3711
ASSOCIATION EQUILIBRIA OF SILVER AND CHLORIDE IONS
Association Equilibria of Silver and Chloride Ions in Liquid Ammonium Nitratewater Mixtures.
11. The Anhydrous Melt
by Mordechai Peleg Department of Inorganic and Analytical Chemistry, The Hebrew University of Jerusalem, Jerusalem, Israel (Receioed January 6, 1971) Publication costs borne completely by The Journal of Phusical Chemistry
The electromotive force of the cells
(M = Na or K) as a function of silver nitrate and ammonium chloride has been measured a t 150" for the mixture and a t 160" for the ammonium ammonium nitrate-sodium nitrate (80-20 g, 81 mol % "4NOa) nitrate-potassiurn nitrate (86-14 g, 89 mol % NH4N0,) mixture. By applyiag the quasi-lattice model for inolteii salts the cation-ligand interaction energy for the Ag+-Cl- pair was evaluated as -5.01 rt 0.05 kcal/ mol for the ammonium nitrate-sodium nitrate mixture and -5.10 0.2 kcal/mol for the ammonium iiitrate-potassium nitrate mixture. The value obtained by extrapolation from concentrated aqueous am0.1 kcallmol) is in fair agreement. monium nitrate solution (-4.9
*
*
Introduction In a previous paper,' we reported measurements for the association constants of the silver ion with the chloride ion in liquid ammonium nitrate-water mixtures = 0.4-1.4) at a tem(mole ratio range, H20:"*S03 perature of 110". The results were discussed in terms of a quasi-lattice model based on that of molten salts as proposed by Braunstein2t3to explain hydration and association equilibria in these highly concentrated aqueous solutions. By applying the equation developed by Braunstein3and extrapolating to zero water content it is possible to obtain a value for the cation-ligand interaction energy. From the above mentioned results a value of 4.9 f 0.5 kcal/mol was obtained for the Agf-C1- complex in anhydrous ammonium nitrate. It was thus of interest to attempt to measure the association constant for the Ag+-C1- pair directly in anhydrous ammonium nitrate and to compare this value with that obtained by applying the quasi-lattice model of Braunstein to the results obtained from the concentrated aqueous ammonium nitrate solutions. While results have been reported for Ag+-Cl- complex formation for most of the alkali metal nitrates and their mixture^,^ to the best of the author's knowledge no results have been published for molten ammonium nitrate or its mixtures. Due to practical difficulties it was n o t possible to measure Ag+-Cl- complex formation in the pure ammonium nitrate system due to corrosion of the silver electrode. However, measurements were possible in mixtures of ammonium nitrate-sodium nitrate and ammonium nitrate-potassium nitrate.
Ag
",NO8 MNOa AgIYOa
1 *:E 1 NHdC1
~
Ag
(1) M. Peleg, J . Phys. Chem., 75, 2060 (1971). (2) J. Braunstein in "Ionic Interactions: Dilute Solutions to Molten Salts," 8. Petrucci, Ed., Academic Press, New York, N . Y., 1971,
Chapter 4. (3) J. Braunstein, J . Phys. Chem., 71, 3402 (1967). (4) Y . T. Hsu, R. B. Escue, and T. H. Tidwell, Jr., J. Electroanal. Chem., 15, 245 (1967). The Journal of Physical Chemistrv, Vol. '76,No. 24,1971
37 12
MORDECHAI PELEG
44.2:6.9:48.9 mol To, mp 14205) controlled to within =to.1". A Keithley 660A differential voltmeter was used to measure the emf of the cell. Preliminary experiments using pure ammonium nitrate at a temperature just above that of its melting point (mp 169.2") gave unstable readings and a continuous drift. This was due to the corrosive effect of the ammonium nitrate melt. Barclay and Crewe6 have studied the thermal decomposition of ammonium nitrate and, while the mechanism of decomposition is complex, it appears that one of the products formed is HNOa. Chloride ions also have been shown to catalyze the thermal decomposition of ammonium nitrate.' A trial experiment carried out on an ammonium nitratesodium nitrate (80-20 g) mixture at 150" showed that a weighed piece of silver placed in this melt under exactly the same experimental conditions as were employed for the emf measurements remained uncorroded after 3 days. Similar experiments using an ammonium nitrate-potassium nitrate (86-14 g) mixture at 160" gave evidence of slight corrosion over a prolonged period of time. Measurements were therefore made using the ammonium nitrate-sodium nitrate mixture as the solvent although a few runs were performed on the ammonium nitrate-potassium nitrate melt. Owing to the low values of the solubility product for silver chloride (-lo+ mole ratio) in the solvents studied,8 low concentrations of the silver ion and the chloride ion had to be employed. The silver ion concentration was thus limited to a range between about 1 X lov4and 3 X (mol/mol of solvent melt), while the chloride concentration was varied between 1 X and 6 X (mol/mol of solvent melt). Preliminary experiments showed that the Nernst equation was obeyed with respect to the stoichiometric concentration of silver in the range under examination (Figure 1). The concentration units are given as mole ratios, R
Rci-
=
R N H C= ~
~ N H , C ~ nNHiNOa
+
nMN0a
where M represents either Na or I< and n is the number of moles of AgKOa,P\'H4N03,;i\lN03,and T\"*Cl. Procedure. A weighed amount of ammonium nitrate-sodium (or potassium) nitrate was added t o the cell together with a known amount of silver nitrate. The cell was stoppered and suspended in the molten salt bath and a stream of dry nitrogen was allowed to pass over the melt. After attainment of thermal equilibrium, the empty reference compartment was placed in position and filled by applying a vacuum to the tube. The indicating silver wire and reference electrode wire were then introduced into their respective compartments. After about 1 hr the potential The Journal of Physical Chemistry, Val. 75, No. 34, 1971
eo 70
-
60
-
50
-
LO
-
-20
- 3.0
-LO
' 9
[RAd]
Figure 1. Typical Nernst plot for silver-silver ion electrode in ammonium nitrate-sodium nitrate solvent (81 mol % ammonium nitrate) at 150".
between the two compartments steadied and showed only a very slight drift of less than 1 mV/hr. This drift was probably due to the very slight evaporation of ammonium nitrate from the solution. This drift was compensated for in the calculations, the magnitude of the correction approaching about 10-15'3,) depending on the time between additions of chloride ions.
Results The emf's of the concentration cell were measured as a function of changing silver ion concentration a t a temperature of 150" for the ammonium nitrate-sodium nitrate mixture and at 160" for the ammonium nitratepotassium nitrate mixture. Owing to the low concentration of silver ions in solution, the change of cell emf on addition of chloride ions to the indicating compartment is related to the stoichiometric activity coeecient of silver nitrate by the formula
where AB is the difference between the emf of the cell in the presence and absence of chloride ions.g Thermodynamic equilibrium constants for association reaction Ag+ C1- = Ag+-Cl- were calculated by the graphical analysis of the activity coefficients as described by Braunstein, et aL1O
+
( 5 ) T. Alexander, Jr., and S. G. Hindin. Ind. Eng. Chem., 39, 1044
(1967). (6) K . S. Barolay and J . M . Crewe, J . Appl. Chem., 17,21 (1967). (7) A. G. Keenan and B. Dimitriades, J . Chem. Phys., 37, 1683 (1962). (8) M. Peleg, to he published. (9) J. &I. C. Hess, J. Braunstein, and H. Braunstein, J. Inorg. Nud. Chem., 26, 811 (1964). (10) J. Braunstein, M. Blander, and R. M. Lindgren, J . Amer. Chem. Soc., 84, 1529 (1962).
ASBOCIATION EQUILIBRIA OF SILVER AND CHLORIDE IONS
3713
or
K1 = -1im
So
RAgNOadO
where
so=
lim RNErCl-0
(a In
YAgNOa
aRNH4Cl
-
)
0, 0
RAgNOs
Typical plots of -log YAgNOa against Rcl- for the ammonium nitrate-sodium nitrate solvent are shown in Figure 2. The potentiometric data appear as supplementary material in the microfilm edition.'l The slopes of'the plots are different from those typically obtained (cf. the shape of a typical plot in Figure 3 of ref 1). The sudden change in slope is probably due to precipitation of silver chloride as the chloride concentration increases. Although the solubility product "constants" ( 1 ' 9 ~ ) are not constant and increase with increasing silver ion content, their variation is not large and may be accounted for, as suggested by Braunstein,lZa by the heat sink a t the electrode-solution interface and/ or by the possibility that the solid phase is not pure AgCl but a solid solution. At higher chloride ion concentration corrosion might also start to occur as suggested by Inman, et a,1.,12b since the chloride ion concentration may reduce the silver ion concentration below a level where corrosion is so significant as to affect the validity of the Nernst equation. However, the initial portions of the curves are valid and could be analyzed using the graphical method (Figure 3) described earlier and this produced a value of K1 = 2120 f 100 (mol/mol of solvent)-l for the ammonium nitrate-sodium nitrate mixture (80-20 g) at a temperature of 150". The equation developed by Blander, et al., l 3 based on the asymmetric quasi-lattice model for molten salts was utilized t o evaluate the Ag-C1 interaction energy K1 = Z ( a
-
1)
(1)
where a = exp(-e,/kT),' e, being the cation-ligand interaction energy and 2 the coordination number, generally chosen as 6. For the mixed nitrate solvent e, was -4.93 f 0.05 kcal/mol. It has been concluded (ref 13, p 230) from an examination of the data available for Ag-C1 complex formation in mixed molten alkali nitrates that e, is linearly proportional to the composition of the binary mixture melts; i.e., the equation €,(mix)
=
N'e,'(soIvent 1)
+ N"e,"(solvent
2)
(2)
will probably hold in every case. Using eq 2 for the present solvent mixture and taking e, for Ag-Cl in pure molten NaN03 to be -4.59 kcal/m01,'~a value of 5.01
2
0
8
6
4
Rc,-
0
1
2
10
IO 5
3 L R A ~~1 0
Figure 3. Limiting slope Sovs. R A ~ in + NHdN0a-NaN03 a t 150". solvent (81 mol % "4NOa) f 0.05 kcal/mol was obtained €or the Ag-C1 interaction energy. Similarly, from the ammonium nitrate-potassium nitrate solvent at 160°, E , for the Ag-C1 complex in pure ammonium nitrate was calculated to be -5.10 =k 0.2 kcal/mol (K1 = 2400 f 400 (mol/mol of solvent)-'), E , = 5.16 f 0.1 kcal/mol for the mixed solvent, while e, in pure potassium nitrate was taken t o be -5,62 kcal/mol.15 The larger value of the uncertainty range was due to the greater instability and drift in the experimental measurements; however, the e, value over-
(11) The potentiometric data appear immediately following this article in the microfilm edition of this volume of the journal. Single copies may be obtained from the Reprint Department, ACS Publications, 1155 Sixteenth St., N.W., Washington, D. C. 20036. Remit check or money order for $3.00 for photocopy or $2.00 for microfiche. (12) (a) J. Braunstein, private communication; (b) D. Inman, B. Jones, and 8 . H. White, J. Inorg. Nucl. Chem., 32, 927 (1970). (13) M. Blander in "Molten Salt Chemistry," M. Blander, Ed., Interscience, New York, N. Y., 1964. (14) D. G. Hill, J. Braunstein, and M. Blander, J. Phys. Chem., 64, 1038 (1960). (15) D. L. Manning, J. Braunstein, and .M.Blander, ibid., 66, 2069 (1962). The Journal of Physical Chemistry, Vol. 76, N o . 94, 1971
MORDECHAI PELEG
37 14 4.0
3.0
Z.,$ Kl 20
)
Figure 4. Plot of Z / K 1 os. mixtures at T = 110'.
05
RH,O
10 15 RH20
for NHrNOa-HsO
laps the value estimated from the NH4NOS-NaN03 data. Assuming a value of e, = -5.01 f 0.1 kcal/mol for the cation-ligand interaction energy to be more correct, a slight difference exists between this value and that obtained by extrapolation from concentrated aqueous ammonium nitrate solution (eo = -4.9 =k 0.5 kcal/ mol)' but the estimates do overlap. Braunstein3 developed the following equation to explain complex formation in concentrated aqueous solution Z/Ki
=
+(P/~BH
(3)
where a! = = e-@", EB is the cation-water interaction energy, and RH is the mole ratio of water. If eq 3 is now replotted using the value for ec obtained from the anhydrous melt, together with the results obtained previously1 from aqueous solutions, a graph as shown in Figure 4 is obtained. A straight line as was suggested earlier' for the region between RH = 0 and RH = 1 is not realized although in quite good correlation with the aqueous data. This suggests that the equilibrium constants reported earlier' had not yet reached the limiting behavior predicted by eq 3, which is only a limiting equation for low water contentsalthough it applies occasionally, empirically, to higher
The Journal of Physical Chemistru, Vol. 7 6 , N o . 84, 1971
water contents.2 Braunstein's assumption2,a in developing eq 3 from the quasi-lattice model of molten salts was that the water dipoles sit on the anion sublattice; i.e., only the cations are hydrated. Presumably in the present case the nitrate anion competes with the ammonium cation for the water molecule and thus eq 3 does not hold. Other factors which might cause deviations from eq 3 include the possibility of change of coordination number and nonadditivity of pairwise interaction.2ga A measure of the limiting slope of the Z / K vs. R H ~plot O leads t o the value e~ = -0.5 kcal/ mol as compared to the previous estimated value of EH = +0.5 kcal/mol.l A value for K12(Ag2CI+)of around 870 can be estimated from the slope plot in Figure 3. The ionic radii of Na, K, and NH4+are usually taken to be 0.98, 1.33, and 1.45 8,respectively, based on X-ray crystallographic studies. Thus it might be expected in the light of the reciprocal Coulomb effect16 that the stability of the Ag-Cl complex in molten ammonium nitrate would be about the same as, or slightly larger than, the value in molten potassium nitrate. The results indicate that the value of ec in ammonium nitrate (-5.01) falls in between those for sodium nitrate (-4.59) and potassium nitrate (-5.62). Thus it appears that the effective radius of the ammonium ion in the molten nitrate is less than that of the potassium ion. This is substantiated by the conclusion drawn by Braunstein, et ul.,l7 that although K+ and NH4+ have the same effective radius in dilute aqueous solution at 25", the activity of water at high temperatures and high salt concentrations in ammonium nitratela is closer to that in sodium nitrate than to that in potassium nitrnte.lg
Acknowledgment. The author wishes to thank Professor J. Braunstein for helpful suggestions. (16) D. L. Manning, R. C. Bansal, J. Braunstein, and M. Blander, J . Amer. Chem. SOC.,84, 2028 (1962) (17) J. Braunstein and H. Braunstein, Inorg. Chem., 8, 1558 (1969). (18) A. N. Campbell, J. B. Fishman, G. Rutherford, T. P. Schaefer, and L. Ross, Can. J . Chem., 34, 151 (1956). (19) L. P. Shpigel and K. P. Misohenko, J . A p p l . Chem. USSR, 40, 659 (1967). I
