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Distribution of nitric acid between water and organic solvents, containing tri-n-octyl amine—II
J. Inorg.Nucl. Chem.,1964,Vol.26, pp. 1085to 1102. PergamonPressLtd. Printedin NorthernIreland
DISTRIBUTION OF NITRIC ACID BETWEEN WATER A N D O R G A N I C SOLVENTS, C O N T A I N I N G TRI-N-OCTYL AMINE--II J. M. P. J. VERSTEGEN Institutt for Atomenergi, Kjeller, Norway (Received 5 April 1963; in revised form 4 July 1963) Abstract--Previous measurements on the distribution of HNOs between water and organic solvents containing TOA are extended. Viscosity and conductivity of the equilibrated amine phases can be explained by ion-pair behaviour in solvents of low dielectric constant. Formation of micelles or colloids is not likely. Hydration occurs and solvation was found in chloroform solutions. It is concluded that a compound (TOAH)NOfHNO3 is formed by hydrogen bonding and that the equilibrium constant depends on the solvent, but not significantly on the temperature. IT was s h o w n earlier ~1) that the d i s t r i b u t i o n of H N O 3 between water a n d organic solvents c o n t a i n i n g t r i - n - o c t y l amine (TOA) at low acidities can be described in terms of i o n - p a i r ( a m i n e - n i t r a t e ) f o r m a t i o n in the organic phase. After stoichiometric n e u t r a l i z a t i o n of the amine, further extraction of acid by the a m i n e occurs3 ~-4) SHEVCHENKO et al. <5) reached the c o n c l u s i o n that a n a m i n e nitrate complex with the c o m p o s i t i o n 1:2 exists. BERTOCCI a n d ROLANDIt6) concluded t h a t no stoichiometric c o m p o u n d between T O A a n d H N O 3 is formed. T~LAT-ERBE~c7) assumed that a series of complexes is b u i l t u p by stepwise equilibria. D u r i n g p r e p a r a t i o n of this paper BARONCELLI et al. ~8) concluded t h a t the d i s t r i b u t i o n of H N O a in the range of high acidities could be ascribed to dissolution of the acid in the amine nitrate, witho u t further c o m p o u n d formation. EXPERIMENTAL Reagents TOA was vacuum distilled and titrated in acetic acid. `9' The molecular weight found was 352"7. Organic TOA phases are in general colourless after shaking with HNOs, but at high acidities a brown or pink solution was often observed. The measured refractive index was n~5 = 1.4494, while the undistiUed sample gave n~ = 1"4491. The density of TOA at different temperatures is given in Table 1. Densities of TOA-benzene and TOA--chloroform mixtures and refractive indexes of the first are shown in Fig. 1 and 2. Fig. 1 shows agreement with data in similar systems,tx°~ Deviations of the densities from ideality are small. i1) j. M. P. J. VERSTEGEN,Trans. Faraday Soc. 58, 1878 (1962). t2) M. DE TRENTINIANand A. CHESNE, Compte Rendu du Colloque de Extraction par Solvent (Madrid, 1959). la) U . BERTOCCI,AERE-R 2933 (1959). ,4~ D. J. CARSWELLand J. J. LAWRANC~,J. Inorg. Nucl. Chem. 11, 69 (1959). ts) V. B. SREVCHENKO,V. S. SPAvllDT,E. A. NENAROKOMOWand K. A. PETROW,Zh. Neorg. Khim. 5, 1852 (1960). ca) U. BERTOCCIand G. ROLANDI,3". Inorg. Nucl. Chem. 23, 323 (1961). ~7, M. Tf~J~AT-ERBEN,Proceedings of the 7th L C.C.C. Stockholm/Uppsala (1962). ,s, F. BARONCF.LLI,G. SCmONAand M. ZIFFERERO,J. Inorg. NucL Chem. 24, 403 (1962). ~9~C. D. WAGNER,R. H. BROWNand E. D. PETERS,J. Amer. Chem. Soc. 69, 2609 (1947). ~t0, K. v. IPENaURG,Thesis, Delft (1962). 1085
1086
J.M.P.J.
VERSTEGEN
TABLE ] ,--DENSITY OF TOA AT DIFFERENTTEMPERATURES Temp. (°C) Density (g/ml)
25.1
35"4
49.9
61"2
75"8
0.80585
0.79950
0.78980
0.78370
0.77385
0.870 i 0.860
1.52
0.85(;
1.51
0.840
1,50
-. 0.830 E E
"-..'--...
1.49
,m 0.820
1.48
~e 0.810 r~ 0.800
1.47 1.46
0.790
1.45 0 100
20 ' 80
4 0' 60
8'0
6 'o 40
20
I00 0
Jo
Vol % TOA Vol % CB Hs
FIG. 1.--Density in g/ml (solid line, left hand scale) and n,Vo (dashed line, right hand
scale) plotted against vol. per cent of the components for the system TOA-benzene. Additivity lines included.
1.5o 140 "-.
130
~'
Jzo
u)
1.10
123 1.00
0.90 0.80 IO0
!
I
/
25 75
5O 50
75 Z5
JO0 VOI % TOA 0 Vol % CHCI 3
FfG. 2.--Density in g/ml plotted against vol. per cent of the components for the system TOA-chloroform. Additivity line excluded.
Distribution of nitric acid--II
1087
Fatuilibrations Equilibrations were performed at 25"0°C but some measurements were also made at 4°C, to study heat effects. Organic phase HNOs concentration is that of Co and total amine concentration by Y.TOA. The range C0/Y~TOA < 1 defines ion-pair formation, while 1 < C0/~TOA < 2 is the region of formation of the 1 : 2 complex, providing the equilibrium constants differ sufficiently. Solutions containing only (TOAH)NO3 are obtained by equilibration of TOA with an empirically determined amount of HNO3 or by stripping excess acid with a pH ~ 8 buffer.
Analytical methods Nitric acid in the organic phase was titrated potentiometrically in an alcohol-water mixture, while N2 was bubbled through the solution. When C0/Y,TOA > 1 two neutralization points are found, at pH ~ 10"5 and at pH ~ 4"5, corresponding to the amine salt and the excess acid respectively. Water determinations were performed by Karl Fischer titrations, using dead stop endpoint indication.
Conductivity measurements Conductivity measurements were performed with standard equipment. The cell constant was verified at short intervals with 0"01 and 0"02 M KC1 solutions.
Viscosity and density measurements Water-calibrated Ostwald viscometers were used for the measurements. Flow times ranged between 18 sec and 4 hr. Densities were measured by weighing temperature equilibrated samples, drained from a 5 ml pipette or pycnometrically.
Infra-red spectra Infra-red spectra were obtained by previously described methods, c~ using variable stainless steel liquid cells with CaFz and NaC1 windows, which were checked for window contamination after each measurement. Spectra of HNOa and DNOs in the range 1 < Co/~TOA < 2 were recorded and the intensity of the CH stretching frequency in chloroform was studied as a function of the amine salt concentration. RESULTS AND DISCUSSION
The dipole moment o f ( T O A H ) N O a and its association behaviour Species w i t h large d i p o l e m o m e n t s associate w h e n dissolved in inert diluents. W h e n all o t h e r p a r a m e t e r s are k e p t constant, the degree o f association d e p e n d s on the dipole m o m e n t o f the solute a n d t h e dielectric c o n s t a n t o f the solvent. A large d i p o l e m o m e n t o f t h e solute and a small dielectric c o n s t a n t o f the solvent cause increasing association. T h e d i p o l e m o m e n t o f t h e solute ( T O A H ) N O a was f o u n d to be 1-21 <11) b u t this value a p p e a r s t o o low. T h e literature provides evidence t h a t d i p o l e m o m e n t s o f q u a t e r n a r y a m m o n i u m salts a n d salts o f t e r t i a r y amines lie between 7 a n d 20. (12-14) A n estimate o f the d i p o l e m o m e n t o f ( T O A H ) N O z f r o m the existing d a t a is based o n three considerations. First, large dipole m o m e n t s are r e p o r t e d for the picrates, in which t h e distance between the positive a n d t h e negative pole is considerable. Second, w h e n smaller anions are i n t r o d u c e d , the distance between t h e poles a n d consequently the d i p o l e m o m e n t decreases. T h i r d , the salts o f t e r t i a r y amines show the a d d i t i o n a l possibility o f h y d r o g e n b r i d g i n g between the a n i o n a n d the n i t r o g e n - b o n d e d h y d r o g e n , which will cause a f u r t h e r decrease o f the distance between the poles a n d o f the d i p o l e m o m e n t . F r o m t h o s e considerations the d i p o l e m o m e n t o f ( T O A H ) N O z can be (~) G. SCIBONA,Comitato Nazionale per le Ricerche Nucleari, CNC-43. (xs) j. A. GEDDm and C. A. KRAUS, Trans. Faraday Soc. 32, 585 (1936). (is) E. A. RICHARDSONand K. H. STERN,J. Amer. Chem. Soc. 82, 1296 (1960). (x4) W. R. GXLKERSONand K. K. SRrVAs"rAVA,J. Phys. Chem. 65, 272 (1961).
1088
J.M.P.J.
VERSTEGEN
estimated to lie between 7 and 12. Species with a dipole moment of that order of magnitude are not likely to exist as monomers in diluents such as benzene. Indeed it was found that salts of tertiary amines do not associate to the same extent as the quaternary compounds. (ls~ However, the tendency towards association is still pronounced. According to a classification introduced by KRAUSc16~, a tertiary amine salts with small anions, such as (TOAH)NO3, belong to his group B. A well investigated representative of this group is tri-n-butyl amine iodide and as a first approximation
8
7 6
d
. 4
././" ././"
~.
~..i--
.~'" ~
I
O0 Ioo
J 20 so
4~ 60
60 40
80 20
tO0 Mole 0 Mole
% TOA % Cell e
FIG. 3.--r/~5 in centipoise against mole per cent for TOA-benzene mixtures.
exp., ..... from eq. 2, .......... from eq. 3. its behaviour can be taken as a measure of that of (TOAH)NO3. Tri-n-butyl iodide shows considerable association at a concentration as low as 10-2 M. (17) Apparently association is a very c o m m o n phenomenon in those systems even at moderate concentrations, such as 0.1 M. A study of the viscosity and the conductivity was undertaken to clarify whether the association leads to colloidal behaviour or to formation of high molecular association products in true solution.
Viscosity The viscosity of T O A was found to be 8.084 centipoise at 25.0°C and the activation energy of viscous flow was Evjs = 5.3 kcal/mole (1) In Figs. 3 and 4 the viscosites r/z5 are given in centipoise at 25.0°C for TOA-benzene <1~ F. M. BATSON and C. A. KRAUS, J. Amer. Chem. Soc. 56, 2017 (1934).
Distribution of nitric acid--II
1089
and TOA-chloroform mixtures. Two empirical treatments, one the cube root relation of KENDALLand MONROE t18) 9~Jl =
X19~1} i -
(1 -- xl)~/2},
(2)
where x I is the molar per cent concentration of component 1, and the other (3)
logi0 4 = Xl log10 41 q- (1 -- Xx) log10 42
where 4 is the fluidity, c19) were applied to the data. In the TOA-CHCla system a 1:1 compound seems likely.
J J ,i .I
/
¢)
I
0 ~00
20 80
FIG.
/
I
I
40 60
6O 40
/ / Jg" /
//"
I
8O 20
~00 0
Mole % Mole %
TOA CHCI 3
4.--~/25 in centipoise against mole per cent for TOA-chloroform mixtures. - -
exp.)
.... from eq. 2, ........ from eq. 3. One may now turn to the amine phases, which have been equilibrated with nitric acid of various concentrations. Formations of high molecular species will cause the viscosity to increase, while colloidal particles exhibit very pronounced viscosity effects. In Fig. 5 the viscosity 725 is given in centipoise at 25.0°C as a function of Co/Y.TOA for various values of Y,TOA in benzene and chloroform. An increase in viscosity occurs at high amine concentrations in benzene as well as in chloroform, but the effect is more marked in benzene, which has the lower dielectric constant. The increase by a factor two is not of the same order of magnitude as that for lyophilic sols in organic diluents, where the viscosity of a 1 per cent solution may be 60 (rubber in benzene) to 1000 (nitrocellulose in acetone) times that of the diluent3 ~°) ilS) j. KENDALLand K. P. MONROE,J. Amer. Chem. Soc. 39, 1802 (1917). (19)S. GLASSrONE,Textbook of Physical Chemistry (2nd Ed), p. 500. Macmillan (1956). (20)E. HATSCHEK,The Viscosity of Liquids. Bell (1928).
1090
J.M.P.J.
VEP,STF_~EN
The viscosities of associated liquids show a temperature coefficient: if r/t and %' are the viscosity of the dispersion medium and of the sol at temperature t, then ~h'/% decreases with rising temperature. For (TOAH)NO3 solutions in benzene (C0/Y,TOA = 1) %'/~t decreases by 5 per cent when Y~TOA = 0.1M, by 15 per cent when ETOA = 0.25M and by 20 per cent when Y,TOA = 0.50M, measured in the 2.0
1.5
¢¢
/D,~ " £]-" ~ ~"O" " - ' ~ "-. . . .
0
-.12.-
// //// /
1.0
1 1 1
I
I 2
I 3
Co/3" TO A FIG. 5.--~z5 in centipoise against C0/Y,T O A for various values o f Y.TOA in benzene (solid lines and data points) and chloroform (dashed lines and open data points). • : Y,T O A = 0 . 5 0 M , • : ~ T O A = 0.25 M, • : Y~TOA = 0.10 M.
temperature range from 8 to 50°C. This decrease is an indication that association takes place in the amine nitrate phase. However, for sols the decrease is generally of the order of a factor two or three. (~°) The amine salt phases were never observed to set to gels. All the experiments point to a certain degree of association in the solutions of (TOAH)NO 3 in benzene and chloroform, but there is no evidence that colloidal species are formed. However, this conclusion is not necessarily valid for diluents with still lower dielectric constant, e.g. hydrocarbons. Here colloid formation might occur and certainly the viscosity effects at low concentrations are pronounced. (1)
Distribution of nitric acid--II
1091
An increase in viscosity by a factor 300 is observed when the concentration of (TOAH)NO a in benzene increases from 0 to 100 weight per cent. In Fig. 6 ~5 incentipoise is plotted against weight per cent for some quaternary ammonium salts, and (TOAH)NOa conforms to the pattern. The quaternary ammonium thiocyanates in benzene show somewhat higher viscosities, (~1) probably because of their larger dipole moments. For tetra-n-butyl ammonium picrate in n-butyl alcohol (2zl the viscosity curve remains below that of (TOAH)NO a, apparently on account of the more polar diluent. 5,10zI 2.10z ~02 50
®
I
2O I0 a (J
+e
5 2 I
05 I
o
51o
IO0
Weight % amine salt
FIG.
--I---© + ®
6.--~hs in centipoise against weight per cent amine salt. (TOAH)NO3 in benzene, tetra-n-butyl ammonium picrate in n-butylalcohol,c~ tetra-iso-amyl ammonium thiocyanate in benzene,'zx' tetra-butyl ammonium thiocyanate in benzene.(~x)
Conductivity
The specific conductance K in ohm -x cm -1 of the amine solutions after equilibration with aqueous HNOa is plotted against Co/ZTOA in Fig. 7. The measurements were performed at 25.0°C. There is a strong increase in conductance in the range C0/Y,TOA < 1 and occasionally at higher acidities. Interpolation at C0/T.TOA = 1 gives us the conductivity data for (TOAH)NOa solutions. The degree of dissociation of the amine nitrate can be estimated from conductivity and viscosity data, according to WALDEN'Srule (~a) 0: = At//60 (4) where A is the molar conductivity and ~7 the viscosity in centipoise. The values obtained are given in Table 2. The degree of dissociation 0: increases with the (=1)R. P. SEWARD,J. Amer. Chem. Sac. 73, 515 (1951). ~2) L. E. STROrqCand C. A. IrOtAUS,J. Amer. Chem, Sac. 72, 166 (1950). (~8)p. WALDEN,Elektrochemie Nichtwiissriger LSsungen Barth, Leipzig (1924).
1092
J . M . P . J . VERSTEGEN
16a ~.,.o"
o
#
10"
E o E
165
1(56
0
I
I
I
I
2
3
Co/Y.roA FIG. 7.--K(~ -z cm -x) against C0/~TOA. Symbols as in Fig. 5. concentration. That, as well as the strong increase in conductance, is characteristic o f salts in low dielectric media.t24, ~s) I t has been suggested t h a t micelles c o u l d be f o r m e d in the amine salt solutions,t26# 7~ a n d it should be k e p t in m i n d t h a t the big increase in c o n d u c t a n c e c o u l d be ascribed to micelle f o r m a t i o n . A t the critical micelle c o n c e n t r a t i o n a steep increase TABLE 2 . - - D E G R E E OF DISSOCIATION g OF
Conc. (TOAH)NO3 (mole/l) 0.50 0'25 0.50
(TOAH)NOs
Diluent Benzene Benzene Chloroform
5 × 10 -a 2 × 10-3 5 × 10-~
tu~ C. A. KRAUSand R. M. Fuoss, J. Amer. Chem. Soc. 15, 21 (1933). ~=5~p. L. MERCmR and C. A. KRAOS,Proc. Natl. Acad. Sci. U.S. 42, 487 (1956). t=e~K. A. ALLEN,J. Phys. Chem. 60, 239 (1956). t=:~K. A. ALLEN,J. Phys. Chem. 60, 943 (1956).
Distribution of nitric acid--II
1093
in conductivity is often observed, as the resistance to motion per unit electrical charge becomes smaller when multiply charged particles are formed. ~'8~ However, the micelle concept in that form is typical of aqueous solutions or at least of water-rich mixtures. Such micelles, with the non-polar carbon chains directed towards the centre and the polar groups, forming the outer sphere and interacting with the diluent, are of course highly improbable under the present conditions. A second objection to the explanation of the conductivity behaviour by means of the classical micelle concept is the wide range of concentrations in which the increase is observed, whereas one would expect it at the critical micelle concentration only. A reversed micelle with the polar groups directed towards the centre and the carbon chains forming the outer sphere has also been suggested, m) Such aggregates would exhibit conductivity behaviour different from the classical micelles and it is not clear how their formation could explain the steep increase in the conductivity. Co-extraction of water In Fig. 8 are given the amounts of water co-extracted when amine solutions are equilibrated with aqueous HNO,. From the experimental results the amounts of water extracted by the amine-diluent mixtures have been subtracted, so that the curves have their origin at zero. The amount of water in the saturated amine-benzene and amine-chloroform mixtures is not significantly different from that in the pure diluents, so formation of a compound TOA-H20 is not likely. The correlation between co-extraction of water and the viscosity agrees with the assumption that co-extracted water and association both stabilize the ion-pair in the low dielectric diluents. The results given here follow the pattern outlined earlier, {1) according to which the co-extraction of water decreases in the diluent with the high dielectric constant although the solubility of the water increases. At higher acidities the water co-extraction increases considerably, as has also been observed in TBP systemsJ29~ A partial explanation lies in the extraction of a hydrated acid species in the diluentJ 3°~ Equilibrations Solubility of HNOa in the pure diluents. HARDY et al. ~°~ showed that the nitric acid solubility in organic diluents can be understood on the basis of a distribution of HNO 3 between water and the organic phase and subsequent dimerization of the acid. Introducing Co, the organic acid molarity and Cua, the undissociated aqueous nitric acid concentration at equilibrium, an equation C O = DCua q- 2Km~o3 D~Cua~ (5) was shown to give an explanation of the processes taking place. Plotting Co/Cua against Cua* the value of the distribution coefficient D is found and that of the dimerization constant KH~ros can be calculated. * Values of C.a were taken from K~WETZ data, reported by Hr~FORD and McK.AY.~sl~ Their original graph, which was large enough to read off fairly accurately was kindly sent to us by Mr. H. A. C. McKAY. i~s~ C. A. KRAUS, J. Amer. Chem. Soc. 60, 129 (1956). izg) K. ALCOCK, S. S. GRIMLEY,T. V. HEALY,J. KENNEDYand H. A. C. McKAY, Trans. Faraday $oc. 52, 39 (1956).
tao)C. J. HARDY,B. F. GREENFIELDand D. SCAgGILL,J. Chem. Soc. 90 (1961). tal) E. HESFORDand H. A. C. MCKAY, Trans. Faraday Soc. 54, 573 (1958).
1094
VE~GEN
J.M.P.J. 0.2~
0.2(
0.15
(--)
o.lo
7//
,
/
/
005
x,,&__ . & /
Y
o
~
I
;
3
Co/ZTOA
FIG. 8.ICO-extraction of H=O in mole/], against Co/~.TOA. Symbols as in Fig. 5.
Values of KHNo8 and D in kerosene and toluene have been reported and some measurements in other diluents have been given. (3°) The work of BERTOCCIand ROLANDI(6) deals with the HNOa solubility in xylene. HARDY'S equation can be applied to their data at the three higher acidities, when the highest value of the parallel determinations at 8.7 M is chosen. The values of D and KH~o= obtained in this way fit nicely into the picture given in Table 3. In column 2 values of the dielectric constant e have been tabulated. TABLE 3.~DISTRIBUTION COEFFICIENT D AND DIMERIZATION CONSTANT KHNOa
Solvent Ker osene Benzene Toluene Xylene o- Dichloro benzene Chloroform
~ N1.9 2.27 2.38 N2-4 9'93 4"8
Distr. coeff. D 6 1 1.2 1"4 1"4 4
x x x x x x
10 -5 10 -8 10 -a 10 -a 10 -a 10 -s
dim. const. K a s % (mole/l) -x 4000 t240 590 300 ~ 100 50
source Ref. This Ref. Ref. This This
30 work 30 6 work work
K=' (mole/l) -z -0-22 0.21 -0'09 if03
Distribution of nitric acid--II
1095
As the isomeric composition of the xylene mixture in Ref. 6 was not given, the value of 2.4 for the dielectric constant is arbitrary. HARDY'S treatment lacks thermodynamic validity, but the results obtained in that way lead to a pattern which can easily be understood from general chemical considerations. In solvents of low dielectric constant the distribution coefficient decreases and the dimerization increases. Chloroform is exceptional in possessing additional possibilities of stabilizing monomeric acid by means of hydrogen bonding. Empirical relations in the high acid range. Concentrations in mole/l are denoted by C, molar activity coefficients by f, the thermodynamic equilibrium constant (based on aqueous and organic activities) by K, and the equilibrium constant based on aqueous activities and organic concentrations by K'. The aqueous HNOa activity 0.12
o.~o
~
]
I
h
O08 o E
0.06
0.04
]3Z
0 02
i
0
I
2
3 Cf,
4
5
6
mole/L
FIG.9.--CcagainstCf for 0"1M TOA solutionsin variousdiluentsat 25°C. - - O - - benzene, - - A - - toluene, - - I I - - o-dichloro benzene, - - O - - chloroform. is denoted by aH~o, and calculated from tabulated values o f T + . (3~) The subscripts a, s and c refer to the free amine, the ion-pair and the 1 : 2 complex respectively. Other terminology is the same as defined before, (1) viz. C i and Cf refer to initial and final aqueous HNO3 concentration and C o = Cs + 2Ce, while ETOA = C a + Cs + Ce. Some typical C~ against Cf curves are shown in Fig. 9. In benzene an empirical relation Ce = 0.19 CfETOA (6) gives a good approximation for 2M < Cf < 6M. The material balance Cf---- C i -- Ce -- ETOA
(7)
combined with Equation (6) leads to
Cf = Ci/[1 + 0-19 ETOA -- function (Y,TOA)] (8) The small deviations from Equation (6) disappear in Equation (8) as 0.19ETOA ~. 1. When C r is plotted against Ci straight lines are found, but small volume changes (s2)
LANOOLr-BrRNSTEIN,Physikalisch-ChemischeTabellen, Eg. IIb, page 1119, Eg. Illc, page 2145.
1096
J.M.P.J.
VERSTEGEN
cause the slope to deviate somewhat from 1/(1 -+- 0.19ZTOA). The empirical relations are given in Table 4. TABLE 4.--EMPIRICAL RELATIONSFOR H N O a EXTRACTION BY (TOAH)NOa-BENZENE PHASES Y,T O A (mole/l.)
Relation
0-465 0-249 0-100
C t = 0 ' 9 2 9 C i -- 0'43
0.050
Ct = 0.975C~ -- 0.01
Cr = 0"967C1 -- 0'22 C r = 0"972C l -- 0"07
Formation of a 1:2 amine-acid complex at high acidities. The quantitative description of the extraction of H N O 3 in the high acid region is based partly on a reconsideration of the results in the ion-pair range (C0/Y~TOA < 1). Under certain conditions the formation of (TOAH)NO 3 could be based on the mass action law. tl) In benzene and chloroform these conditions are fulfilled if the initial TOA concentration _< 0.1 M and the final (TOAH)NO 3 concentration < ~--~0.02 M. In benzene these conditions are not fulfilled at initial TOA concentrations _>0.25 M, even when the amount of organic amine salt remains very small. Here the formation of (TOAH)NO3 takes place according to a mechanism which differs from the usual mass-action expression. In chloroform a similar effect might be observed, but the extractive power for acids of concentrated amine solutions in that diluent is too large to allow accurate measurements. It was reported earlier tl) that in modified dodecane the TOA concentration had to be kept as low as 0.05 M to observe mass-action behaviour. Analogous observations in TOA -- H2SO 4 systems showed that the formation of amine sulphate in benzene occurred on a mass action basis as long as the TOA concentration was <0.25 M. At higher amine concentration (0.50 M) mass action behaviour was not observed.t 2n) The general trend behind these observations is that when the initial amine concentration increases, a critical value is reached, beyond which the formation of amine salt can not be explained by a mass action mechanism. Or, the behaviour of the amine salt depends on the amount of free amine present and that suggests some sort of interaction between the two. If the highest amine concentration with a mass action range is defined as the critical amine concentration (CAC) a table can be made to show the influence of various parameters (Table 5). TABLE 5.--INFLUENCE OF PARAMETERSON C A C C A C (mole/l.) 0-25 <0.10 <0.05 0.10 0.05
amine
a q u e o u s acid
diluent
Tri-n-octyl Tri-n-octyl Di-n-decyl Tri-n-octyl Tri-n-octyl
H2SO4 H~SO~ H2SO, HNOs HNO3
Benzene Kerosene Benzene Benzene M o d . dodecane
source Ref. Ref. Ref. This Ref.
26 33 27 work I
Distribution of nitric acid--II
1097
In comparison secondary amine salts have a large polar group, whereas kerosene and modified dodecane have lower dielectric constants. Probably a lower CAC and a greater tendency towards third phase formation correspond to a greater difference in polar character between solvent and solute. Additional evidence for TOA-(TOAH)NOa interaction is obtained as follows. It has been shown before,t2e,27,a3,a4~ that empirical formulae such as Kemp. = aH~rOa [TOA] ~
(9)
can often describe the results. A relation of the same form with n = ] can be applied to the (TOAH)NO3-benzene system in the concentration ranges where the law of mass action is not obeyed (that means in the whole range 0 < C0/Y.TOA < 1 when Z T O A = 0.25 and 0.50 M and for (TOAH)NOz concentrations > ~ 0.02 M in the cases that Y,TOA = 0.05 and 0-10 M). When an~o, is plotted against TOA on a logarithmic scale, four parallel straight lines are found, and from the slope n can be derived. The lines for different ETOA values are not superimposed. This is analogous to results in sulphate systems as reported by ALLEN (26) who showed that interaction of free amine with the amine salt or its high molecular association products may provide an explanation of the effect. The side reactions discussed above and in the section on viscosity cause deviations from mass-action behaviour, such that in a certain range of concentrations Equation (9) has to be preferred. Another approach to the problem is to express the deviations in terms of the activity coefficients. This is a purely formal treatment, which does not lead to a better insight into the processes taking place, so long as activity coefficient behaviour is not fully understood in the system under consideration. However, in the present case useful information was obtained and the treatment provides evidence for the existence of a 1 : 2 amine-nitric acid complex. The gross equilibrium constant K1 of TOA -}- H + + NO~- ~ (TOAH) NO 3
(10)
is obtained by interpolation from earlier work. tl~
K1 =f~(TOAH)NO3/fa TOA a~ro 3=fJfaK; -= 0.48 × 106 (mole/l) -2
(11)
and
fs/f. = 0.48 × 106/K;
(lla)
When .fJfa is plotted against C0/YTOA the curves in Fig. 10 are obtained. For Y~TOA = 0.05 and 0.10 M a region exists where the processes can be explained on basis of the mass action law. Here the r a t i o f J f a is arbitrarily set equal to one. That does not necessarily mean that in this region (TOAH)NO3 exhibits monomeric behaviour. The gross quantity K1 might contain equilibrium constants arising from possible side reactions. The deviations offa from one in the binary mixture TOA-benzene are discussed in the section on reagents. They are apparently small. An explanation of the shape of the f~/fa curves is based on two considerations. First, mutual association of amine salt molecules causes the thermodynamic concentration Co/n to fall considerably below the stoichiometric one, C O(n is the average number of (TOAH)NOa c38~j. M. P. J. VERSa~GENand J. A. A. KETELAAR,J. Phys. Chem. 66, 126 (1962). t34~ j. M. P. J. VERSXEGENand J. A. A. KETELAAR, Trans. Faraday Soc. 57, 1527 (1961). 13
1098
J . M . P . J . VERSTEGEN
molecules per association product). Second, association of TOA with high molecular amine salt species decreases the amount of free amine present. Both effects can be accounted for by activity coefficients < 1. The fs/f~ curves show three regions of interest. In the first, at low values of C0/~VTOA,fJf~ decreases. Here the amount of unconverted amine is dominant and the formation of amine salt association products has not yet proceeded very far, so that most of the TOA remains in solution. The shape of the curve is mainly determined by the mutual association of amine salt molecules and the corresponding
0.2
0.4
0.6
0.8
1.0
Co/~TOA
Fro. lO.--fJf=
against Co/~TOA in the range 0 < C o / ~ T O A < 1 at 25°C. . . . . . . . . . . ~ T O A = 0.47 M , - ~ T O A ---- 0.25 M ,
. . . . . . . . . ~ETOA= 0.10 M, - - - ~TOA = 0"05 M. decrease offs. At high values of Co/Y~TOAfJfa increases. Here most of the TOA has been converted to the nitrate and appreciable quantities of associated amine salt species are present. The shape of the curve is mainly determined by the interaction of TOA with those species and by the corresponding decrease infa. In the region of intermediate acidities the two effects compensate one another and the ratio fdfa may remain fairly well constant. The same treatment can be applied to the data in the region 1 < Co/~TOA < 2. It is assumed that a 1:2 complex is formed between TOA and HNO3, according to the equilibrium (TOAH)NO8 + H + + NOa- ~ (TOAH)NOa'HNOa,
(12)
The gross equilibrium constant is Ks -----fe(TOAH)NOa'HNOa/j~(TOAH(NOa) arr~ros =
fe/fBK~'
(13)
In Fig. 11 K2' is plotted against Co/ZTOA. Except for a constant the curves represent plots offdfc against Co/ZTOA in the range I < C0/ZTOA < 2 and for different values of ZTOA. There are considerable deviations from ideality. The increase of the fs/fc curves when Co/ZTOA is slightly larger than 1 is probably due to association of the 1 : 2 complex with the amine salt. Another range where fs/fe increases is found at acidities corresponding to
Distribution of nitric acid--II
1099
1.6 < Co/ZTOA < 2. In this range the viscosity decreases and the co-extraction of water increases. A possible explanation is that, on account of the increasingly polar character of the organic phase, the stabilizing influence of association is replaced by that of hydration. The species de-associate to form hydrated lower molecular products. Special interest attaches to the points of intersection of the fs/fe and fs/fa curves with the ordinate Co/ZTOA = 1, obtained by extrapolation. Purely formally they represent the ratio of the molar activity coefficients fs for different (TOAH)NOz concentrations. From Fig. 10 even the absolute values o f f i
t \
o.6~ o~1~\\
0.7
"\
'\
kl
0.4 •
0,2
'"
'~
-" --~"~'~..--~
. . . . . . .
~
.-~'~/
0.1
0
I I2
I 1.6
L 1.4
| 1.8
2 0
CoI'~TOA FIG. II.--K~' against C0/~TOA in the range 1 < C0/ETOA < 2 at 25 ° C. Symbols as in Fig. 10.
can be derived, providing the assumption is valid that fJfa--~ 1 when Co ~ 0. Although the values of fs obtained in that way are probably not meaningful, the extrapolations to C0/ZTOA = 1 from the low acid side (0 < Co/Y.TOA < 1) and from the high acid side (1 < C0/ZTOA < 2), yield the same ratio of fs values for different amine salt concentrations. This provides some evidence that a 1:2 complex is indeed formed. Further reasons for believing that when 1 < C0/~TOA < 2 formation of a 1:2 complex is predominant will be given now. First, when Cf ~-, 6 M (where Co/Y.TOAis slightly larger than 2) new phenomena occur, indicated by an increasing acid extraction, which was not reproducible. Extraction of HNO 3 by the diluent does not account for the whole effect. This observation provides another example of anomalous extraction behaviour such as was first reported by McDowH~L et al. (~) Energy differences between possible 1:3 and 1:4 complexes are probably small, and after phase settling non-equilibrium species may still exist, because of slow diffusion to the interface. A second reason to believe in tss) W. J. McDOWELL and K. A. ALLE~, J. Phys. Chem. 65, 1358 (1961).
1100
.l.M.P.J. V~RST~GEN
the formation of the 1:2 complex is the rough proportionality between ETOA and excess acid concentration. Consequently about the same ratio Ce/Cs is found at the same value of Cf, when the amine concentration varies from 0.05 to 0.5 M. Third, at least in the region I < Co/ETOA < 1.6 the Ks' values approach constancy, when the amine nitrate concentration decreases (Fig. 11). A fourth reason is the influence of the diluent on the formation of a 1 : 2 complex, as shown in Fig. 9. In Table 3, column 6, values of K s' are compared at Ce/Cs = 1 and for 0.10 M TOA solutions. It has been suggested tr's) that the distribution of HNO 3 between aqueous solutions and amine nitrate phases can be explained as simple dissolution of the acid in a more polar organic solution. The trend given in Table 3, column 6, is at variance with this conception. The amount of HNO3 extracted into the organic amine nitrate phases, increases with decreasing dielectric constant of the diluent, while in the pure diluents the reverse is true. If the role of the amine salt is only to lend polar properties to the diluents, one would expect the original sequence to be maintained. In view of the data presented above, a better explanation of the distribution at high acidities is based on the stabilization of the amine nitrate salt, which is formed at low acidities. (~ The ion-pair needs to surround itself with stabilizing molecules, especially in the low dielectric diluents. Nitric acid may act as a stabilizing molecule and it will be extracted to a greater extent when stabilization by the diluent is unimportant. Chloroform shows stabilizing properties, much greater than would be expected from its dielectric constant. The amine nitrate salt is probably solvated in that diluent and the interaction takes place by hydrogen bonding. From our measurements at 4 and 25°C in benzene systems it was found that the distribution at high acidities was insensitive to temperature. The same was found at our Institute for the distribution of HNO 3 between aqueous solutions and a solution of tri-lauryl amine in modified dodecane. (36~ The main impression is that the second HNO 3 molecule is bonded to the amine nitrate by means of hydrogen bonding. The 5 kcal/mole gained in that way, could be balanced by side effects, such as de-association, partial dehydration etc.
Infra-red spectra It has been postulated (1) that higher complexes are formed in the amine nitrate system even in cases where as much as 40 per cent of the amine remains available for salt formation. From the present paper the conclusion may be drawn that the equilibrium constants for the first two complexes differ by a factor of 10e. Therefore, no appreciable amounts of free amine exist in equilibrium with 1:2 complex. The prior erroneous assumption, which does not affect the general trend reported earlier, <1) was primarily caused by interpretation of the C i against Cf curves. Typical plots in benzene showed a rapid increase of Cf but in the present paper it is shown that this is due to the shape of thefJJk curves (Fig. 10). Infra-red spectra provide evidence that after reaching a value Co/]~TOA = 1, the salt concentration remains constant. When spectra are recorded for (TOAH)NO 8 in CHCla, compensated with TOA in CHCI 3, the chloroform CH stretching frequency is found at 2998 cm -~ and a range of transmission >100 per cent is observed at 2780 cm -1. The explanation is as follows. In the reference cell, some CHCI a molecules (36~
R. BAc, Institutt for Atomenergi. Personal communication.
Distribution of nitric acid--II
1101
are hydrogen-bonded to TOA. This causes the CH frequency to split into two branches, a low frequency component at 2780 cm -1, due to interaction with the amine, and a component at the usual frequency of ~ 3 0 1 9 cm-1. taT~ In the sample cell HNOa replaces CHC13 in its interaction with TOA, but now solvation of the amine nitrate takes place. Here the CH frequency is split into a branch at ,-~ 3019 cm -~, compensated in the reference beam, and one at lower frequency, due to interaction between the chloroform hydrogen and the amine nitrate oxygen. The A
B I
T
Trr
I 28
i 25
I 22
I 28
I 25
I 22x102 cm-I
FIG. 12.--The shape of the N÷-H . . . . (ONO2)- band at different organic acidities
C0/]~TOA in chloroform (A) and benzene (B) as diluent. I: Co/5'.TOA< 1, II: Co/~TOA ~ 1, III: Co/ETOA :> 1. latter emerges from the spectrum as a peak at 2998 cm -1, with an intensity proportional to the (TOAH)NO 3 concentration. It was found that after reaching a value C0/ETOA = 1 the intensity of the peak, and consequently the (TOAH)NOz concentration, remained constant. Interaction between the amine nitrate group and the second H N O 8 molecule may lead to a weakening of the central hydrogen bond in the amine salt ion-pair. Indeed it was observed that the N + -- H . . . (ONOz)- band at 2430 cm -x in the spectra of (TOAH)NO3 in benzene and chloroform ~1~ shifts towards higher frequencies when C0/]~TOA > 1. In CHC13 this leads to a splitting with a new frequency at 2560 cm -1, which is found as a weak shoulder when Co/ETOA < 1 (Fig. 12). Nitric acid in the range 1 < Co/ETOA < 2 and CHC13 in the range 0 < C0/ETOA < 1 form hydrogen bonds with the oxygens of the amine salt anion. There is a competition between excess H N O 3 and CHCI~, consistent with the low equilibrium constant found for the 1:2 complex in that diluent. The observed infra-red frequencies of excess HNOz (DNOz) are given in Table 6. The spectra are difficult to interpret, because in the region of interest the frequencies of the ion-pair nitrate shift under influence of the excess acid. Assignments were made using References 38-40. ~37~G. HERZBERG, Infra-red and Raman Spectra of Polyatomic Molecules p. 316. Van Nostrand, New York (1945).
1102
J. M. P. J. VEP-,STF,GEN
TABLE6.--INFRA-REDFREQUENCIESOF EXCESSHNOs AND DNOs IN (TOAH)NOs Excess HNO 8 in ( TOAH)NOa--CeH e (cm- 1)
Exc,~s HNO s in (TOAH)NOa-.CHCI 8
Excess DNO s in (TOAH)NOa-CeH e
(eat-1 )
HNO s yap.
(era -:t)
HNO$ liq.
D N O s DNO a Assignyap. liq. mcnt (ss-4o)
(cna-1)
~3100 1645
~3200 1640
2230 1609
3560 1710
3400 1675
2627 1685
2470 1645
1295
1310
1298
1320
1300
1313
1300
948
938
938
886
925
888
915
HO stretch NO asymm. stretch I~O symm. stretch NOH stretch
The frequencies suggest the structure f3
(CsH17)s N--H. .... O - - N ~
. . . . . . HONO~
(electrical charges omitted).
Acknowledgement--This paper is published by permission of the Managing Director of Institutt for Atomenergi and the work forms part of the programme of a Dutch-Norwegian Joint Project. Mrs. C. M. KOFP~RtrD,Mr. I. A. HtrgoERE and Miss W. A. TAP assisted in various facets of the work. css~C. K. INGOLDand D. J. MmLEN,J. Chem. Soc. 2612 (1950). ts,J H. Cot]n, C. K. INGOLDand H. G. POOLE,3". Chem. Soc. 4272 (1952). t40~R. A. MARCUSand J. M. FRESCO,./. Chem. Phys. 27, 564 (1957).
DISTRIBUTION OF NITRIC ACID BETWEEN WATER A N D O R G A N I C SOLVENTS, C O N T A I N I N G TRI-N-OCTYL AMINE--II J. M. P. J. VERSTEGEN Institutt for Atomenergi, Kjeller, Norway (Received 5 April 1963; in revised form 4 July 1963) Abstract--Previous measurements on the distribution of HNOs between water and organic solvents containing TOA are extended. Viscosity and conductivity of the equilibrated amine phases can be explained by ion-pair behaviour in solvents of low dielectric constant. Formation of micelles or colloids is not likely. Hydration occurs and solvation was found in chloroform solutions. It is concluded that a compound (TOAH)NOfHNO3 is formed by hydrogen bonding and that the equilibrium constant depends on the solvent, but not significantly on the temperature. IT was s h o w n earlier ~1) that the d i s t r i b u t i o n of H N O 3 between water a n d organic solvents c o n t a i n i n g t r i - n - o c t y l amine (TOA) at low acidities can be described in terms of i o n - p a i r ( a m i n e - n i t r a t e ) f o r m a t i o n in the organic phase. After stoichiometric n e u t r a l i z a t i o n of the amine, further extraction of acid by the a m i n e occurs3 ~-4) SHEVCHENKO et al. <5) reached the c o n c l u s i o n that a n a m i n e nitrate complex with the c o m p o s i t i o n 1:2 exists. BERTOCCI a n d ROLANDIt6) concluded t h a t no stoichiometric c o m p o u n d between T O A a n d H N O 3 is formed. T~LAT-ERBE~c7) assumed that a series of complexes is b u i l t u p by stepwise equilibria. D u r i n g p r e p a r a t i o n of this paper BARONCELLI et al. ~8) concluded t h a t the d i s t r i b u t i o n of H N O a in the range of high acidities could be ascribed to dissolution of the acid in the amine nitrate, witho u t further c o m p o u n d formation. EXPERIMENTAL Reagents TOA was vacuum distilled and titrated in acetic acid. `9' The molecular weight found was 352"7. Organic TOA phases are in general colourless after shaking with HNOs, but at high acidities a brown or pink solution was often observed. The measured refractive index was n~5 = 1.4494, while the undistiUed sample gave n~ = 1"4491. The density of TOA at different temperatures is given in Table 1. Densities of TOA-benzene and TOA--chloroform mixtures and refractive indexes of the first are shown in Fig. 1 and 2. Fig. 1 shows agreement with data in similar systems,tx°~ Deviations of the densities from ideality are small. i1) j. M. P. J. VERSTEGEN,Trans. Faraday Soc. 58, 1878 (1962). t2) M. DE TRENTINIANand A. CHESNE, Compte Rendu du Colloque de Extraction par Solvent (Madrid, 1959). la) U . BERTOCCI,AERE-R 2933 (1959). ,4~ D. J. CARSWELLand J. J. LAWRANC~,J. Inorg. Nucl. Chem. 11, 69 (1959). ts) V. B. SREVCHENKO,V. S. SPAvllDT,E. A. NENAROKOMOWand K. A. PETROW,Zh. Neorg. Khim. 5, 1852 (1960). ca) U. BERTOCCIand G. ROLANDI,3". Inorg. Nucl. Chem. 23, 323 (1961). ~7, M. Tf~J~AT-ERBEN,Proceedings of the 7th L C.C.C. Stockholm/Uppsala (1962). ,s, F. BARONCF.LLI,G. SCmONAand M. ZIFFERERO,J. Inorg. NucL Chem. 24, 403 (1962). ~9~C. D. WAGNER,R. H. BROWNand E. D. PETERS,J. Amer. Chem. Soc. 69, 2609 (1947). ~t0, K. v. IPENaURG,Thesis, Delft (1962). 1085
1086
J.M.P.J.
VERSTEGEN
TABLE ] ,--DENSITY OF TOA AT DIFFERENTTEMPERATURES Temp. (°C) Density (g/ml)
25.1
35"4
49.9
61"2
75"8
0.80585
0.79950
0.78980
0.78370
0.77385
0.870 i 0.860
1.52
0.85(;
1.51
0.840
1,50
-. 0.830 E E
"-..'--...
1.49
,m 0.820
1.48
~e 0.810 r~ 0.800
1.47 1.46
0.790
1.45 0 100
20 ' 80
4 0' 60
8'0
6 'o 40
20
I00 0
Jo
Vol % TOA Vol % CB Hs
FIG. 1.--Density in g/ml (solid line, left hand scale) and n,Vo (dashed line, right hand
scale) plotted against vol. per cent of the components for the system TOA-benzene. Additivity lines included.
1.5o 140 "-.
130
~'
Jzo
u)
1.10
123 1.00
0.90 0.80 IO0
!
I
/
25 75
5O 50
75 Z5
JO0 VOI % TOA 0 Vol % CHCI 3
FfG. 2.--Density in g/ml plotted against vol. per cent of the components for the system TOA-chloroform. Additivity line excluded.
Distribution of nitric acid--II
1087
Fatuilibrations Equilibrations were performed at 25"0°C but some measurements were also made at 4°C, to study heat effects. Organic phase HNOs concentration is that of Co and total amine concentration by Y.TOA. The range C0/Y~TOA < 1 defines ion-pair formation, while 1 < C0/~TOA < 2 is the region of formation of the 1 : 2 complex, providing the equilibrium constants differ sufficiently. Solutions containing only (TOAH)NO3 are obtained by equilibration of TOA with an empirically determined amount of HNO3 or by stripping excess acid with a pH ~ 8 buffer.
Analytical methods Nitric acid in the organic phase was titrated potentiometrically in an alcohol-water mixture, while N2 was bubbled through the solution. When C0/Y,TOA > 1 two neutralization points are found, at pH ~ 10"5 and at pH ~ 4"5, corresponding to the amine salt and the excess acid respectively. Water determinations were performed by Karl Fischer titrations, using dead stop endpoint indication.
Conductivity measurements Conductivity measurements were performed with standard equipment. The cell constant was verified at short intervals with 0"01 and 0"02 M KC1 solutions.
Viscosity and density measurements Water-calibrated Ostwald viscometers were used for the measurements. Flow times ranged between 18 sec and 4 hr. Densities were measured by weighing temperature equilibrated samples, drained from a 5 ml pipette or pycnometrically.
Infra-red spectra Infra-red spectra were obtained by previously described methods, c~ using variable stainless steel liquid cells with CaFz and NaC1 windows, which were checked for window contamination after each measurement. Spectra of HNOa and DNOs in the range 1 < Co/~TOA < 2 were recorded and the intensity of the CH stretching frequency in chloroform was studied as a function of the amine salt concentration. RESULTS AND DISCUSSION
The dipole moment o f ( T O A H ) N O a and its association behaviour Species w i t h large d i p o l e m o m e n t s associate w h e n dissolved in inert diluents. W h e n all o t h e r p a r a m e t e r s are k e p t constant, the degree o f association d e p e n d s on the dipole m o m e n t o f the solute a n d t h e dielectric c o n s t a n t o f the solvent. A large d i p o l e m o m e n t o f t h e solute and a small dielectric c o n s t a n t o f the solvent cause increasing association. T h e d i p o l e m o m e n t o f t h e solute ( T O A H ) N O a was f o u n d to be 1-21 <11) b u t this value a p p e a r s t o o low. T h e literature provides evidence t h a t d i p o l e m o m e n t s o f q u a t e r n a r y a m m o n i u m salts a n d salts o f t e r t i a r y amines lie between 7 a n d 20. (12-14) A n estimate o f the d i p o l e m o m e n t o f ( T O A H ) N O z f r o m the existing d a t a is based o n three considerations. First, large dipole m o m e n t s are r e p o r t e d for the picrates, in which t h e distance between the positive a n d t h e negative pole is considerable. Second, w h e n smaller anions are i n t r o d u c e d , the distance between t h e poles a n d consequently the d i p o l e m o m e n t decreases. T h i r d , the salts o f t e r t i a r y amines show the a d d i t i o n a l possibility o f h y d r o g e n b r i d g i n g between the a n i o n a n d the n i t r o g e n - b o n d e d h y d r o g e n , which will cause a f u r t h e r decrease o f the distance between the poles a n d o f the d i p o l e m o m e n t . F r o m t h o s e considerations the d i p o l e m o m e n t o f ( T O A H ) N O z can be (~) G. SCIBONA,Comitato Nazionale per le Ricerche Nucleari, CNC-43. (xs) j. A. GEDDm and C. A. KRAUS, Trans. Faraday Soc. 32, 585 (1936). (is) E. A. RICHARDSONand K. H. STERN,J. Amer. Chem. Soc. 82, 1296 (1960). (x4) W. R. GXLKERSONand K. K. SRrVAs"rAVA,J. Phys. Chem. 65, 272 (1961).
1088
J.M.P.J.
VERSTEGEN
estimated to lie between 7 and 12. Species with a dipole moment of that order of magnitude are not likely to exist as monomers in diluents such as benzene. Indeed it was found that salts of tertiary amines do not associate to the same extent as the quaternary compounds. (ls~ However, the tendency towards association is still pronounced. According to a classification introduced by KRAUSc16~, a tertiary amine salts with small anions, such as (TOAH)NO3, belong to his group B. A well investigated representative of this group is tri-n-butyl amine iodide and as a first approximation
8
7 6
d
. 4
././" ././"
~.
~..i--
.~'" ~
I
O0 Ioo
J 20 so
4~ 60
60 40
80 20
tO0 Mole 0 Mole
% TOA % Cell e
FIG. 3.--r/~5 in centipoise against mole per cent for TOA-benzene mixtures.
exp., ..... from eq. 2, .......... from eq. 3. its behaviour can be taken as a measure of that of (TOAH)NO3. Tri-n-butyl iodide shows considerable association at a concentration as low as 10-2 M. (17) Apparently association is a very c o m m o n phenomenon in those systems even at moderate concentrations, such as 0.1 M. A study of the viscosity and the conductivity was undertaken to clarify whether the association leads to colloidal behaviour or to formation of high molecular association products in true solution.
Viscosity The viscosity of T O A was found to be 8.084 centipoise at 25.0°C and the activation energy of viscous flow was Evjs = 5.3 kcal/mole (1) In Figs. 3 and 4 the viscosites r/z5 are given in centipoise at 25.0°C for TOA-benzene <1~ F. M. BATSON and C. A. KRAUS, J. Amer. Chem. Soc. 56, 2017 (1934).
Distribution of nitric acid--II
1089
and TOA-chloroform mixtures. Two empirical treatments, one the cube root relation of KENDALLand MONROE t18) 9~Jl =
X19~1} i -
(1 -- xl)~/2},
(2)
where x I is the molar per cent concentration of component 1, and the other (3)
logi0 4 = Xl log10 41 q- (1 -- Xx) log10 42
where 4 is the fluidity, c19) were applied to the data. In the TOA-CHCla system a 1:1 compound seems likely.
J J ,i .I
/
¢)
I
0 ~00
20 80
FIG.
/
I
I
40 60
6O 40
/ / Jg" /
//"
I
8O 20
~00 0
Mole % Mole %
TOA CHCI 3
4.--~/25 in centipoise against mole per cent for TOA-chloroform mixtures. - -
exp.)
.... from eq. 2, ........ from eq. 3. One may now turn to the amine phases, which have been equilibrated with nitric acid of various concentrations. Formations of high molecular species will cause the viscosity to increase, while colloidal particles exhibit very pronounced viscosity effects. In Fig. 5 the viscosity 725 is given in centipoise at 25.0°C as a function of Co/Y.TOA for various values of Y,TOA in benzene and chloroform. An increase in viscosity occurs at high amine concentrations in benzene as well as in chloroform, but the effect is more marked in benzene, which has the lower dielectric constant. The increase by a factor two is not of the same order of magnitude as that for lyophilic sols in organic diluents, where the viscosity of a 1 per cent solution may be 60 (rubber in benzene) to 1000 (nitrocellulose in acetone) times that of the diluent3 ~°) ilS) j. KENDALLand K. P. MONROE,J. Amer. Chem. Soc. 39, 1802 (1917). (19)S. GLASSrONE,Textbook of Physical Chemistry (2nd Ed), p. 500. Macmillan (1956). (20)E. HATSCHEK,The Viscosity of Liquids. Bell (1928).
1090
J.M.P.J.
VEP,STF_~EN
The viscosities of associated liquids show a temperature coefficient: if r/t and %' are the viscosity of the dispersion medium and of the sol at temperature t, then ~h'/% decreases with rising temperature. For (TOAH)NO3 solutions in benzene (C0/Y,TOA = 1) %'/~t decreases by 5 per cent when Y~TOA = 0.1M, by 15 per cent when ETOA = 0.25M and by 20 per cent when Y,TOA = 0.50M, measured in the 2.0
1.5
¢¢
/D,~ " £]-" ~ ~"O" " - ' ~ "-. . . .
0
-.12.-
// //// /
1.0
1 1 1
I
I 2
I 3
Co/3" TO A FIG. 5.--~z5 in centipoise against C0/Y,T O A for various values o f Y.TOA in benzene (solid lines and data points) and chloroform (dashed lines and open data points). • : Y,T O A = 0 . 5 0 M , • : ~ T O A = 0.25 M, • : Y~TOA = 0.10 M.
temperature range from 8 to 50°C. This decrease is an indication that association takes place in the amine nitrate phase. However, for sols the decrease is generally of the order of a factor two or three. (~°) The amine salt phases were never observed to set to gels. All the experiments point to a certain degree of association in the solutions of (TOAH)NO 3 in benzene and chloroform, but there is no evidence that colloidal species are formed. However, this conclusion is not necessarily valid for diluents with still lower dielectric constant, e.g. hydrocarbons. Here colloid formation might occur and certainly the viscosity effects at low concentrations are pronounced. (1)
Distribution of nitric acid--II
1091
An increase in viscosity by a factor 300 is observed when the concentration of (TOAH)NO a in benzene increases from 0 to 100 weight per cent. In Fig. 6 ~5 incentipoise is plotted against weight per cent for some quaternary ammonium salts, and (TOAH)NOa conforms to the pattern. The quaternary ammonium thiocyanates in benzene show somewhat higher viscosities, (~1) probably because of their larger dipole moments. For tetra-n-butyl ammonium picrate in n-butyl alcohol (2zl the viscosity curve remains below that of (TOAH)NO a, apparently on account of the more polar diluent. 5,10zI 2.10z ~02 50
®
I
2O I0 a (J
+e
5 2 I
05 I
o
51o
IO0
Weight % amine salt
FIG.
--I---© + ®
6.--~hs in centipoise against weight per cent amine salt. (TOAH)NO3 in benzene, tetra-n-butyl ammonium picrate in n-butylalcohol,c~ tetra-iso-amyl ammonium thiocyanate in benzene,'zx' tetra-butyl ammonium thiocyanate in benzene.(~x)
Conductivity
The specific conductance K in ohm -x cm -1 of the amine solutions after equilibration with aqueous HNOa is plotted against Co/ZTOA in Fig. 7. The measurements were performed at 25.0°C. There is a strong increase in conductance in the range C0/Y,TOA < 1 and occasionally at higher acidities. Interpolation at C0/T.TOA = 1 gives us the conductivity data for (TOAH)NOa solutions. The degree of dissociation of the amine nitrate can be estimated from conductivity and viscosity data, according to WALDEN'Srule (~a) 0: = At//60 (4) where A is the molar conductivity and ~7 the viscosity in centipoise. The values obtained are given in Table 2. The degree of dissociation 0: increases with the (=1)R. P. SEWARD,J. Amer. Chem. Sac. 73, 515 (1951). ~2) L. E. STROrqCand C. A. IrOtAUS,J. Amer. Chem, Sac. 72, 166 (1950). (~8)p. WALDEN,Elektrochemie Nichtwiissriger LSsungen Barth, Leipzig (1924).
1092
J . M . P . J . VERSTEGEN
16a ~.,.o"
o
#
10"
E o E
165
1(56
0
I
I
I
I
2
3
Co/Y.roA FIG. 7.--K(~ -z cm -x) against C0/~TOA. Symbols as in Fig. 5. concentration. That, as well as the strong increase in conductance, is characteristic o f salts in low dielectric media.t24, ~s) I t has been suggested t h a t micelles c o u l d be f o r m e d in the amine salt solutions,t26# 7~ a n d it should be k e p t in m i n d t h a t the big increase in c o n d u c t a n c e c o u l d be ascribed to micelle f o r m a t i o n . A t the critical micelle c o n c e n t r a t i o n a steep increase TABLE 2 . - - D E G R E E OF DISSOCIATION g OF
Conc. (TOAH)NO3 (mole/l) 0.50 0'25 0.50
(TOAH)NOs
Diluent Benzene Benzene Chloroform
5 × 10 -a 2 × 10-3 5 × 10-~
tu~ C. A. KRAUSand R. M. Fuoss, J. Amer. Chem. Soc. 15, 21 (1933). ~=5~p. L. MERCmR and C. A. KRAOS,Proc. Natl. Acad. Sci. U.S. 42, 487 (1956). t=e~K. A. ALLEN,J. Phys. Chem. 60, 239 (1956). t=:~K. A. ALLEN,J. Phys. Chem. 60, 943 (1956).
Distribution of nitric acid--II
1093
in conductivity is often observed, as the resistance to motion per unit electrical charge becomes smaller when multiply charged particles are formed. ~'8~ However, the micelle concept in that form is typical of aqueous solutions or at least of water-rich mixtures. Such micelles, with the non-polar carbon chains directed towards the centre and the polar groups, forming the outer sphere and interacting with the diluent, are of course highly improbable under the present conditions. A second objection to the explanation of the conductivity behaviour by means of the classical micelle concept is the wide range of concentrations in which the increase is observed, whereas one would expect it at the critical micelle concentration only. A reversed micelle with the polar groups directed towards the centre and the carbon chains forming the outer sphere has also been suggested, m) Such aggregates would exhibit conductivity behaviour different from the classical micelles and it is not clear how their formation could explain the steep increase in the conductivity. Co-extraction of water In Fig. 8 are given the amounts of water co-extracted when amine solutions are equilibrated with aqueous HNO,. From the experimental results the amounts of water extracted by the amine-diluent mixtures have been subtracted, so that the curves have their origin at zero. The amount of water in the saturated amine-benzene and amine-chloroform mixtures is not significantly different from that in the pure diluents, so formation of a compound TOA-H20 is not likely. The correlation between co-extraction of water and the viscosity agrees with the assumption that co-extracted water and association both stabilize the ion-pair in the low dielectric diluents. The results given here follow the pattern outlined earlier, {1) according to which the co-extraction of water decreases in the diluent with the high dielectric constant although the solubility of the water increases. At higher acidities the water co-extraction increases considerably, as has also been observed in TBP systemsJ29~ A partial explanation lies in the extraction of a hydrated acid species in the diluentJ 3°~ Equilibrations Solubility of HNOa in the pure diluents. HARDY et al. ~°~ showed that the nitric acid solubility in organic diluents can be understood on the basis of a distribution of HNO 3 between water and the organic phase and subsequent dimerization of the acid. Introducing Co, the organic acid molarity and Cua, the undissociated aqueous nitric acid concentration at equilibrium, an equation C O = DCua q- 2Km~o3 D~Cua~ (5) was shown to give an explanation of the processes taking place. Plotting Co/Cua against Cua* the value of the distribution coefficient D is found and that of the dimerization constant KH~ros can be calculated. * Values of C.a were taken from K~WETZ data, reported by Hr~FORD and McK.AY.~sl~ Their original graph, which was large enough to read off fairly accurately was kindly sent to us by Mr. H. A. C. McKAY. i~s~ C. A. KRAUS, J. Amer. Chem. Soc. 60, 129 (1956). izg) K. ALCOCK, S. S. GRIMLEY,T. V. HEALY,J. KENNEDYand H. A. C. McKAY, Trans. Faraday $oc. 52, 39 (1956).
tao)C. J. HARDY,B. F. GREENFIELDand D. SCAgGILL,J. Chem. Soc. 90 (1961). tal) E. HESFORDand H. A. C. MCKAY, Trans. Faraday Soc. 54, 573 (1958).
1094
VE~GEN
J.M.P.J. 0.2~
0.2(
0.15
(--)
o.lo
7//
,
/
/
005
x,,&__ . & /
Y
o
~
I
;
3
Co/ZTOA
FIG. 8.ICO-extraction of H=O in mole/], against Co/~.TOA. Symbols as in Fig. 5.
Values of KHNo8 and D in kerosene and toluene have been reported and some measurements in other diluents have been given. (3°) The work of BERTOCCIand ROLANDI(6) deals with the HNOa solubility in xylene. HARDY'S equation can be applied to their data at the three higher acidities, when the highest value of the parallel determinations at 8.7 M is chosen. The values of D and KH~o= obtained in this way fit nicely into the picture given in Table 3. In column 2 values of the dielectric constant e have been tabulated. TABLE 3.~DISTRIBUTION COEFFICIENT D AND DIMERIZATION CONSTANT KHNOa
Solvent Ker osene Benzene Toluene Xylene o- Dichloro benzene Chloroform
~ N1.9 2.27 2.38 N2-4 9'93 4"8
Distr. coeff. D 6 1 1.2 1"4 1"4 4
x x x x x x
10 -5 10 -8 10 -a 10 -a 10 -a 10 -s
dim. const. K a s % (mole/l) -x 4000 t240 590 300 ~ 100 50
source Ref. This Ref. Ref. This This
30 work 30 6 work work
K=' (mole/l) -z -0-22 0.21 -0'09 if03
Distribution of nitric acid--II
1095
As the isomeric composition of the xylene mixture in Ref. 6 was not given, the value of 2.4 for the dielectric constant is arbitrary. HARDY'S treatment lacks thermodynamic validity, but the results obtained in that way lead to a pattern which can easily be understood from general chemical considerations. In solvents of low dielectric constant the distribution coefficient decreases and the dimerization increases. Chloroform is exceptional in possessing additional possibilities of stabilizing monomeric acid by means of hydrogen bonding. Empirical relations in the high acid range. Concentrations in mole/l are denoted by C, molar activity coefficients by f, the thermodynamic equilibrium constant (based on aqueous and organic activities) by K, and the equilibrium constant based on aqueous activities and organic concentrations by K'. The aqueous HNOa activity 0.12
o.~o
~
]
I
h
O08 o E
0.06
0.04
]3Z
0 02
i
0
I
2
3 Cf,
4
5
6
mole/L
FIG.9.--CcagainstCf for 0"1M TOA solutionsin variousdiluentsat 25°C. - - O - - benzene, - - A - - toluene, - - I I - - o-dichloro benzene, - - O - - chloroform. is denoted by aH~o, and calculated from tabulated values o f T + . (3~) The subscripts a, s and c refer to the free amine, the ion-pair and the 1 : 2 complex respectively. Other terminology is the same as defined before, (1) viz. C i and Cf refer to initial and final aqueous HNO3 concentration and C o = Cs + 2Ce, while ETOA = C a + Cs + Ce. Some typical C~ against Cf curves are shown in Fig. 9. In benzene an empirical relation Ce = 0.19 CfETOA (6) gives a good approximation for 2M < Cf < 6M. The material balance Cf---- C i -- Ce -- ETOA
(7)
combined with Equation (6) leads to
Cf = Ci/[1 + 0-19 ETOA -- function (Y,TOA)] (8) The small deviations from Equation (6) disappear in Equation (8) as 0.19ETOA ~. 1. When C r is plotted against Ci straight lines are found, but small volume changes (s2)
LANOOLr-BrRNSTEIN,Physikalisch-ChemischeTabellen, Eg. IIb, page 1119, Eg. Illc, page 2145.
1096
J.M.P.J.
VERSTEGEN
cause the slope to deviate somewhat from 1/(1 -+- 0.19ZTOA). The empirical relations are given in Table 4. TABLE 4.--EMPIRICAL RELATIONSFOR H N O a EXTRACTION BY (TOAH)NOa-BENZENE PHASES Y,T O A (mole/l.)
Relation
0-465 0-249 0-100
C t = 0 ' 9 2 9 C i -- 0'43
0.050
Ct = 0.975C~ -- 0.01
Cr = 0"967C1 -- 0'22 C r = 0"972C l -- 0"07
Formation of a 1:2 amine-acid complex at high acidities. The quantitative description of the extraction of H N O 3 in the high acid region is based partly on a reconsideration of the results in the ion-pair range (C0/Y~TOA < 1). Under certain conditions the formation of (TOAH)NO 3 could be based on the mass action law. tl) In benzene and chloroform these conditions are fulfilled if the initial TOA concentration _< 0.1 M and the final (TOAH)NO 3 concentration < ~--~0.02 M. In benzene these conditions are not fulfilled at initial TOA concentrations _>0.25 M, even when the amount of organic amine salt remains very small. Here the formation of (TOAH)NO3 takes place according to a mechanism which differs from the usual mass-action expression. In chloroform a similar effect might be observed, but the extractive power for acids of concentrated amine solutions in that diluent is too large to allow accurate measurements. It was reported earlier tl) that in modified dodecane the TOA concentration had to be kept as low as 0.05 M to observe mass-action behaviour. Analogous observations in TOA -- H2SO 4 systems showed that the formation of amine sulphate in benzene occurred on a mass action basis as long as the TOA concentration was <0.25 M. At higher amine concentration (0.50 M) mass action behaviour was not observed.t 2n) The general trend behind these observations is that when the initial amine concentration increases, a critical value is reached, beyond which the formation of amine salt can not be explained by a mass action mechanism. Or, the behaviour of the amine salt depends on the amount of free amine present and that suggests some sort of interaction between the two. If the highest amine concentration with a mass action range is defined as the critical amine concentration (CAC) a table can be made to show the influence of various parameters (Table 5). TABLE 5.--INFLUENCE OF PARAMETERSON C A C C A C (mole/l.) 0-25 <0.10 <0.05 0.10 0.05
amine
a q u e o u s acid
diluent
Tri-n-octyl Tri-n-octyl Di-n-decyl Tri-n-octyl Tri-n-octyl
H2SO4 H~SO~ H2SO, HNOs HNO3
Benzene Kerosene Benzene Benzene M o d . dodecane
source Ref. Ref. Ref. This Ref.
26 33 27 work I
Distribution of nitric acid--II
1097
In comparison secondary amine salts have a large polar group, whereas kerosene and modified dodecane have lower dielectric constants. Probably a lower CAC and a greater tendency towards third phase formation correspond to a greater difference in polar character between solvent and solute. Additional evidence for TOA-(TOAH)NOa interaction is obtained as follows. It has been shown before,t2e,27,a3,a4~ that empirical formulae such as Kemp. = aH~rOa [TOA] ~
(9)
can often describe the results. A relation of the same form with n = ] can be applied to the (TOAH)NO3-benzene system in the concentration ranges where the law of mass action is not obeyed (that means in the whole range 0 < C0/Y.TOA < 1 when Z T O A = 0.25 and 0.50 M and for (TOAH)NOz concentrations > ~ 0.02 M in the cases that Y,TOA = 0.05 and 0-10 M). When an~o, is plotted against TOA on a logarithmic scale, four parallel straight lines are found, and from the slope n can be derived. The lines for different ETOA values are not superimposed. This is analogous to results in sulphate systems as reported by ALLEN (26) who showed that interaction of free amine with the amine salt or its high molecular association products may provide an explanation of the effect. The side reactions discussed above and in the section on viscosity cause deviations from mass-action behaviour, such that in a certain range of concentrations Equation (9) has to be preferred. Another approach to the problem is to express the deviations in terms of the activity coefficients. This is a purely formal treatment, which does not lead to a better insight into the processes taking place, so long as activity coefficient behaviour is not fully understood in the system under consideration. However, in the present case useful information was obtained and the treatment provides evidence for the existence of a 1 : 2 amine-nitric acid complex. The gross equilibrium constant K1 of TOA -}- H + + NO~- ~ (TOAH) NO 3
(10)
is obtained by interpolation from earlier work. tl~
K1 =f~(TOAH)NO3/fa TOA a~ro 3=fJfaK; -= 0.48 × 106 (mole/l) -2
(11)
and
fs/f. = 0.48 × 106/K;
(lla)
When .fJfa is plotted against C0/YTOA the curves in Fig. 10 are obtained. For Y~TOA = 0.05 and 0.10 M a region exists where the processes can be explained on basis of the mass action law. Here the r a t i o f J f a is arbitrarily set equal to one. That does not necessarily mean that in this region (TOAH)NO3 exhibits monomeric behaviour. The gross quantity K1 might contain equilibrium constants arising from possible side reactions. The deviations offa from one in the binary mixture TOA-benzene are discussed in the section on reagents. They are apparently small. An explanation of the shape of the f~/fa curves is based on two considerations. First, mutual association of amine salt molecules causes the thermodynamic concentration Co/n to fall considerably below the stoichiometric one, C O(n is the average number of (TOAH)NOa c38~j. M. P. J. VERSa~GENand J. A. A. KETELAAR,J. Phys. Chem. 66, 126 (1962). t34~ j. M. P. J. VERSXEGENand J. A. A. KETELAAR, Trans. Faraday Soc. 57, 1527 (1961). 13
1098
J . M . P . J . VERSTEGEN
molecules per association product). Second, association of TOA with high molecular amine salt species decreases the amount of free amine present. Both effects can be accounted for by activity coefficients < 1. The fs/f~ curves show three regions of interest. In the first, at low values of C0/~VTOA,fJf~ decreases. Here the amount of unconverted amine is dominant and the formation of amine salt association products has not yet proceeded very far, so that most of the TOA remains in solution. The shape of the curve is mainly determined by the mutual association of amine salt molecules and the corresponding
0.2
0.4
0.6
0.8
1.0
Co/~TOA
Fro. lO.--fJf=
against Co/~TOA in the range 0 < C o / ~ T O A < 1 at 25°C. . . . . . . . . . . ~ T O A = 0.47 M , - ~ T O A ---- 0.25 M ,
. . . . . . . . . ~ETOA= 0.10 M, - - - ~TOA = 0"05 M. decrease offs. At high values of Co/Y~TOAfJfa increases. Here most of the TOA has been converted to the nitrate and appreciable quantities of associated amine salt species are present. The shape of the curve is mainly determined by the interaction of TOA with those species and by the corresponding decrease infa. In the region of intermediate acidities the two effects compensate one another and the ratio fdfa may remain fairly well constant. The same treatment can be applied to the data in the region 1 < Co/~TOA < 2. It is assumed that a 1:2 complex is formed between TOA and HNO3, according to the equilibrium (TOAH)NO8 + H + + NOa- ~ (TOAH)NOa'HNOa,
(12)
The gross equilibrium constant is Ks -----fe(TOAH)NOa'HNOa/j~(TOAH(NOa) arr~ros =
fe/fBK~'
(13)
In Fig. 11 K2' is plotted against Co/ZTOA. Except for a constant the curves represent plots offdfc against Co/ZTOA in the range I < C0/ZTOA < 2 and for different values of ZTOA. There are considerable deviations from ideality. The increase of the fs/fc curves when Co/ZTOA is slightly larger than 1 is probably due to association of the 1 : 2 complex with the amine salt. Another range where fs/fe increases is found at acidities corresponding to
Distribution of nitric acid--II
1099
1.6 < Co/ZTOA < 2. In this range the viscosity decreases and the co-extraction of water increases. A possible explanation is that, on account of the increasingly polar character of the organic phase, the stabilizing influence of association is replaced by that of hydration. The species de-associate to form hydrated lower molecular products. Special interest attaches to the points of intersection of the fs/fe and fs/fa curves with the ordinate Co/ZTOA = 1, obtained by extrapolation. Purely formally they represent the ratio of the molar activity coefficients fs for different (TOAH)NOz concentrations. From Fig. 10 even the absolute values o f f i
t \
o.6~ o~1~\\
0.7
"\
'\
kl
0.4 •
0,2
'"
'~
-" --~"~'~..--~
. . . . . . .
~
.-~'~/
0.1
0
I I2
I 1.6
L 1.4
| 1.8
2 0
CoI'~TOA FIG. II.--K~' against C0/~TOA in the range 1 < C0/ETOA < 2 at 25 ° C. Symbols as in Fig. 10.
can be derived, providing the assumption is valid that fJfa--~ 1 when Co ~ 0. Although the values of fs obtained in that way are probably not meaningful, the extrapolations to C0/ZTOA = 1 from the low acid side (0 < Co/Y.TOA < 1) and from the high acid side (1 < C0/ZTOA < 2), yield the same ratio of fs values for different amine salt concentrations. This provides some evidence that a 1:2 complex is indeed formed. Further reasons for believing that when 1 < C0/~TOA < 2 formation of a 1:2 complex is predominant will be given now. First, when Cf ~-, 6 M (where Co/Y.TOAis slightly larger than 2) new phenomena occur, indicated by an increasing acid extraction, which was not reproducible. Extraction of HNO 3 by the diluent does not account for the whole effect. This observation provides another example of anomalous extraction behaviour such as was first reported by McDowH~L et al. (~) Energy differences between possible 1:3 and 1:4 complexes are probably small, and after phase settling non-equilibrium species may still exist, because of slow diffusion to the interface. A second reason to believe in tss) W. J. McDOWELL and K. A. ALLE~, J. Phys. Chem. 65, 1358 (1961).
1100
.l.M.P.J. V~RST~GEN
the formation of the 1:2 complex is the rough proportionality between ETOA and excess acid concentration. Consequently about the same ratio Ce/Cs is found at the same value of Cf, when the amine concentration varies from 0.05 to 0.5 M. Third, at least in the region I < Co/ETOA < 1.6 the Ks' values approach constancy, when the amine nitrate concentration decreases (Fig. 11). A fourth reason is the influence of the diluent on the formation of a 1 : 2 complex, as shown in Fig. 9. In Table 3, column 6, values of K s' are compared at Ce/Cs = 1 and for 0.10 M TOA solutions. It has been suggested tr's) that the distribution of HNO 3 between aqueous solutions and amine nitrate phases can be explained as simple dissolution of the acid in a more polar organic solution. The trend given in Table 3, column 6, is at variance with this conception. The amount of HNO3 extracted into the organic amine nitrate phases, increases with decreasing dielectric constant of the diluent, while in the pure diluents the reverse is true. If the role of the amine salt is only to lend polar properties to the diluents, one would expect the original sequence to be maintained. In view of the data presented above, a better explanation of the distribution at high acidities is based on the stabilization of the amine nitrate salt, which is formed at low acidities. (~ The ion-pair needs to surround itself with stabilizing molecules, especially in the low dielectric diluents. Nitric acid may act as a stabilizing molecule and it will be extracted to a greater extent when stabilization by the diluent is unimportant. Chloroform shows stabilizing properties, much greater than would be expected from its dielectric constant. The amine nitrate salt is probably solvated in that diluent and the interaction takes place by hydrogen bonding. From our measurements at 4 and 25°C in benzene systems it was found that the distribution at high acidities was insensitive to temperature. The same was found at our Institute for the distribution of HNO 3 between aqueous solutions and a solution of tri-lauryl amine in modified dodecane. (36~ The main impression is that the second HNO 3 molecule is bonded to the amine nitrate by means of hydrogen bonding. The 5 kcal/mole gained in that way, could be balanced by side effects, such as de-association, partial dehydration etc.
Infra-red spectra It has been postulated (1) that higher complexes are formed in the amine nitrate system even in cases where as much as 40 per cent of the amine remains available for salt formation. From the present paper the conclusion may be drawn that the equilibrium constants for the first two complexes differ by a factor of 10e. Therefore, no appreciable amounts of free amine exist in equilibrium with 1:2 complex. The prior erroneous assumption, which does not affect the general trend reported earlier, <1) was primarily caused by interpretation of the C i against Cf curves. Typical plots in benzene showed a rapid increase of Cf but in the present paper it is shown that this is due to the shape of thefJJk curves (Fig. 10). Infra-red spectra provide evidence that after reaching a value Co/]~TOA = 1, the salt concentration remains constant. When spectra are recorded for (TOAH)NO 8 in CHCla, compensated with TOA in CHCI 3, the chloroform CH stretching frequency is found at 2998 cm -~ and a range of transmission >100 per cent is observed at 2780 cm -1. The explanation is as follows. In the reference cell, some CHCI a molecules (36~
R. BAc, Institutt for Atomenergi. Personal communication.
Distribution of nitric acid--II
1101
are hydrogen-bonded to TOA. This causes the CH frequency to split into two branches, a low frequency component at 2780 cm -1, due to interaction with the amine, and a component at the usual frequency of ~ 3 0 1 9 cm-1. taT~ In the sample cell HNOa replaces CHC13 in its interaction with TOA, but now solvation of the amine nitrate takes place. Here the CH frequency is split into a branch at ,-~ 3019 cm -~, compensated in the reference beam, and one at lower frequency, due to interaction between the chloroform hydrogen and the amine nitrate oxygen. The A
B I
T
Trr
I 28
i 25
I 22
I 28
I 25
I 22x102 cm-I
FIG. 12.--The shape of the N÷-H . . . . (ONO2)- band at different organic acidities
C0/]~TOA in chloroform (A) and benzene (B) as diluent. I: Co/5'.TOA< 1, II: Co/~TOA ~ 1, III: Co/ETOA :> 1. latter emerges from the spectrum as a peak at 2998 cm -1, with an intensity proportional to the (TOAH)NO 3 concentration. It was found that after reaching a value C0/ETOA = 1 the intensity of the peak, and consequently the (TOAH)NOz concentration, remained constant. Interaction between the amine nitrate group and the second H N O 8 molecule may lead to a weakening of the central hydrogen bond in the amine salt ion-pair. Indeed it was observed that the N + -- H . . . (ONOz)- band at 2430 cm -x in the spectra of (TOAH)NO3 in benzene and chloroform ~1~ shifts towards higher frequencies when C0/]~TOA > 1. In CHC13 this leads to a splitting with a new frequency at 2560 cm -1, which is found as a weak shoulder when Co/ETOA < 1 (Fig. 12). Nitric acid in the range 1 < Co/ETOA < 2 and CHC13 in the range 0 < C0/ETOA < 1 form hydrogen bonds with the oxygens of the amine salt anion. There is a competition between excess H N O 3 and CHCI~, consistent with the low equilibrium constant found for the 1:2 complex in that diluent. The observed infra-red frequencies of excess HNOz (DNOz) are given in Table 6. The spectra are difficult to interpret, because in the region of interest the frequencies of the ion-pair nitrate shift under influence of the excess acid. Assignments were made using References 38-40. ~37~G. HERZBERG, Infra-red and Raman Spectra of Polyatomic Molecules p. 316. Van Nostrand, New York (1945).
1102
J. M. P. J. VEP-,STF,GEN
TABLE6.--INFRA-REDFREQUENCIESOF EXCESSHNOs AND DNOs IN (TOAH)NOs Excess HNO 8 in ( TOAH)NOa--CeH e (cm- 1)
Exc,~s HNO s in (TOAH)NOa-.CHCI 8
Excess DNO s in (TOAH)NOa-CeH e
(eat-1 )
HNO s yap.
(era -:t)
HNO$ liq.
D N O s DNO a Assignyap. liq. mcnt (ss-4o)
(cna-1)
~3100 1645
~3200 1640
2230 1609
3560 1710
3400 1675
2627 1685
2470 1645
1295
1310
1298
1320
1300
1313
1300
948
938
938
886
925
888
915
HO stretch NO asymm. stretch I~O symm. stretch NOH stretch
The frequencies suggest the structure f3
(CsH17)s N--H. .... O - - N ~
. . . . . . HONO~
(electrical charges omitted).
Acknowledgement--This paper is published by permission of the Managing Director of Institutt for Atomenergi and the work forms part of the programme of a Dutch-Norwegian Joint Project. Mrs. C. M. KOFP~RtrD,Mr. I. A. HtrgoERE and Miss W. A. TAP assisted in various facets of the work. css~C. K. INGOLDand D. J. MmLEN,J. Chem. Soc. 2612 (1950). ts,J H. Cot]n, C. K. INGOLDand H. G. POOLE,3". Chem. Soc. 4272 (1952). t40~R. A. MARCUSand J. M. FRESCO,./. Chem. Phys. 27, 564 (1957).