Solvent extraction of some anions from molten NH4NO3·2H2O with tetraheptyl-ammonium nitrate

J. inorg,nucl.Chem.,1968,Vol.30, pp. 1963to 1969. PergamonPress. Printedin Great Britain

SOLVENT EXTRACTION OF SOME ANIONS FROM MOLTEN NH4NO3"2H20 WITH TETRAHEPTYLAMMONIUM NITRATE R. M. NIKOLI(~ and 1..I. G A L The Boris Kidri6 Institute of Nuclear Sciences, Belgrade, Yugoslavia

(Received4 December 1967) Abstract--The distribution of anions between an aqueous melt of composition NH4NOz.2H20 and an organic solution of tetraheptylammonium nitrate in a biphenyl-naphthalene diluent was measured at 55, 70 and 85°C. At low concentrations of solutes in both phases, a simple anion-exchange NH4Y + TNO:~(org.) = TY(org.)+ N H4NO:~. ( T * = tetraheptylammonium, Y - = anionic species) governs the distribution. For Y - = CI , Br-, ReO4-, AgCIz-, and AgBr2-, equilibrium constants for the anion-exchange reaction were determined. The results are discussed with respect to a simple thermodynamic model for the distribution of charged species between a quasi-crystalline liquid electrolyte and an organic solvent.

THE DISTRIBUTION of anions between an aqueous solution and an organic solvent

containing a quarternary alkylammonium salt was interpreted by Scibona et al.[ 1] in terms of a simple anion-exchange equilibrium between the two phases. They have also shown that some thermodynamic quantities for this equilibrium can be calculated by means of an "electrostatic model" based on Born's charging expression and the dissociation energy of ion-pairs. Recently it was found that tetraheptylammonium nitrate in a polyphenyl diluent extracts anions from an anhydrous lithium-potassium nitrate melt at 140-160°C[2]. In this case too, the distribution of anions between the two phases mainly depends on the anionexchange equilibrium, but the heterogenous equilibrium constant cannot be reproduced by the above model. The present work investigates the distribution of some anions between an aqueous melt of composition NH4NO3-2H20 and an organic solution of tetraheptylammonium nitrate in a biphenyl-naphthalene mixture. The melt, which is liquid above 40°C, is intermediate between concentrated aqueous electrolyte solutions and anhydrous molten salts. Such a system can furnish additional information as to whether the same basic equilibrium governs the distribution of anions irrespective of the nature of the electrolyte phase. It would also be desirable to develop a suitable model for the distribution of charged species between a melt and an organic solvent containing ion-pairs. EXPERIMENTAL

Reagents and radioactive tracers. Tetraheptylammonium nitrate was prepared from tetraheptylammonium iodide (Eastman White Label) as described earlier[2]. The organic diluent is an eutectic 1. G. Scibona, J. E. Byrum, K. Kimura andJ. W. lrvine,Jr, SolventExtractionChemistry(Proceedings of the Gothenburg Conference), pp. 398-407. North-Holland, Amsterdam (1967). 2. I.J. Gal, J. Mendez andJ. W. lrvine,Jr, inorg. Chem. In press. 3. V. H. Troutner, U.S. Atomic Energy Commission, Report No. HW-57431 (1958). 1963

1964

R. M. NIKOLI(~ and I. J. G A L

mixture of biphenyl and naphthalene (55 mole per cent biphenyl) which melts at 39.5°C[3]. At temperatures used in the present investigation, this mixture behaves toward the nitrate melt as a chemically inert diluent of low vapor pressure. The aqueous melt of composition NH4NO3"2H20 was prepared from dry ammonium nitrate (reagent grade) and distilled water. Radioactive tracers in the form of NH~S2Br, NH43sCI, NH41S°ReO~ and ~°mAgNO3 were introduced in the melt. Their preparation and purity are described elsewhere[l, 2]. Into some experiments inactive NH4CI or NH4Br was added to the melt. Distribution measurements. Equal weights, usually about 3 g, of the organic solution and the melt were equilibrated in a glass stoppered tube. Several tubes were placed in a thermostat, and good mixing of the phases was obtained by the vibrating-motion of a mechanical arm on which the glass tubes were fixed. The equilibration temperatures were 55, 70 and 85°C, and the equilibration time was 30 rain. Preliminary experiments have shown that in all cases equilibrium is reached after at least 20 rain. The phases separated easily after standing for some time in the thermostat. Firstly, samples of the organic phase were taken in vials, then the whole organic phase was withdrawn with a glass pipette, and finally samples of the melt were taken. After weighing, the melt samples were dissolved in water (in the case of Ag, in 6 M ammonia), and the organic samples in benzene. The samples were counted in a well-type gamma scintillation counter. The molal distribution coefficient, D, was calculated as the ratio of the counting rate per gram organic diluent to the counting rate per gram nitrate melt. Every distribution coefficient is the average value of at least two independent equilibration experiments. Usually an agreement in D values within 1-5 per cent was obtained, except in the range 10 -~ to 10 -4, where discrepancies up to 10 per cent were accepted. All concentrations are expressed in molality. RESULTS

AND DISCUSSION

The molal distribution coefficients (D) of CI-, Br% ReO4-, and of the silver halide complexes AgCIz- and AgBrz- depend on the tetraheptylammonium nitrate molality (mr) of the organic solution as shown in Table 1. Silver chloride and bromide were extracted in the presence of 5.10 -2 molal NH4CI and 1" 10 -2 molal NH4Br, respectively, Solubility data of silver halides in the N H 4 N O s ' 2 H 2 0 melt[4] indicate that under these conditions the silver tracer is completely in the form of an anionic complex. A series of experiments (not presented in Table 1) has shown that the distribution coefficients of chloride and bromide do not depend on the concentration of these anions in the melt up to a molality of 0.1. The distribution coefficients of these halides are low enough so that in the case of AgCI~- and AgBr2- extraction the quarternary salt practically remains in the nitrate form. A distinct temperature effect in the extraction of R e O c and AgBr2- was observed, while in the case of bromide it was negligible. The same is probably true for chloride, which has a very small value of D. The sequence of extraction obviously depends on the size of anions: Cl- < Br- < R e O 4 - , and also AgCI~- < AgBr2-.

Equilibrium constants and thermodynamic quantities By analogy with extractions from dilute aqueous solutions and from molten nitrate salts, it can be assumed that the distribution of an anion Y - between the two phases depends on the ion-exchange reaction NH4Y + TNO3(o) ----TY(o) + NH4NO3. 4. I.J. Gal, To be published.

(1)

1965

Solvent extraction of some anions T a b l e 1. - T h e d i s t r i b u t i o n c o e f f i c i e n t s ( D ) a n d t h e e q u i l i b r i u m q u o t i e n t s ( Q ) I n i t i a l m o l a l i t i e s in m e l t : C I - 1" 10 -3, B r - 1-10 -3, R e O 4 - 1' 10 -3, A g ÷ 1" 10 -~

D

m t

(molality)

55 °

70 °

Q 85 °

55 °

70 °

85 °

C11.00.10 -3 5 . 0 1 - 1 0 -2 2 ' 0 4 - 1 0 -2

3 . 8 . 1 0 -4 2 . 3 . 1 0 -4 I" 1"10-4

0.03 0-04 0.05

Br-

1.00.10 -1 5 . 0 7 . 1 0 -2 2 . 1 3 . 1 0 -2 1.00.10 -2 5 . 0 1 . 1 0 -3

9 . 3 2 . 1 0 -3 4 . 8 7 . 1 0 -3 2 . 5 0 . 1 0 -3 1.39.10 -3 7.1 -10 -4

8 . 9 2 . 1 0 -3 4"92" 10 -z 2 . 4 3 . 1 0 -3 1.40.10 -3 7-4 .10 -4

8 . 9 3 . 1 0 -3 4 . 6 1 . 1 0 -3 2 . 2 0 . 1 0 -3 1-26.10 -3 6.5 -10 -4

0-80 0-83 1-0 1-2 1.2

0-77 0-84 0-98 1-2 1.3

0.77 0.78 0"91 1-1 1.1

7.92 3.84 1.48 6 . 1 4 . 1 0 -1 2 . 8 6 . 1 0 -1 1.05.10 -1 5"05" 10 -2 2 . 5 4 - 1 0 -2 9 . 6 6 - 1 0 -3

6.03 974 683 2.80 948 667 1.03 849 611 4 . 5 9 . 1 0 -3 778 523 2 . 0 7 . 1 0 -~ 694 494 7 . 8 0 . 1 0 -2 604 450 3"85-10 -2 600 436 1-98' 10 -2 603 435 8 - 0 1 . 1 0 -3 600 410 A g C I 2 - ( m e l t 5 . 1 0 -2 m o l a l in N H 4 C I )

520 488 427 396 358 334 332 339 340

2.47 1-25 2 . 4 6 . 1 0 -I 1.07.10 -1 4"42" 10 -2 2 . 2 0 . 1 0 -2 1 . 1 0 . 1 0 -2 4 . 4 7 . 1 0 -3

213 214 212 185 190 190 189 190 A g B r 2 - ( m e l t 1.10 -2 m o l a l in N H 4 B r )

ReO4-

9 . 9 8 . 1 0 -2 4 . 9 5 . 1 0 -2 2 . 0 9 . 1 0 -2 9 . 9 8 . 1 0 -3 4 - 9 9 . 1 0 -:3 2.01.10-3 9-98" 10 -4 5-03" 10 -4 2 . 0 3 . 1 0 -4

11.3 5.48 2.06 9 . 0 1 . 1 0 -1 4 . 0 2 . 1 0 -1 1 . 4 1 . 1 0 -1 6 - 9 5 . 1 0 -2 3.52" 10 -2 1.41.10 -2

1-00-10 -1 5 . 0 1 . 1 0 -2 1.00. I0 -z 5-00-10 -3 2 . 0 0 . 1 0 -3 9.96.10-4 4.99.10-4 2.10.10-4

1.00-10 -1 5 . 0 1 . 1 0 -z 2 . 0 0 . 1 0 -z 1.00.10 -3 5 . 0 0 . 1 0 -3 2 . 0 0 - 1 0 -3 1.00.10 3 5 . 1 1 . 1 0 -4 1"99"10-4

390 185 76.1 38-5 17"6 7-10 3-75 1.92 0"691

175 86.1 42.9 21-9 10-0 4-11 2"17 1-09 0"397

115 65.2 26.3 13'1 5'89 2-40 1.27 0.644 0"232

3-4.104 3"2"104 3'2'104 3"3"104 3"0"104 3-1-104 3-2-104 3-2.104 3"0"104

1-5" 104 1"5"104 1"8'104 1"9'104 1"7"104 1"8"104 1"9"104 1-8-104 1"7"104

1 ' 0 ' 104 1"1"104 1"1"104 I'1"104 1"0"104 1'0'104 1"1"104 1"1"104 1"0"104

1966

R. M. NIKOLI(~ and I. J. GAL

Here, T ÷ is the tetraheptylammonium cation, and the subscript (o) denotes the species in the organic phase. Equation (1) cannot interpret the distribution data at all concentrations of the quarternary salt (which tends to associate in organic solvents of low dielectric constant). However, at low concentration of all solutes in both phases, it can be assumed that Equation (1) holds. Defining the distribution coefficient of the anion Y- by D = (TYro)) / (NH4Y) the equilibrium constant for the ion-exchange reaction equals K = Dm"FoF = QFoF mt

(2)

Here, Q = Dm,/mt is an equifibrium quotient, mn and mt are the molalities of NH4NOa and TNO3to) in the melt and organic phase, respectively, while Fo and F are activity coefficient ratios of the species in the organic phase and melt. From the thermodynamic viewpoint, it is perfectly acceptable to include all solutesolute interactions in the organic phase, even the association of ion-pairs, into Fo. However, adequate reference states for the activity coefficients must be chosen. As usual, it will be assumed that Fo and F tend to unity when all solutes in both phases are extremely diluted. (This also implies that the activity coefficient of NH4NO3 in the pure melt equals unity). In that case, FoF = 1 and K----Q. Table 1 shows that for mt values below 10 -2, the equilibrium quotient Q tends to a constant value. Assuming then FoF - 1, the equilibrium constant K can be evaluated. In the case of chloride, the distribution coefficients are too low for extrapolating Q so that only a rough estimate of K is possible. In evaluating the equilibrium constants the F value (for the melt phase) is assumed to equal unity. This is probably true within the experimental error when the solute concentration in the melt is below 10 -2 molal[2]. In the case of AgClzextraction, owing to the presence of 5"10 -2 molal NH4CI, the F value, although constant, might be slightly different from unity. Unfortunately, the activity coefficients of halides in the melt are not known, and no correction of the K value for AgCI~- can be made at present. The equilibrium constants, the standard free energy, enthalpy and entropy changes for the anion-exchange equilibria are summarized in Table 2. The thermodynamic quantities were calculated as usual: AG ° = - R T I n K , AH ° = --R d(lnK)/d (l/T), and AS° = (AH ° - AG°)/T. Plots of log K vs. I/T are presented in Fig. 1. A comparison of the data in Table 2 with those obtained in aqueous solutions at 25°C show, in general, that the thermodynamic quantities are of the same order of magnitude and sign, but the equilibrium constants are considerably lower for the mett. The latter fact is in qualitative agreement with the negative value of AH °. In the case of ReO4- extraction, for instance, the data in Table 2 should be compared with K = 1.2.103 and AH ° -------7.1 kcal/mole obtained at 25°C from aqueous solutions with tetraheptylammonium nitrate in benzene [1]. A simple thermodynamic model. The "electrostatic model" proposed to evaluate AG ° on a theoretical basis [1] is hardly applicable to equilibria involving melts. One reason for its inadequacy is the difference in structure between a

Solvent extraction of some anions

1967

,.5 t

,0[

3"0

2"5

1

~ O J

2!0

J

I

....,-- • " " "

2"9

3"0

,?

3"~

Fig. 1. Plots of log K vs. I/T for ReO4- and AgBr2- extraction.

molten salt and a dilute aqueous solution. It is felt that the quasi-crystalline structure of melts and concentrated aqueous solutions requires another approach in developing a suitable model. The standard free energy change for Equation (1), for instance, can be computed by means of the following thermodynamic steps: (l) The molecule NH4Y is transferred from the quasi-lattice of the melt into the gas phase and then completely dissociated into NH4 ÷ and Y - at infinite separation. The molal free energy change for this step is equal to the potential energy increase given by Born's expression:

Here, A is the Madelung constant for the "lattice", N Avogadro's number, e is the elementary charge of an ion, rnm + rv is the interionic separation given as the sum of cation and anion radii, and n is the repulsive potential exponent*. • Small contributions due to the van der Waais and other terms can be neglected, particularly if only the difference in energy between NH4Y and NH4NO.~ is required.

R. M. N I K O L I t ~ and I. J. G A L

1968

Table 2. - Equilibrium constants (K) and thermodynamic quantities Y- *

°C

K

AG ° 343 (kcal/mole)

CI-

70

-0.05

+2

Br-

55 70 85

1.2 1.2 1.1

-1.24

55 70 85

600 430 335

--4.13

AgCI2-

70

190

--3.58

AgBr2-

55 70 85

3'2"104 1.8.104 1.0.104

-6.68

ReO4-

A[-I°343 (kcal/mole)

A5%43 (e.u.)

-4.49

--1.0

-9.66

--8.7

*See Equation (1). (2) In the next step, Y - is transferred from the gas phase into the organic phase. The contribution to the molal free energy is equal to (Ne2/2rv)(1/eo- 1). Here, ~o is the dielectric constant of the organic phase. (3) Y - in the organic phase, initially at infinite distance d® from the cation T ÷, forms an ion-pair TY at equilibrium separation ry + rT. The molal free energy change for this step is

13

~o \ry + rT

(4-6) In an identical way, three reverse steps are required to transfer the nitrate ion from the organic phase into the melt. The contributions to the molal free energy are:

N_e~[ 1 'o \d®

1

\

Ne2 /

1

rNo3-+-rr)+2--~o3~l--~)

ANe 2 ( 1 _ ! I rNH,+rNO3\ n/"

Summing up the six terms and taking into account that K = exp (--G°/RT), this model gives logK=

Ne2[(

1)( 1

2 . 3 - ~ R T A I-- n

rNH~-~re

1

-

) (1

rNH 4 q- rNO3

1

+ ---(.--Eo r T -~- rNOa

1 ) r T -~ r v

+0 tl 1)tl ] We propose to test Equation (3) for Y - = CI-, Br-, ReO4-, and compare the K values with the experimental equilibrium constants in Table 2. (For linear AgCI2and AgBr 2- no adequate values for the "radii" are available). The following numerical values of the parameters are substituted in Equation (3).

Solvent extraction of some anions

1969

The dielectric constant and ionic radii, eo = 2.53 (both biphenyl and naphthalene have the same dielectric constant 2.53 at 70°C [5]). The Pauling radii rc~---1.81, rBr = 1"95, rNH,= 1"43, and the recently reported values[l] of rNo.~= 2.01, rp~o, = 2.45 and ra- = 5.80 (in Angstrom units) are used. The Madelung constant. Standard values reported for solid crystals obviously cannot be applied to a quasi-crystalline lattice of liquids. Diffraction data of molten salts [6] have demonstrated that around each ion, in average, (4-6) nearest neighbours of opposite charge are symmetrically packed. The average anion-cation distance is, surprisingly, usually shorter than in the corresponding solid near the m.p. However, at larger distances the structure around the ion is "smeared out". Taking into account the predominantly short-range order, let us assume that a highly symmetric NaCI type of lattice prevails at a short distance around an ion. Considering only the first nearest cube surrounding an anion, it can be shown that it intercepts six cations on the cube faces, twelve anions on the cube edges and eight cations at the cube corners. Hence, A = 6/2--12/4~/2+ 8/8V'3 = 1.45. As a first approximation, this value seems to be more realistic for a melt than the standard Madelung constants for crystals. (Strictly speaking, the model does not take into account that the next-nearest neighbours of an Y- ion, i.e. the NO:lions, differ in size from Y-. However, for the present purpose such fine details in the next-nearest neighbours interaction are neglected.) For the repulsive potential exponent, n = 9 was adopted [7]. Table 3 . - C o m p a r i s o n of equilibrium constants (70 °) K y*

CIBrReO4-

exp.

theor.

-0.05 1.2 430

0-07 0.6 400

*See Equation ( 1).

In Table 3, K values calculated by means of Equation (3) are compared with those obtained experimentally. As seen, the model predicts the right sequence of extraction of anions and at least the order of magnitude of the equilibrium constant. In the case of ReO4- a very good agreement is obtained, but this is probably fortuitous. Although Equation (3) is based on a crude and necessarily approximate model, it is felt that the accuracy in computing the equilibrium constants can be increased by refining some parameters, especially the Madelung constant and the interionic distances. 5. C. D. Hodgman (Editor), Handbook of Chemistry and Physics. Chem. Rubber Publ. Co., Cleveland (1964).

6. M. Blander (Editor), Molten Salt Chemistry, Chap. 2.1nterscience, New York (1964). 7. E.A. Moelwyn-Hughes, Physical Chemistry, 2nd Edn. Pergamon Press, Oxford, 1964.