represent the junction potential, E j, introduced through the glass membrane of the reference electrode. In NaNOa-KNOa melts, however, Ej was practically independent of temperature in the range 25o-35 °0 ( - 2 2 mV). Since the same type of glass was used in both melts, it appears that Ej depends on the melt used, and is not a simple function of temperature. It should, however, be corrected for when calculating the concentration of Ag + in the melt from the Nernst equation. Curves representing the titration of various quantities of AgN03 in fused IKNO3 at 35 o° are given in Fig. 2,A. The curves of Fig. 2,B are for the same titration in NaNOa-KN03 at 25 o°. These curves, typical for some 8o experiments carried out in the two melts in the temperature range 25o-425 °, show distinct potential drops of the silver indicator electrode at the points of equivalence. In KN03, the potential inflexions were about 5oo mV, whereas in NaNO3-KNOa t h e y ranged between 3oo and 350 mV. The difference between the two values is apparently related to the acid character of the two solvents, the sodium melt being more acid.

CB)

g -20~.

400

600 -4 tU

600 i

i

o.I

i

i

J

02

i

i

Na2

02

i

o.1

03

i

i

02

(g)

Fig. 2. P o t e n t i o m e t r i c titration of AgNOs with Na202 in: (A), KNO3 at 35o°; (B), equimolar NaNO3-KNOa at 25 °0 . ( A - - I ) , 1 . 1 3 2 ; ( 2 ) , 2 . 1 o 6 ; ( 3 ) , 3 . 4 2 5 ; ( 4 ) , 4 . 8 o 2 ; ( 5 ) , 7 .262 × I o - z m o l e k g 1 AgNOa. (t3--1), 2.350; (2), 3.531; (3), 4.582; (4), 5 .88° × IO-2 mole kg-1 AgNO3.

The solubility product, Ks, of Ag20 was calculated from the titration curves assuming the applicability of the Nernst relation at high Ag+-dilution. The molal concentration of Ag +, C2, was calculated from the relation: E =2.303

(RT/F)log (C1/C2)+Ej

(3)

for three successive potentials directly following the equivalence points, and displaced 25 mV from one another. CI in eqn. (3) is the concentration of silver in the reference half-cell (o.12o mole kg-1) and Ej is the junction potential at the particular temperaj . Electroanal. Chem., 17 (1968) 137-143

14o

A . M . SHAMS EL DIN, T. GOUDA, A. A. EL I-IOSARY

ture. The corresponding molal concentration of 02--ion was taken as: Co 2- = f ( w - w e )

(4)

where w and we are, respectively, the weight of Na202 added to establish a particular potential, and that consumed in the complete precipitation of Ag20; f is a proportionality factor for converting weights into molalities. As the concentration of free Ag +ion after precipitation was too low compared to that of the free oxide-ion, it was not included in eqn. (4). D a t a for the calculation of the molal solubility product of Ag20 in molten TABLE 1 C A L C U L A T I O N OF T H E M O L A L S O L U B I L I T Y

Temp. (°C)

Initial AgNOa (mole kg-l,zo ~)

PRODUCT OF

Ag20 IN

MOLTEN ALKALI NITRATES

E -- Ej

CA9+

Co 2-

K s

(mV)

(mole kg-l.zo 7)

(mole kg-l.±o a)

(mole a kg-S.3:o 16)

(A) MoltenKNOs 35o

4.8o2

--655 --68o --705

5.797 3.624 2.271

1.282 1.923 3.845

375

3.572

--659 --684 --7o9

8.691 5.537 3.542

1.282 2-564 5.64 °

m e a n --

4.31 2.53 1.98 2.70

mean =

9.69 7 .86 7.o7 8.20

400

3.507

--661 --686 --711

12.97 ° 8.420 5.451

1.282 2.564 3.84 °

21.6o 18.17 11.4o m e a n ~ 16.1o

425

2.o13

--669 --694 --719

17.15o 11.3oo 7.677

1.282 2.564 5.270

37.70 32.76 31.o6 mean=33.6o

(B) MoltenKNOa-NaNOs

K,.~o 14

250

5 .880

--453 --478 --503

4.991 2.860 1.634

o.641 1.923 4.486

1.6o 1.57 1.2o m e a n = 1.5o

275

8.630

--453 --478 --5o3

7.948 4.670 2.751

1.282 3.205 5.77 °

8.1o 6.98 4.36 5.55

300

2.51o

--453 --478 --503

12.14o 7.315 4.397

0.769 1.923 4.842

mean=

11.33 lO.3O 9.36 mean

325

3.7 °0

--453 --478 --503

17.75o lO.92o 6.718

=

0.769 1.282 3.846 mean=

35 °

6-344

--453 --478 --503

25.240 15.82o 9.913

j . Electroanal. Chem., 17 (1968) 137-143

o.641 1.282 1.923

lO.8o

24.23 15.36 17.36 18.9o

40.86 32.09 18.9o m e a n = 27.2o

SOLUBILITY OF _A_g20 IN MOLTEN NITRATES

141

KNO~ at various temperatures are given in Table I,A. The average Ks-values listed in the table represent the mean of at least 12 calculations on four solutions of different initial AgNO3 concentration. Table I,B gives the same data for molten NaNO3KNOB in the temperature range 250-35 o°. The values of Ks for both melts, computed from potentials not far removed from the equivalence points, are--within a factor of 2--reasonably consistent to justify a consideration of their mean values. At comparable temperatures (35o°), Ks in NaNOs-KNO3 is ca. lO s times as large as in KNOs. This increased solubility, in the former m e l t is d u e - - a s is also the smaller potential drop at the inflexion points of the titration c u r v e s - - t o the fact that the sodium melt is more acid than the corresponding potassium melt 24. Calculation of Ks from potentials displaced more than 50 mV from the equivalence points gave smaller values than those listed in Table I, and the decrease was larger, the larger the quantity of the free oxide-ion donor in the melt. This behaviour indicates the tendency of Ag20 to redissolve, probably as AGO-. The concentration of the free silverion, CAg+, to which the silver indicator electrode would respond, would be given as CA~+=V(KsKc/C2Ago-) rather than V(Ks/Co2-); where Kc is the formation constant of the argentate ion. An estimation of the magnitude of Kc from the variation of the potential of the silver electrode with the amount of Na20~ added after complete precipitation, is conceivable but impossible in these relatively basic melts as no steady-state potentials can be measured. Thus, whereas before the equivalence points both the silver and the oxygen electrodes acquired steady-state potentials that were practically time-independent, the potentials measured after the inflexions changed with time in a particular manner. Each new addition of Na202 caused the potential to drop instantaneously to negative values, but after a time these drifted once again towards positive values. The rate of potential change depended on the ambient temperature and on the quantity of Na~O2 added to the melt. In KN03, an end potential varying between - 5 5 0 and - 6 0 0 mV was usually attained. This behaviour is due to the reaction of Na~02 with the nitrate base electrolyte to yield pyronitrate, N20 74- 25. The rate of this reaction was found to be independent of the presence of other material in the melt, e.g., Ag~O, CrO42- or PO4 ~- 25. In the construction of the lower parts of the titration curves, therefore, the potentials of the indicator electrodes were always measured 3 min after each Na202-addition, during which time pyronitrate-formation was considered not to interfere seriously with the Ag/Ag20 couple. These potentials were fairly reproducible. The density of fused KN03 over a wide range of temperature was given b y LORENZ, FREI AND JABS as26: d = 2 . o 4 4 - o . o o o 6 t g cm-3

(5)

where t is the temperature ill degrees centigrade. This allows the calculation of the

molar solubility product of Ag~O in this melt. Since molar and molal concentrations are related as: m o l a r i t y = m o l a l i t y x d , the average molar Ks of Ag20 in molten KNOs is 1.67, 4.93, 9.45 and 19.28 x lO -15 mole31-3 at 350, 375,400 and 425 °, successively. Another possibility for determining the solubility product of Ag20 without the complications discussed above is to apply an oxide-ion indicator electrode for the evaluation of the concentration of the free oxide-ion, Co ~-, along the upper parts of j . Electroanal. Chem., 17 (1968) I37-I43

142

A.M. SHAMSEL DIN, T. GOUDA,A. A. EL ttOSARY

the neutralization curves before the points of inflexion. For this purpose we investigated the suitability of the oxygen (Pt) electrode for following the precipitation of Ag20 in molten KN03. Excellent titration curves were always obtained with sharp potential inflexions (ca. 500 mV) which coincided with those measured with the silver indicator electrode. For these curves Co 2- was determined from the potential of the oxygen electrode at 25, 50 and 75% neutralization. Use was made of the E°-value of the system O2/O ~- theoretically calculated for this melt at 35 o°, viz., - 7 9 6 . 6 mV with respect to the same reference 27. The corresponding values of CAg+ were taken as the initial (molar) concentrations minus those precipitated. From these considerations, Ks was calculated to be 6.76, 9.1o, 33 and 161 x lO -15 moleal-3 at 350, 375, 4 °0 and 425 °, respectively. Although 2-8 times larger than those calculated from the silver electrode, these values are of the same order of magnitude, and support the previous conclusion that the oxygen electrode behaves reversibly in well-buffered acid melts 2s (see, however, refs. 27-29 for the behaviour in unbuffered melts). The difference between the two sets of figures is most probably the result of the simplification introduced ill the calculation of Co 2-, viz., the application of E ° of the system 02/0 2- at 35 °0 also for other temperatures, without consideration of any possible temperature coefficient. -/2

-13

o

~o -7,;

-I,5

i

,,

1.4

1.5

1.6

t

i

L7

l.a

I

1.9

i

2.0

1 / T X 103

Fig. 3. Log K, vs. I/T plots for Ag~O. (i), molal solubility product in KNOa; (2) molar solubility product in KNO~; (3), molal solubility product in equimolar NaNO~-KNO~. In Fig. 3 the values of Ks are plotted as a function of the reciprocal of the absolute temperature. For KNOa, two lines are given, representing molal and molar values, both calculated from the potential of the silver electrode. From the slopes of the lines of Fig. 3 the heats of dissolution, AH, of Ag20 were calculated to be 28.5 ° kcal mole -1 in KNOa and 14.82 kcal mole -1 in NaNOa-KNOa. The difference in AH in the two solvents reflects a corresponding difference in AG (pKs) as well as AS of the dissolution reaction. At 35 o°, AS is --25.26 cal deg -~ mole-1 in KNO3, and -33.51 cal deg -1 mole -~ in NaN0a-KN0a. SUMMARY The applicability of the Nernst equation for the Ag/Ag(I) couple in fused KNOa j . Electroanal. Chem., I7 (I968) I37-143

SOLUBILITY OF

Ag20 IN

MOLTEN NITRATES

143

in the temperature range 350-425 ° and in fused N a N 0 a - K N 0 a in the range 250-35 °0 was established. Various amounts of AgN0a in the two melts were titrated potentiometrically with Naz02 and the course of reaction was followed with a silver or an oxygen indicator electrode. Average values for the solubility product of Agz0 in the two melts were calculated from potentials not far removed from the points of equivalence. Agz0 is more soluble in N a N 0 3 - K N 0 3 than in KN0a. Differences in tile heats of dissolution of the oxide in the melts are due to both AG (pKs) and AS. REFERENCES I S. N. FLENGAS AND E. K. RIDEAL, Proc. Roy. Soe. London, A 233 (1956) 442. 2 S. N. FLENGAS, J. Chem. Soe., (1956) 534. 3 M. BLANDER, F. F. BLANKENSHIP AND R. F. NEWTON, J. Phys. Chem., 63 (1959) 1259. 4 J. JORDAN, J. MEIER, E. J. BILLINGHAM AND J'. PENDERGRAST, Anal. Chem., 31 (1959) 1439; 32 (i96o) 651; Nature, 187 (I96O) 318. 5 D. L. MANNING AND M. BLANDER, Inorg. Chem., I (1962) 594. 6 H. T. TIEN AND G. W. HAERINGTON, Inorg. Chem., 2 (1963) 369. 7 M. BLANDER AND E. B. LUCHSlNGER, J. Am. Chem. Soc., 86 (1964) 319 . 8 G. G. BOMBI, M. FIORANI AND G. A. MAZZOCCHIN, J. Electroanal. Chem., 9 (1965) 457. 9 1:~. S. SETHI AND H. C. GAUR, S y m p o s i u m on Electrochemistry, Delhi, October, 1964. IO R. CIGEN AND N. MANNERSTRAND, Acta Chem. Scand., 18 (1964) 1755, 2203. I I H. J. ARNIKAR, D. K. SHARMA AND R. TRIPATHI,]rnd. J. Chem., 3 (1965) 7. 12 H. T. TIEN, J. Phys. Chem., 69 (1965) 3763 . 13 M. FIORANI, G. G. BOMBI AND G. A. MAZZOCCHIN, J. Electroanal. Chem., 13 (1967) 167. 14 Yu. S. LYALIKOV, E. A. LEVlNSON, G. I. TODOROVAAND V. V. NIKOLAEVA, Uch. Za D. Kishinevsk. Sos. Univ., 56 (196o) 91. 15 G. G. BOMBI AND M. FIORANI, Talanta, 12 (1965) lO53. 1 6 H. S. SWOFFORD AND H. A. LAITINEN, J. Electrochem. Soc., i i o (1963) 814. 17 H. L u x , Z. Elektrochem., 45 (1939) 303. 18 H. A. LAITINEN AND B. B. BHATIA, J. Electrochem. Soc., lO 7 (196o) 705 . 19 G. DELARUE, J. Electroanal. Chem., i (196o) 285; Bull. Soe. Chim. France (196o) 1654. 20 A. M. SHAMS E1 DIN AND ~k. A. A. GERGES, J. Electroanal. Chem., 4 (1962) 309. 21 A. M. SHAMS EL DIN AND A.. A. EL I-IOSARY, Electrochim. Acta, in press. 22 Handbook of Chemistry and Physics, 43rd ed., The Chemical R u b b e r Pub. Co., 1961, p. 648. 23 A. M. SHAMS EL DIN, A. A. EL HOSARY AND A. A. A. GERGES, J. Electroanal. Chem., 6 (1963) 131. 24 A. M. SHAMS EL DIN AND A. A. EL HOSARY, Electrochim. Acta, in press. 25 A. M. SHAMS EL DIN AND A. A. EL HOSARY,Electrochim. Acta, in press. 26 R. LORENZ, H. FREI AND A. JABS, Z. Physih. Chim., 61 (19o8) 468. 27 A. M. SHAMS EL DIN AND A. A. GERGES, Electroehim. Acta, 9 (1964) 613. 28 A . M . SHAMS EL DIN AND A. A. SERGES, Electrochemistry, edited b y A. FRIEND AND V. GUTMAN, P e r g a m o n Press, Oxford 1964, p. 562. 29 &. M. SHAMS EL DIN, Electrochim. Acta, 7 (1962) 285.

j . Electroanal. Chem., 17 (1968) 137-143