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Distribution of nitric acid between water and amberlite LA-1 liquid anion exchanger
J. Imrs. NucLChsm., 19~2, VoL 24, pp. 1417to 1428. PmllamoaPmmLtd. Prlatsd ia SaSiS~!
DISTRIBUTION OF NITRIC ACID BETWEEN WATER AND A M B E R L I T E LA-1 L I Q U I D A N I O N E X C H A N G E R A. S. K~Tm* and I. T. PLATZNn Departmmt of Inoqpmic and Analyt/c~ ~ , The Hebrew University of Jaqumlem, J e ~ b m % Israel (Rec¢iwd 20 D~ember 1961) AbUrad--The solute-solvent inUnctkm betweemAmberlite LA-I, a Ions-chain secondary Amineextractant d|molved in carbon Wtrmchiceddeor bennme and 8qtmeus nitric acid wluti~n up to ~ 16 M has been inveetilsted by physical end chmni~ mmumrmn~ts. The flint reaction is an acid-base equilibrium leadin_a to the formation of amine nitrate. The ~ comtlmU were evahmted u El - 3.8 x 10s and $-4 x 10s in carbon tetrachlott~ and bemme m~m~vely. The meond and third.mactiom immured involve 8mine aitmte and meciated nitrk scid nmlmdm, leadin8 to the f e r n m ~ of a molemlar a d ~ a mmlmmtd/dmtifi~ as It4HNI~OrHNOs and to the ~ l a r t / t ~ of m e e i m d
nitria.c/d ~
the .qumm and o r s . ~ ptme.. The ~
~
p~h~
quotiemm Imve thz wdue8 JC2 -- 1.85 for carbon tetrachloride and K2 ffi 2.24 for end the ~ constant Ires the value F~d -- 0"I _+0~ for I M amim ~ in both dilmms. The eSeets of the formation of the amine nitrate ~ or of the retracted m o l e m ~ addit/on ~ er both, are not takm into ~ t . Bvidmm w u elmlmd that the q l r e ~ t i o n of the speciss in the orsank pham is larllely a peet-equililm~m
~mmon. IT has been observed(l-~ that the extraction of nitric acid by Ions-chain amines is not simply a case of an acid-baN reaction hadin8 to the formation of amine nitrate, since potentiometrk: neutralization curves of the acid contained in the equilibrium Orllanic phase give two endpoints. The amount of base consumed from the first to the second endpoint is equivalent to the amount of amine present, and the amount required to reach the first endpoint is a function of the conomttmtion of nitric acid in the aqueous phase at equih'brium. This ind/cates that the acid is bound in two di~emnt ways. * PmJmt addmm: l ) q m m t m t of O ~ t r y
and L~J~ratory for Nuclear 9cimee,
Msmschmem Insl/tute of Technelo~, ~ 39, Mass. I~) E. A.-h4lAsoN end V. C. A. VAVOmU~,Report A.~U-4~31 (1959); Report TI]D-ST20 (1960); V. C. A. VAVOmU~and E. A. MASON, Summary Rel~rt, Dept. of Nucleer. lVLLT.(I~0). {2) U. Bmt3~cx~ AEKE/R-2933 (1959).
(3) D. J. Ctawwma. and J. J. I.~WUNCS,7./mort. Ntr.t C/ram. 11, 69 (1959). (4) D. E. ~ snd C. F. CctattaN, Report OKNL-30$1 (1961). (s) (a) G. SCmONA,It~U~ Comm./~'~i. Jets., Report CNC-43 (1960); (b) G. SCatONAand M. Zn,mmmm,/b/d. CNBN-$1 (1961). (e) V. B. ~ , V. S. Sin/Ira', E. A. Nl~AmOZOtmVand K. A. F~,,tov, Z~. N~oqr. ~. S, 18.52 (1~0). (~) O. I. ZAH~OV-N~tcmOV and A. V. Ocmmq, 2~. Neoq,. ~k/m. 6, 1936 (1961). 1417
1418
A.S. KeR'r~s and I. T. PL~TZNmt
Several studies have been published in the last two years on the extraction of actinides¢2, 3, 4, 8), lanthanides,¢9) technetium and rhenium,C10, 11) some fission products¢l, 12, 13~ and heavy metals04~ from nitrate systems by amines. Some work has also ~ devoted to the r61e of nitric acid. In MASON'Slaborat0ry¢l~ the extraction of nitric acid in the range 2-10 M by a number of primary, secondaryland tertiary amines has been investigated. The nitric acid in the organic phase in excess of that equivalent to the amine concentration was found to be proportional to the concentration of the acid in the equilibrium aqueous phase for the tertiary amines, but a less simple relationship was observed for primary and secondary amines. CARSWELL and LAWRENCE,¢3)BERTOCCI(2) and COLF~N and HORNER¢4~ provide support for the linear relation with tertiary amines, and according to the work at Oak Ridge the linear relationship appears to hold for all amines, including primary and secondary ones.¢4) Thus, the concentration of extracted nitric acid can be related t0the concentrations of amine and aqueous acid by the expression (HNO3~xc~)ors/(Amine)ors(HNO3)~q ----ce where (HNO3e~m) is the organic nitric acid concentration above the stoichiometric ratio 1~1 with the amine. The available results fit this relatio~hip with values of c~between 0.16-0.18 for tertiary, 0,10-0.14 for secondary and 0-06-0.11 for primary amines. We have found that Amberlite LA-I, a secondary amine, follows the above relationship in the range between 3 and 12 M nitric acid in the initial aqueous solution, the constant u having a value of 0.13-0.14. Although this relationship is valid for a ~vide range of acid loadings and amine concentrations it cannot be interpreted as the mechanism of the reaction, and represents only an empirie~ relationship. SHEVCHENKOet al.¢6~ have studied the extraction of nitric acid by trioctylamine, and propose as the extraction mechanism complex formation between the amine nitrate and excess nitric acid according to the reaction (TOA'HNO3)o~-Ha+ ~(NO~-)a --* [(TOA'HNO3)'HNO3]o The calculated equilibrium constants have values of/C--~ 0.09 and 0.13 for the diluents carbon tetrachloride and xylene respectively, K ~howed satisfactory constancy in the range of 1-5 M aqueous nitric acid, but was measured for only one Cs~W, E. KEDER, J. C. SHEPPARDand A. S. WILSON,jr. Inorg. Nucl. Chem. 12, 327 (1960); J. C. SHEPP~,tu, Report HW-51958 (1957); A. S. WILSONand W. E. KEDmX,Report HW-62908 (1959); A. S. WILSON,Report HW-68207 (1961); W. E. KEDEn, J. L. RYAN and.A.S. WZLSON,Y. Inorg. Nuel. Chem. In press; F. IcmKAwA and S. URUNO,Bull. Chem. Soc. Japan, 33, 569 (1960); (~. J. HANDY, D. SC~XOILLand J. M, FLETCHER, J. Inorg. Nucl. Chem. 7, 257 (1958);'1"K. B. BROTHEaS,C. F. COLEMAN,D. J. CROUSE, J. V. DENmand J. G. MOORE,Y. Inorg. Nucl. Chem. 7, 85 (1958); H. G. PETnOW,Report TID~5357 (1960); F. R. BRuce, R. E. BLAN¢Oand J. C. BX~SEE,Report CF-58-11-91 (1959); B. WEAVERand D. E. HORNER,J. Chem. Engng. Data, 5, 260 (1960); M. DE TRENTrNL~Nand A. CHESNE,Report CEA-1426 (1960). (9) F. ICHIKAVA,Bull. Chem. Soc. Japan, 34, 183 (1961). CJ0~G. E. BOyD and G..V. ~LAt~SON,?. Phys. Chem fCJ, 988 (1960); A. S. KERTESand A. BECK,J. Chem. SbC. 1926 (1961); A. BECK. Ph.D. Thesis, The Hebrew University, Jerusalem (1961). ~1.~F. L. CULLmLAnnual Progress RepotL ORNL-2993 (1960). ¢t2~E. A. MASONand V. C. VAUGrmN],Report AECU-4239 (1959); Report AECU-4351 (1959). ¢13~V. B. SHEvcHEt,n~oand V. S. SMIVT, Radiokhimija, 3, 121 (1961); W. KNOCHand R. LXNDER,Z. Elektrochemie, 64, 1020(1960); W. KN]ocH,Z. Nat urforschung, 16a, 525 (1961). (t4) E. A. MASON]and R. E. Sg.AVDALL,Report TID-11196 (1960).
Distribution of nitric acid between water and ~ t e
LA-1 liquid anion exchanger 1419
concentration of amine (0.47 M). An unsuccessful attempt was made to: fit the results obtained in this study to ~ ¢ above reaction; the calculated, values o f K were not satisfactorily constant. In a,papar published while this report was being completed, ZAHAROV-NAgcISOV et al.~ have interpreted the equilibrium data in the system 0.25 M triheptylamine n!trateraqueous nitric acid (1.3-4 M) by the extraction of undissociated nitric acid, and postulated the existence of the species (THA.HNO3)HNO3 wittha formation constant K - 7.2=I=0.2. These authors also found that in carbon tetrachloride the organic solution of amine or amine nitrate is non-ideal; and constant values for K could not be obtained. The purpose of this investigation w ~ to ascertain whether the amine nitrate and nitric acid interact, and if so to what extent. In addition to a study of equilibrium data for the acid distribution, this investigation is devoted particularly to the physicochemical behaviour of the loaded organic layer, its viscosity, density and conductivity: Similar studies with phosphoric acld esters as extractants(XS, 16) have been successfully employed in. clarifying the heterogeneous equilibria in these solvent extraction processes. Amberllte LA-I was chosen as extractanL since it is one of the few orga~onitrogen compounds selected as suitable for process use.(l~ EXPERIMENTAL Materials. AmberHte LA.I* is a secondary amine, with a molecular weight of'4t6as determined in a potentiometric titration using per~loric acid.(ts) In the free base form,:a~ supplied bythe manufacturer, it is a viscous, amber-coloured liquid; very slightly soluble in aqueous solutions (10--20 p.p.m.). The diluents used were carbon tetracMoride and benzene, both of the highest quality, and they were used without further pm-illmti0n. Extraction procedure. Equal volmnes, urmtlly 5 ml, of the amine solution and the aqueous nitric acid solution were mixed in a t5 ml graduated centrifnge tubeandshaken mechankally for 15 vain. Certain experiments in which a different technique was employed are demm-ibedlater. After equilibration with 0-89M amine solutions when acid solutions of concentration greater than ,~ IOM were used, swe~i'na of the orllanic phase occurred irrespectivelyof the diluent. In these systems,prolonged centrifuilation at maximum speed wasnecessary tO wodu~ phase separation, in order.to enable volume .change to be measured. As a precaution against changes observed inthe organic phase which are described later, thesamples were treated one after another through all stages of the work, and in addition were, in so far as possible, protected from light. Work on each .sample took about 30 rain. As preliminary experiments had shown only a slight tomperature dependence, extractions were performed at room temperature (22-26°C) ~ where otherwise stated. Colour c h a n ~ , from light to dark brown, were otserved on ~ the amine solution into contact with the more concentrated nitric acid~ No attempt was made to investigate this phenomenon. Analytical procedure. Aliquots of the phases were diluted~ the o~ganic phase with an 8:5:1 methanol:acetone:water mixture, and the acid content determined potmtiometrically * We are indebted to Messrs. Rohm and Haas Company for kindly providing a free sample of the amine. ~ls~ A. S. K ~ T u , J. Inorg. NucL Chem. 14, 104 (1960); A. SI KIUtTnand V. K.~TES,Canadian J. Chem. 3S, 612 (1960); Y. AppL Chem. 10, 287 (1960); A. S. KEST~Sand M. HALPEaN, J. Inorg. Nucl. Chem. 16, 308 (1961). ~t6) A. S. K~T~S, A. BEctc and Y. HAnOUSHA,J. Inorg. Nud. Chem. 21, 108 (1961). (tT~C. F. COLEMAN,K. B. BROWN,J. G. Moon~ and D, J. Qtous~, Industr. Engng. Chem. 50, 1756 (1958). (ls~ CH. Bont~, Bull. Soc. C/u'm. France, 980, .1088 (1958); J. BIZOTand B. T~LLON, Bull. Soc. Chim. France, 122 (1959).
1420
A. S. KERTES and I. T. PLATZNBR
using a Beckman Zeromatic pH-meter. Two endpoints were observed in titrations of amine phases which were in contact with more concentrated nitric acid solutions. Turbidity, which does not interfere with the pH readings, appears just at the point when the excess acid (over the stoicheiometric ratio of 1: 1 amine to acid) is neutralized. Blank extractions with the diluents alone indicated that appreciable amounts of nitric acid are extracted into benzene, (5. 19) particularly at the higher acid concentrations (> 10M). The benzene blank based on the volume percentage of benzene in the organic solution, determined separately, was therefore subtracted from the total acid in the organic phase. The values for organic acid content given in the tables and graphs throughout this work are those for extraction by the amine alone. Physical chemical measurements. The density, viscosity and conductivity of the homogeneous organic layers were measured immediately after separation, according to the procedure previously described. US) Measurement of amine nitrate formation constant. The equilibrium constant for the acid-base reaction leading to the formation of amine nitrate was studied by converting the base into the nitrate, equilibrating the amine nitrate dissolved in the organic phase with water, and measuring the hydrogen ion activity in the aqueous phase. us. 20) A pure amine salt solution in benzene or carbon tetrachloride was prepared by converting the base to the nitrate form by shaking it with an aqueous nitric acid solution of slightly greater acidity, and removing excess acid by three washings with aqueous nitric acid solution of slightly lower acidity. The results referred to later in Fig. 5 show that the amine nitrate so obtained in the organic phase contained neither free amine base nor free nitric acid. A series of dilutions of this amine nitrate stock solution with the appropriate diluent was made down to 0,0015 M. Ten ml aliquots of the amine nitrate solutions were equilibrated with equal volumes of triply-distilled water, and the hydrogen ion activity of the aqueous equilibrium phases was determined by pH measurements using a Cambridge pH-meter to an accuracy of 0·01 pH units. Degree ofagitation and "post-equilibrium" changes of the loaded organic phases. ALLEN and McDoWELL(21l have observed that in amine systems the final distribution of the solute is not entirely independent of the way in which equilibrium is reached. Following the experimental procedure suggested by them, results were obtained for the extraction of nitric acid by both vigorous and gentle agitation at 27° to·l°C. The results presented in Table 1 show that the organic acid loading varies with contact time and type of agitation TABLE 1.-DISTRIBU110N OF NITRIC ACID AT VARIOUS INITIAL CONCENTRATIONS BETWEEN WATER AND 0,45 M AMINE IN CARBON TETRACHLORIDE AS A FUNCTION OF TIME AND TYPE OF AOITATION Slow equilibration Violent Initial equilib. aqueous NaOH consumed for 5 ml acid(M) of the organic phase 1 day 3 days 7 days 15 mins 8·75 mlO'259NNaOHforaminenitrate 8·75 8·90 9·05 9·80 mlO'259NNaOHfortotalacid 9·70 9·80 9·85 2-154 [HN03]0, total organic acid loading 0·504 0·508 0,511 0·508 8·75 8·75 8·90 9·40 mlO·259 NNaOH for amine nitrate 12·05 11·80 11·95 12·55 4·194 mlO·259 N NaOH for total acid 0·633 [HN03]0, total organic acid loading 0·613 0,620 0·652 8·75 8·75 9,00 9·45 mlO·259 NNaOH for amine nitrate IH5 14·85 14·95 15·70 6'518 mlO·259 NNaOH for total acid 0·784 [HN03]0, total organic acid loading 0·772 0·776 0·815 8·75 mlO'259NNaOHforaminenitrate 8·75 9·15 9·45 17·80 17·40 17,50 18·25 8'527 ml 0·259 N NaOH for total acid 0·923 [HN03]0, total organic acid loading 0,904 0·908 0·946 C. J. HARDY, B. F. GREENFIELD and D. SCAROILL, J. Chem. Soc. 90 (1961). (20) L. NEWMAN and P. KLOTZ, J. Phys. Chem. 65t 796 (1961). (21l K. A. ALLEN and W. J. McDoWELL, J. Phys. Chem. 64t 877 (1960).
(19)
Distribution of nitric acid between water and Amberlite LA-I liquid anion exchanser 1421 in a manner similar to that observed by ALL~ and McDOWeLL. Furthermore, as our results show, the amount of base necessary to neutralize the amine nitrate varies with the time of contact of the two phases. In addition, the loaded orlpmic phase, previously separated from the aqueous layer, and containing niti~c acid in conc~tratiom above.the I'1 stoicheiometric ratio, undergoes chanlles both in density and viscosity; Fig. 1 shows v i g ~ i t y meamrements
560
520
/"
u 480 -
/o
nf ° ,~
440
i
280
200 0
I
i
24
t
t
i
I
411 72 Aqelnq time,
]
[
96
I
, t20
hr
~ o . 1.--Pcat-equfl/brlum c h a n l ~ in the v t s c ~ t y of the loaded ~ (0.89 M amine equilibrated with 9-24 M nitric acid solution).
phase
carried out on allein8 solutions. A plot of time of emux in an Ostwald vi~zeity pipette is 8iveu rather than true v i ~ e i t y values, since diflimltim were encountered in density determinatiom using the usual Spreusel type pycnometer. The cause of change in the organic phase has not been det~...'ne~L Possible causes are re-ananllemeuts in the orsanic phase to yield different final amine-nitric acid species, which may be due either to some unknown post-equihlxium ~ t i o n effect,, s. = ) o r t o simpler chemical reactions such as oxidation or nm'uctk~ Tlm~ore, we do not feel that the anomalous solvent extraction equih'bria are emirely due to violence of allitation, as sulmested by A ~ J ~ and McDonaLd.c21) As additional evidence for this view, we point out that ueither the time nor the way of mixing was found to be critical in at least two different systans u~ing other eu~aes, viz. amin~ n i t r a ~ nitrate(z) and tertiary amine-sulphuric acid.(23) In any case, this post.equilibrium chanlle, shown in Table 1 and Fig. 1, raises considerable doubts concerning the applkal~lity of the equilibrium results for ailed solutions to an interpretation of extraction mechanism. It was t b e ~ o r e decided that rapid extraction equilibria would be simpler to interpret than those performed under chanlling conditions. Conaeqtle~tly, 15 min of violent shat-lna Was taken as an adequate and conveuieut sh__akin8time for valid equilibrium interlmmttion. (22) K. A. ~ , 3. Phys. Chem. 60, 239, 943 (1956); 62, 1119 (1958); 3. Amer. Chem. Soc. 80, 4133 (1958); W. J. McDowm.L and C. F. BAss, J. Phys. Chem. 67, 777 (1958). (z3) j. M. P. J. Vm~rmmq and J. A. A. ~
Tram. Faraday So¢. $7, 1527 (1961).
1422
A.S. K~TES and I. T. PLATZNER RESULTS A N D DISCUSSION
Physico-chemica/ properties of the organic layer. The density and viscosity o f the equilibrated organic phases as a function of the ratio [I-INO3]o/[R2HN]0 are shown in Figs. 2 and 3. The curves for carbon tetrachloride diluent (0.89 M amine) each 1'30
eje" •
_J E t~ "o >,°
1"28
/
1'26
/
e/
CCL4
7
I.:>4
S t t CsH6 0'9;
c c3 0-90
0'88
o.86
0
I
I
I
2
I
3
[ H N 0 3 ] o /rR2HN]o ratio in trle org pnose
FIG. 2.--Density of the equilibrium organic phase as a function of the nitric acid:amine ratio in the organic phase. Initial amine concentration 0.89 M.
/\\ ,...'
o.
I(~
~.
i'
~
_
CCL4
s 4
"t' C6H6
Z 0
i
I
[HN03] o /[RzHN]o
I
2
I
3
4
ratio in the or 9 phase
l:~o. 3.--Viscosity of the equilibrium organic phase as a function of the nitric
acid:amine ratio in the organic phase. Initial amine concentration 0.89 M.
Distribution of nitric acid between water and Amberlite LA-I liquid anion exchanger 1423 show two intersections between straight lines of different slopes. The first intersection corresponds to the amount of acid necessary to neutralize the amine in the organic phase, producing the amine nitrate, R2HNHNO3. The second intersection is at a ratio of 2:1, indicating that the amine nitrate has interacted with more nitric acid to give a molecular addition compound which may be tentatively identified as R2HNHNO3.HNO3. The corresponding curves for benzene diluent are only plotted to a ratio of about two, since appropriate corrections could not be made for the benzene blank. Measurements were carried out also with 0.45 and 0.25 M amine solutions in carbon tetrachloride, but the high diluent content in these solutions caused the viscosity and density values to level off.(7) A change in reaction mechanism is also indicated by the conductivity curves in Fig. 4. Both the amine and the amine
•0
9
C¢L41
~ 7g u "o
g
-
._u
-
• C~%,"
/
o. se SQ 0
t 2 3 [HNOB]e/[IR~H.]o rotio in the orq phose
FIG. 4.---Specific conductance of the equilibrium organic phase as a function of the nitric acid:amine ratio in the organic phase. Initial amine concentration 0.89 M. nitrate solutions showed no detectable conductivity with the equipment available, and the onset of conductivity is believed to be due to formation of a conducting species, such as the addition ~'ompound tentatively defined above. This suggests that the extent of ion-pair dissociation m the organic phase is likely to be important. A rough estimate of the acid dissociation as calculated by WALDEN'Srule (//2/60), shows that it does not exceed five per cent at the highest acid and amine concentrations. Partial ionization of tributyl phosphate-nitric acid compounds at the highest nitric acid concentrations, found from infra-red absorption data by TUCK,(Z4) may be indicative of some similarity between the two systems. Formation of amine nitrate. The neutralization reaction between amine and nitric acid up to a [HNO3]o/[R2HN]0 ratio of unity may be represented as (R2HN)o+ H + + (NO3-)a ~ (R2HNHNO3)o (1) where R2HN stands for the amine base and R2HNHNO 3 for the amine nitrate, and the subscripts --a-- and --0-- indicate the aqueous and organic phases respectively. ~2~)D. G. TUCK,J. Chem. Soc. 2783 (1958).
1424
A.S. K~TeS and I. T. PLATZ~
The thermodynamic equilibrium constant for this reaction is: 7AN.[R2HNHNO3]o (2) KI 7A:[R2HNJo{H+}a{NO~}, ?AN is the activity coefficient of the amine nitrate and YAis that of the amine base both in the organic phase; their ratio can be taken as unity, since they refer to low concentrations and to species that are relatively alike with respect to the medium. Under the experimental conditions described, {H+} a, the hydrogen ion activity, can be taken as equal to the nitrate ion activity {NO;}a. and is also numerically equal to the concentration of free amino base [R2HN]o. Assuming ideal immiscibility of the phases, [R2HNHNO3]o equals the initial amine nitrate concentration less {H+},. Thus, transforming (2) into a logarithmic form, we obtain --log ([R2HNHNO3]~--{H+}a) ---- 3pH+pKI (3) A plot of the leA-hand expression vs. the pH readings should give a straight line with a slope of three and an intercept equal to pK1. Fig. 5 shows such plots, with intercepts K1 = 3.8)< 10s and K 1 = 5.4× 10s for carbon totrachlorido and benzene media
respectively.
CCt4 -3"0
-2.5
o z
-2"0
Z "I" N
.~. -,.5
-I.0
[
,.s
n 2.0
I
I
J
2.4
2.8
3.2
pH
l~O. 5.bPlot of Equation (3) for the determination of/{'i. As seen from the plots, the method of calculation obviously fails when the amino nitrate concentration exceeds 0,05 M. Since deviations from linearity increase gradually with the amino salt concentration in the organic phase, it appears that the limitations are imposed rather by the inability to determine activity coefficients for the organic solutes, than by a chango in the mechanism.(20)
Distribution of nitric acid between water and Amberlite LA-1 liquid anion exchanser 1425
Extraction of acid by amine nitrate. The extent of the extraction of nitric acid by amine nitrate as a function of its equilibrium aqueous concentration is shown in Figs. 6 and 7 for carbon tetrachloride and benzene as diluents, respectively.
/ e/
...J"
~
. ...//'J" o ~o...
"
0"89 M
/
./07/°/~,. j,~,,=o ,,,~.,~
J 25M
0
L.
I
I
4
I
I
I
8
I
12
J
[H NO3] ° M
Fzo. 6.--Eqnilibrium distribution of nitric acid between water and carbon tetrachloride solution of the amine. 3
I
Os
I
1
at-
/
1
0
/×
/
:o.., 1 2' L
/
o°"
I
/ /
• Se'e
4
8 [H NO~]o,
12 M
I'~. 7.--Equifibriurn distribution of nitric acid between water and benzene solution of the amine.
1426
A. S. KERTESand I. T. PLATZNER
It was found that 18 M nitric acid is completely miscible with undiluted amine nitrate; this might be taken as evidence for the extraction of associated nitric acid molecules rather than ionized nitric acid. Similar observations and interpretations were made by SHULER(25)and by COLLOPYand co-workers(26) in nitric acid extraction by tributyl phosphate. The amine nitrate forms a molecular addition compound with associated nitric acid and the amine nitrate-nitric acid species formed may be an amine analogue of the tributyl phosphate-nitric acid species TBP.HNO3 or even TBP.(HNO3)2. The comparable shapes of the isotherms for extraction of nitric acid by TBP and by amine nitrate emphasise this similarity (compare Figs. 6 and 7 with Fig. 2 of reference 25). The above reasoning implies the existence of the equilibrium (R2HN HNO3)o + (HNOa)a ~ (R2HNHNO3.HNO3)o (4) with the equilibrium concentration product quotient K2 = [R2HNHNO3.HNO3]o (5) [R2HNHNO3]o[HNO3Ja where [HNO3]a is the concentration of undissociated, or ion=paired, nitric acid in the equilibrium aqueous phase, which can be calculated from the data of KRAWETZ.(27) Knowing both the total organic acid concentration and the initial concentration of amine in the organic phase, we may write [R2HNHNO3.HNO3]o = [HNO3Jo--[R2HN]o and [R2HNHNO3]o = [R2HN]o--[R2HNHNO3.HNO3]o -----2[R2HN]o--[HNO3]o Substituting these values in Equation (5), values for the quotient K2 have been calculated (Table 2). The constancy of the values of K2c ° ' = 1.85 and K2c'n' = 2.24 is considered remarkable in view of the fact that both ionic strength and the media change markedly as the aqueous nitric acid content increases, and considering that the activity coefficients of the solutes were taken as unity. The mass-action equilibrium (4) obviously fails to explain the extraction mechanism when the ratio [HNOa]o/[R2HN]o exceeds two. There are two possible ways in which further nitric acid, exceeding the above ratio of two, can be extracted into the organic phase. The first is by partition of associated nitric acid between its aqueous solution and the amine nitrate-nitric acid complex, and the second by a further complex formation reaction leading to the formation of molecular addition compounds of the type R2HNHNO3.(HNO3)~. with x ,> 2. It is not possible at present to decide which process is responsible for the phenomenon. This fact is not surprising, in view of the widely contradicting interpretations of solute-solvent interaction in the more thoroughly studied system tributyl phosphate-nitric acid (reference 25, and references therein). However, since the physico-chemical measurements on the loaded organic phase (Figs. 2 and 3) do not indicate any further compound formation beyond R2HNHNOa.HNO3, we believe that nitric acid is extracted by amine nitrate by two (25) W. E. SHULER,Report DP-513 (1960). (26) T. J. COLLOPYand J. F. BLUM,J. Phys. Chem. 64, 1324 (1960); T. J. COLLOPYand J. H. CAVENDISH,3". Phys. Chem. 64, 1328 (1960). (27) W. J. HAMER(Editor), The Structure of Electrolytic Solutions, p. 43. J. Wiley, New York (I 959).
Distribution of nitric acid between water and Amberlite LA-1 liquid anion exchanger 1427 TABLE 2.--CALCULATEDEQUILIBRIUMCONCENTRATIONPRODUCTQUOTIENTEOR THE REACTIONOF ASSOCIATEDNITRICACIDWITH AMINENITRATEUSINGEQUATION(5) HNO3 (M)
HNO3 (M)
[R2HN]0 Initial Equil. Equil. (M) aqueous aqueous org.
/(2
[R2HN]0 Initial Equil. Equil. (M) aqueous aqueous org.
K2
Carbon tetrachloride diluent 0-89
0-45
2.15 2.90 3.46 4.19 5-16 6.52 8.53 9.24 2-15 2-90
1-20 1.87 2-36 2.99 3.86 5.06 6.88 7.47 1.65 2.35
0"96 1.04 1.09 1.19 1-31 1.46 1.64 1.76 0.51 0.55
2.14 2.18 1.93 1.88 1.77 1.82 2.27 1.81 1 "93 1.90
0.45
0.25
3.46 4-19 5-16 6.52 2.15 2-90 3.46 4.19 5-16 6.52
2.88 3 "56 4 "47 5.75 1.88 2.58 3-12 3.83 4.76 6.08
0.59 0-63 0-70 0.77 0.28 0.31 0.34 0.37 0.40 0.43
1.75 1.78 1.72 1-89 1.71 1.68 1.81 1.82 1.77 1-76
4.19 5.16 6.52 2.15 2.90 3.46 4.19 5-16 6.52
3-55 4.44 5-71 1-87 2.58 3.12 3.82 4.74 6.06
0.66 0-72 0.79 0.29 0.32 0.35 0.38 0.41 0.45
2-18 2.21 2-36 2.12 2.05 2.14 2-08 2-12 2.63
Benzene diluent* 0.89
0.45
2-15 2.90 3-46 4.19 5-16 6.52 8.53 2.15 2.90 3.46
1.16 1.86 2.31 2-96 3.79 4.99 6.75 1.65 2.34 2.84
0.99 1.06 1.13 1.23 1.36 1.50 1-76 0.51 0.56 0.60
3.16 2-36 2-30 2-29 2-24 2.31 2.31 2.20 2.22 2.07
0.45 0.25
* Corrected for the solubility of the acid in pure benzene. mechanisms only: the first by the equilibrium (4) and the second by physical distribution without further complex formation. The first mechanism is chiefly operative when the amine is equilibrated with aqueous nitric acid solutions up to about 6-5 M, and the second one takes place virtually only at the highest acid concentrations. Consequently, the equilibrium conditions can be described simultaneously by equation (4) and by the distribution constant of associated nitric acid [HNO3]oM Kd -- [HNO3].
(6)
where [HNO~]o M stands for the amount of associated nitric acid in the organic phase which does not enter into compound formation with the amine and [HNO3] a stands for the amount of associated nitric acid in the equilibrium aqueous phase. Assuming ideally immiscible phases, we can write [R2HN]o = [R2HNHNO3]o+[R2HNHNOa.HNO3]o and [HNO3]o = [ R 2 H N H N O 3 ] o + 2 [ R 2 H N H N O a . H N O a ] c + [ H N O 3 ] o M
1428
A.S. IfamTmand I. T. PLATZl~m
Solving Equations (4) and (6) for IG we finally obtain [HNO3]o - [HNO3],
[R2HN]o(1-t-2[I-INO3]aK2) [HNO3].(I+[I-INO3].
(7)
2)
Calculations according to Equation (7) showed that Kd is proportional to the amine concentration in the diluent, and the distribution constant has a value of Kd ----0.10-4-0.02 for 1 M amine in both the diluents employed. The values computed obviously reflect the inaccuracy of 42 and the non-ideality of the organic phase at high acid loadings. Nevertheless, the facts that Kd is constant for a given amine concentration and proportional to the amine concentration confirm the mechanism postulated for the extraction of nitric acid by amine nitrate.
Acknowledgement~The authors wish to express their appreciation to Dr. ANNABECKfor her participation in rnnny valuable discussions concerned with this investigation.
DISTRIBUTION OF NITRIC ACID BETWEEN WATER AND A M B E R L I T E LA-1 L I Q U I D A N I O N E X C H A N G E R A. S. K~Tm* and I. T. PLATZNn Departmmt of Inoqpmic and Analyt/c~ ~ , The Hebrew University of Jaqumlem, J e ~ b m % Israel (Rec¢iwd 20 D~ember 1961) AbUrad--The solute-solvent inUnctkm betweemAmberlite LA-I, a Ions-chain secondary Amineextractant d|molved in carbon Wtrmchiceddeor bennme and 8qtmeus nitric acid wluti~n up to ~ 16 M has been inveetilsted by physical end chmni~ mmumrmn~ts. The flint reaction is an acid-base equilibrium leadin_a to the formation of amine nitrate. The ~ comtlmU were evahmted u El - 3.8 x 10s and $-4 x 10s in carbon tetrachlott~ and bemme m~m~vely. The meond and third.mactiom immured involve 8mine aitmte and meciated nitrk scid nmlmdm, leadin8 to the f e r n m ~ of a molemlar a d ~ a mmlmmtd/dmtifi~ as It4HNI~OrHNOs and to the ~ l a r t / t ~ of m e e i m d
nitria.c/d ~
the .qumm and o r s . ~ ptme.. The ~
~
p~h~
quotiemm Imve thz wdue8 JC2 -- 1.85 for carbon tetrachloride and K2 ffi 2.24 for end the ~ constant Ires the value F~d -- 0"I _+0~ for I M amim ~ in both dilmms. The eSeets of the formation of the amine nitrate ~ or of the retracted m o l e m ~ addit/on ~ er both, are not takm into ~ t . Bvidmm w u elmlmd that the q l r e ~ t i o n of the speciss in the orsank pham is larllely a peet-equililm~m
~mmon. IT has been observed(l-~ that the extraction of nitric acid by Ions-chain amines is not simply a case of an acid-baN reaction hadin8 to the formation of amine nitrate, since potentiometrk: neutralization curves of the acid contained in the equilibrium Orllanic phase give two endpoints. The amount of base consumed from the first to the second endpoint is equivalent to the amount of amine present, and the amount required to reach the first endpoint is a function of the conomttmtion of nitric acid in the aqueous phase at equih'brium. This ind/cates that the acid is bound in two di~emnt ways. * PmJmt addmm: l ) q m m t m t of O ~ t r y
and L~J~ratory for Nuclear 9cimee,
Msmschmem Insl/tute of Technelo~, ~ 39, Mass. I~) E. A.-h4lAsoN end V. C. A. VAVOmU~,Report A.~U-4~31 (1959); Report TI]D-ST20 (1960); V. C. A. VAVOmU~and E. A. MASON, Summary Rel~rt, Dept. of Nucleer. lVLLT.(I~0). {2) U. Bmt3~cx~ AEKE/R-2933 (1959).
(3) D. J. Ctawwma. and J. J. I.~WUNCS,7./mort. Ntr.t C/ram. 11, 69 (1959). (4) D. E. ~ snd C. F. CctattaN, Report OKNL-30$1 (1961). (s) (a) G. SCmONA,It~U~ Comm./~'~i. Jets., Report CNC-43 (1960); (b) G. SCatONAand M. Zn,mmmm,/b/d. CNBN-$1 (1961). (e) V. B. ~ , V. S. Sin/Ira', E. A. Nl~AmOZOtmVand K. A. F~,,tov, Z~. N~oqr. ~. S, 18.52 (1~0). (~) O. I. ZAH~OV-N~tcmOV and A. V. Ocmmq, 2~. Neoq,. ~k/m. 6, 1936 (1961). 1417
1418
A.S. KeR'r~s and I. T. PL~TZNmt
Several studies have been published in the last two years on the extraction of actinides¢2, 3, 4, 8), lanthanides,¢9) technetium and rhenium,C10, 11) some fission products¢l, 12, 13~ and heavy metals04~ from nitrate systems by amines. Some work has also ~ devoted to the r61e of nitric acid. In MASON'Slaborat0ry¢l~ the extraction of nitric acid in the range 2-10 M by a number of primary, secondaryland tertiary amines has been investigated. The nitric acid in the organic phase in excess of that equivalent to the amine concentration was found to be proportional to the concentration of the acid in the equilibrium aqueous phase for the tertiary amines, but a less simple relationship was observed for primary and secondary amines. CARSWELL and LAWRENCE,¢3)BERTOCCI(2) and COLF~N and HORNER¢4~ provide support for the linear relation with tertiary amines, and according to the work at Oak Ridge the linear relationship appears to hold for all amines, including primary and secondary ones.¢4) Thus, the concentration of extracted nitric acid can be related t0the concentrations of amine and aqueous acid by the expression (HNO3~xc~)ors/(Amine)ors(HNO3)~q ----ce where (HNO3e~m) is the organic nitric acid concentration above the stoichiometric ratio 1~1 with the amine. The available results fit this relatio~hip with values of c~between 0.16-0.18 for tertiary, 0,10-0.14 for secondary and 0-06-0.11 for primary amines. We have found that Amberlite LA-I, a secondary amine, follows the above relationship in the range between 3 and 12 M nitric acid in the initial aqueous solution, the constant u having a value of 0.13-0.14. Although this relationship is valid for a ~vide range of acid loadings and amine concentrations it cannot be interpreted as the mechanism of the reaction, and represents only an empirie~ relationship. SHEVCHENKOet al.¢6~ have studied the extraction of nitric acid by trioctylamine, and propose as the extraction mechanism complex formation between the amine nitrate and excess nitric acid according to the reaction (TOA'HNO3)o~-Ha+ ~(NO~-)a --* [(TOA'HNO3)'HNO3]o The calculated equilibrium constants have values of/C--~ 0.09 and 0.13 for the diluents carbon tetrachloride and xylene respectively, K ~howed satisfactory constancy in the range of 1-5 M aqueous nitric acid, but was measured for only one Cs~W, E. KEDER, J. C. SHEPPARDand A. S. WILSON,jr. Inorg. Nucl. Chem. 12, 327 (1960); J. C. SHEPP~,tu, Report HW-51958 (1957); A. S. WILSONand W. E. KEDmX,Report HW-62908 (1959); A. S. WILSON,Report HW-68207 (1961); W. E. KEDEn, J. L. RYAN and.A.S. WZLSON,Y. Inorg. Nuel. Chem. In press; F. IcmKAwA and S. URUNO,Bull. Chem. Soc. Japan, 33, 569 (1960); (~. J. HANDY, D. SC~XOILLand J. M, FLETCHER, J. Inorg. Nucl. Chem. 7, 257 (1958);'1"K. B. BROTHEaS,C. F. COLEMAN,D. J. CROUSE, J. V. DENmand J. G. MOORE,Y. Inorg. Nucl. Chem. 7, 85 (1958); H. G. PETnOW,Report TID~5357 (1960); F. R. BRuce, R. E. BLAN¢Oand J. C. BX~SEE,Report CF-58-11-91 (1959); B. WEAVERand D. E. HORNER,J. Chem. Engng. Data, 5, 260 (1960); M. DE TRENTrNL~Nand A. CHESNE,Report CEA-1426 (1960). (9) F. ICHIKAVA,Bull. Chem. Soc. Japan, 34, 183 (1961). CJ0~G. E. BOyD and G..V. ~LAt~SON,?. Phys. Chem fCJ, 988 (1960); A. S. KERTESand A. BECK,J. Chem. SbC. 1926 (1961); A. BECK. Ph.D. Thesis, The Hebrew University, Jerusalem (1961). ~1.~F. L. CULLmLAnnual Progress RepotL ORNL-2993 (1960). ¢t2~E. A. MASONand V. C. VAUGrmN],Report AECU-4239 (1959); Report AECU-4351 (1959). ¢13~V. B. SHEvcHEt,n~oand V. S. SMIVT, Radiokhimija, 3, 121 (1961); W. KNOCHand R. LXNDER,Z. Elektrochemie, 64, 1020(1960); W. KN]ocH,Z. Nat urforschung, 16a, 525 (1961). (t4) E. A. MASON]and R. E. Sg.AVDALL,Report TID-11196 (1960).
Distribution of nitric acid between water and ~ t e
LA-1 liquid anion exchanger 1419
concentration of amine (0.47 M). An unsuccessful attempt was made to: fit the results obtained in this study to ~ ¢ above reaction; the calculated, values o f K were not satisfactorily constant. In a,papar published while this report was being completed, ZAHAROV-NAgcISOV et al.~ have interpreted the equilibrium data in the system 0.25 M triheptylamine n!trateraqueous nitric acid (1.3-4 M) by the extraction of undissociated nitric acid, and postulated the existence of the species (THA.HNO3)HNO3 wittha formation constant K - 7.2=I=0.2. These authors also found that in carbon tetrachloride the organic solution of amine or amine nitrate is non-ideal; and constant values for K could not be obtained. The purpose of this investigation w ~ to ascertain whether the amine nitrate and nitric acid interact, and if so to what extent. In addition to a study of equilibrium data for the acid distribution, this investigation is devoted particularly to the physicochemical behaviour of the loaded organic layer, its viscosity, density and conductivity: Similar studies with phosphoric acld esters as extractants(XS, 16) have been successfully employed in. clarifying the heterogeneous equilibria in these solvent extraction processes. Amberllte LA-I was chosen as extractanL since it is one of the few orga~onitrogen compounds selected as suitable for process use.(l~ EXPERIMENTAL Materials. AmberHte LA.I* is a secondary amine, with a molecular weight of'4t6as determined in a potentiometric titration using per~loric acid.(ts) In the free base form,:a~ supplied bythe manufacturer, it is a viscous, amber-coloured liquid; very slightly soluble in aqueous solutions (10--20 p.p.m.). The diluents used were carbon tetracMoride and benzene, both of the highest quality, and they were used without further pm-illmti0n. Extraction procedure. Equal volmnes, urmtlly 5 ml, of the amine solution and the aqueous nitric acid solution were mixed in a t5 ml graduated centrifnge tubeandshaken mechankally for 15 vain. Certain experiments in which a different technique was employed are demm-ibedlater. After equilibration with 0-89M amine solutions when acid solutions of concentration greater than ,~ IOM were used, swe~i'na of the orllanic phase occurred irrespectivelyof the diluent. In these systems,prolonged centrifuilation at maximum speed wasnecessary tO wodu~ phase separation, in order.to enable volume .change to be measured. As a precaution against changes observed inthe organic phase which are described later, thesamples were treated one after another through all stages of the work, and in addition were, in so far as possible, protected from light. Work on each .sample took about 30 rain. As preliminary experiments had shown only a slight tomperature dependence, extractions were performed at room temperature (22-26°C) ~ where otherwise stated. Colour c h a n ~ , from light to dark brown, were otserved on ~ the amine solution into contact with the more concentrated nitric acid~ No attempt was made to investigate this phenomenon. Analytical procedure. Aliquots of the phases were diluted~ the o~ganic phase with an 8:5:1 methanol:acetone:water mixture, and the acid content determined potmtiometrically * We are indebted to Messrs. Rohm and Haas Company for kindly providing a free sample of the amine. ~ls~ A. S. K ~ T u , J. Inorg. NucL Chem. 14, 104 (1960); A. SI KIUtTnand V. K.~TES,Canadian J. Chem. 3S, 612 (1960); Y. AppL Chem. 10, 287 (1960); A. S. KEST~Sand M. HALPEaN, J. Inorg. Nucl. Chem. 16, 308 (1961). ~t6) A. S. K~T~S, A. BEctc and Y. HAnOUSHA,J. Inorg. Nud. Chem. 21, 108 (1961). (tT~C. F. COLEMAN,K. B. BROWN,J. G. Moon~ and D, J. Qtous~, Industr. Engng. Chem. 50, 1756 (1958). (ls~ CH. Bont~, Bull. Soc. C/u'm. France, 980, .1088 (1958); J. BIZOTand B. T~LLON, Bull. Soc. Chim. France, 122 (1959).
1420
A. S. KERTES and I. T. PLATZNBR
using a Beckman Zeromatic pH-meter. Two endpoints were observed in titrations of amine phases which were in contact with more concentrated nitric acid solutions. Turbidity, which does not interfere with the pH readings, appears just at the point when the excess acid (over the stoicheiometric ratio of 1: 1 amine to acid) is neutralized. Blank extractions with the diluents alone indicated that appreciable amounts of nitric acid are extracted into benzene, (5. 19) particularly at the higher acid concentrations (> 10M). The benzene blank based on the volume percentage of benzene in the organic solution, determined separately, was therefore subtracted from the total acid in the organic phase. The values for organic acid content given in the tables and graphs throughout this work are those for extraction by the amine alone. Physical chemical measurements. The density, viscosity and conductivity of the homogeneous organic layers were measured immediately after separation, according to the procedure previously described. US) Measurement of amine nitrate formation constant. The equilibrium constant for the acid-base reaction leading to the formation of amine nitrate was studied by converting the base into the nitrate, equilibrating the amine nitrate dissolved in the organic phase with water, and measuring the hydrogen ion activity in the aqueous phase. us. 20) A pure amine salt solution in benzene or carbon tetrachloride was prepared by converting the base to the nitrate form by shaking it with an aqueous nitric acid solution of slightly greater acidity, and removing excess acid by three washings with aqueous nitric acid solution of slightly lower acidity. The results referred to later in Fig. 5 show that the amine nitrate so obtained in the organic phase contained neither free amine base nor free nitric acid. A series of dilutions of this amine nitrate stock solution with the appropriate diluent was made down to 0,0015 M. Ten ml aliquots of the amine nitrate solutions were equilibrated with equal volumes of triply-distilled water, and the hydrogen ion activity of the aqueous equilibrium phases was determined by pH measurements using a Cambridge pH-meter to an accuracy of 0·01 pH units. Degree ofagitation and "post-equilibrium" changes of the loaded organic phases. ALLEN and McDoWELL(21l have observed that in amine systems the final distribution of the solute is not entirely independent of the way in which equilibrium is reached. Following the experimental procedure suggested by them, results were obtained for the extraction of nitric acid by both vigorous and gentle agitation at 27° to·l°C. The results presented in Table 1 show that the organic acid loading varies with contact time and type of agitation TABLE 1.-DISTRIBU110N OF NITRIC ACID AT VARIOUS INITIAL CONCENTRATIONS BETWEEN WATER AND 0,45 M AMINE IN CARBON TETRACHLORIDE AS A FUNCTION OF TIME AND TYPE OF AOITATION Slow equilibration Violent Initial equilib. aqueous NaOH consumed for 5 ml acid(M) of the organic phase 1 day 3 days 7 days 15 mins 8·75 mlO'259NNaOHforaminenitrate 8·75 8·90 9·05 9·80 mlO'259NNaOHfortotalacid 9·70 9·80 9·85 2-154 [HN03]0, total organic acid loading 0·504 0·508 0,511 0·508 8·75 8·75 8·90 9·40 mlO·259 NNaOH for amine nitrate 12·05 11·80 11·95 12·55 4·194 mlO·259 N NaOH for total acid 0·633 [HN03]0, total organic acid loading 0·613 0,620 0·652 8·75 8·75 9,00 9·45 mlO·259 NNaOH for amine nitrate IH5 14·85 14·95 15·70 6'518 mlO·259 NNaOH for total acid 0·784 [HN03]0, total organic acid loading 0·772 0·776 0·815 8·75 mlO'259NNaOHforaminenitrate 8·75 9·15 9·45 17·80 17·40 17,50 18·25 8'527 ml 0·259 N NaOH for total acid 0·923 [HN03]0, total organic acid loading 0,904 0·908 0·946 C. J. HARDY, B. F. GREENFIELD and D. SCAROILL, J. Chem. Soc. 90 (1961). (20) L. NEWMAN and P. KLOTZ, J. Phys. Chem. 65t 796 (1961). (21l K. A. ALLEN and W. J. McDoWELL, J. Phys. Chem. 64t 877 (1960).
(19)
Distribution of nitric acid between water and Amberlite LA-I liquid anion exchanser 1421 in a manner similar to that observed by ALL~ and McDOWeLL. Furthermore, as our results show, the amount of base necessary to neutralize the amine nitrate varies with the time of contact of the two phases. In addition, the loaded orlpmic phase, previously separated from the aqueous layer, and containing niti~c acid in conc~tratiom above.the I'1 stoicheiometric ratio, undergoes chanlles both in density and viscosity; Fig. 1 shows v i g ~ i t y meamrements
560
520
/"
u 480 -
/o
nf ° ,~
440
i
280
200 0
I
i
24
t
t
i
I
411 72 Aqelnq time,
]
[
96
I
, t20
hr
~ o . 1.--Pcat-equfl/brlum c h a n l ~ in the v t s c ~ t y of the loaded ~ (0.89 M amine equilibrated with 9-24 M nitric acid solution).
phase
carried out on allein8 solutions. A plot of time of emux in an Ostwald vi~zeity pipette is 8iveu rather than true v i ~ e i t y values, since diflimltim were encountered in density determinatiom using the usual Spreusel type pycnometer. The cause of change in the organic phase has not been det~...'ne~L Possible causes are re-ananllemeuts in the orsanic phase to yield different final amine-nitric acid species, which may be due either to some unknown post-equihlxium ~ t i o n effect,, s. = ) o r t o simpler chemical reactions such as oxidation or nm'uctk~ Tlm~ore, we do not feel that the anomalous solvent extraction equih'bria are emirely due to violence of allitation, as sulmested by A ~ J ~ and McDonaLd.c21) As additional evidence for this view, we point out that ueither the time nor the way of mixing was found to be critical in at least two different systans u~ing other eu~aes, viz. amin~ n i t r a ~ nitrate(z) and tertiary amine-sulphuric acid.(23) In any case, this post.equilibrium chanlle, shown in Table 1 and Fig. 1, raises considerable doubts concerning the applkal~lity of the equilibrium results for ailed solutions to an interpretation of extraction mechanism. It was t b e ~ o r e decided that rapid extraction equilibria would be simpler to interpret than those performed under chanlling conditions. Conaeqtle~tly, 15 min of violent shat-lna Was taken as an adequate and conveuieut sh__akin8time for valid equilibrium interlmmttion. (22) K. A. ~ , 3. Phys. Chem. 60, 239, 943 (1956); 62, 1119 (1958); 3. Amer. Chem. Soc. 80, 4133 (1958); W. J. McDowm.L and C. F. BAss, J. Phys. Chem. 67, 777 (1958). (z3) j. M. P. J. Vm~rmmq and J. A. A. ~
Tram. Faraday So¢. $7, 1527 (1961).
1422
A.S. K~TES and I. T. PLATZNER RESULTS A N D DISCUSSION
Physico-chemica/ properties of the organic layer. The density and viscosity o f the equilibrated organic phases as a function of the ratio [I-INO3]o/[R2HN]0 are shown in Figs. 2 and 3. The curves for carbon tetrachloride diluent (0.89 M amine) each 1'30
eje" •
_J E t~ "o >,°
1"28
/
1'26
/
e/
CCL4
7
I.:>4
S t t CsH6 0'9;
c c3 0-90
0'88
o.86
0
I
I
I
2
I
3
[ H N 0 3 ] o /rR2HN]o ratio in trle org pnose
FIG. 2.--Density of the equilibrium organic phase as a function of the nitric acid:amine ratio in the organic phase. Initial amine concentration 0.89 M.
/\\ ,...'
o.
I(~
~.
i'
~
_
CCL4
s 4
"t' C6H6
Z 0
i
I
[HN03] o /[RzHN]o
I
2
I
3
4
ratio in the or 9 phase
l:~o. 3.--Viscosity of the equilibrium organic phase as a function of the nitric
acid:amine ratio in the organic phase. Initial amine concentration 0.89 M.
Distribution of nitric acid between water and Amberlite LA-I liquid anion exchanger 1423 show two intersections between straight lines of different slopes. The first intersection corresponds to the amount of acid necessary to neutralize the amine in the organic phase, producing the amine nitrate, R2HNHNO3. The second intersection is at a ratio of 2:1, indicating that the amine nitrate has interacted with more nitric acid to give a molecular addition compound which may be tentatively identified as R2HNHNO3.HNO3. The corresponding curves for benzene diluent are only plotted to a ratio of about two, since appropriate corrections could not be made for the benzene blank. Measurements were carried out also with 0.45 and 0.25 M amine solutions in carbon tetrachloride, but the high diluent content in these solutions caused the viscosity and density values to level off.(7) A change in reaction mechanism is also indicated by the conductivity curves in Fig. 4. Both the amine and the amine
•0
9
C¢L41
~ 7g u "o
g
-
._u
-
• C~%,"
/
o. se SQ 0
t 2 3 [HNOB]e/[IR~H.]o rotio in the orq phose
FIG. 4.---Specific conductance of the equilibrium organic phase as a function of the nitric acid:amine ratio in the organic phase. Initial amine concentration 0.89 M. nitrate solutions showed no detectable conductivity with the equipment available, and the onset of conductivity is believed to be due to formation of a conducting species, such as the addition ~'ompound tentatively defined above. This suggests that the extent of ion-pair dissociation m the organic phase is likely to be important. A rough estimate of the acid dissociation as calculated by WALDEN'Srule (//2/60), shows that it does not exceed five per cent at the highest acid and amine concentrations. Partial ionization of tributyl phosphate-nitric acid compounds at the highest nitric acid concentrations, found from infra-red absorption data by TUCK,(Z4) may be indicative of some similarity between the two systems. Formation of amine nitrate. The neutralization reaction between amine and nitric acid up to a [HNO3]o/[R2HN]0 ratio of unity may be represented as (R2HN)o+ H + + (NO3-)a ~ (R2HNHNO3)o (1) where R2HN stands for the amine base and R2HNHNO 3 for the amine nitrate, and the subscripts --a-- and --0-- indicate the aqueous and organic phases respectively. ~2~)D. G. TUCK,J. Chem. Soc. 2783 (1958).
1424
A.S. K~TeS and I. T. PLATZ~
The thermodynamic equilibrium constant for this reaction is: 7AN.[R2HNHNO3]o (2) KI 7A:[R2HNJo{H+}a{NO~}, ?AN is the activity coefficient of the amine nitrate and YAis that of the amine base both in the organic phase; their ratio can be taken as unity, since they refer to low concentrations and to species that are relatively alike with respect to the medium. Under the experimental conditions described, {H+} a, the hydrogen ion activity, can be taken as equal to the nitrate ion activity {NO;}a. and is also numerically equal to the concentration of free amino base [R2HN]o. Assuming ideal immiscibility of the phases, [R2HNHNO3]o equals the initial amine nitrate concentration less {H+},. Thus, transforming (2) into a logarithmic form, we obtain --log ([R2HNHNO3]~--{H+}a) ---- 3pH+pKI (3) A plot of the leA-hand expression vs. the pH readings should give a straight line with a slope of three and an intercept equal to pK1. Fig. 5 shows such plots, with intercepts K1 = 3.8)< 10s and K 1 = 5.4× 10s for carbon totrachlorido and benzene media
respectively.
CCt4 -3"0
-2.5
o z
-2"0
Z "I" N
.~. -,.5
-I.0
[
,.s
n 2.0
I
I
J
2.4
2.8
3.2
pH
l~O. 5.bPlot of Equation (3) for the determination of/{'i. As seen from the plots, the method of calculation obviously fails when the amino nitrate concentration exceeds 0,05 M. Since deviations from linearity increase gradually with the amino salt concentration in the organic phase, it appears that the limitations are imposed rather by the inability to determine activity coefficients for the organic solutes, than by a chango in the mechanism.(20)
Distribution of nitric acid between water and Amberlite LA-1 liquid anion exchanser 1425
Extraction of acid by amine nitrate. The extent of the extraction of nitric acid by amine nitrate as a function of its equilibrium aqueous concentration is shown in Figs. 6 and 7 for carbon tetrachloride and benzene as diluents, respectively.
/ e/
...J"
~
. ...//'J" o ~o...
"
0"89 M
/
./07/°/~,. j,~,,=o ,,,~.,~
J 25M
0
L.
I
I
4
I
I
I
8
I
12
J
[H NO3] ° M
Fzo. 6.--Eqnilibrium distribution of nitric acid between water and carbon tetrachloride solution of the amine. 3
I
Os
I
1
at-
/
1
0
/×
/
:o.., 1 2' L
/
o°"
I
/ /
• Se'e
4
8 [H NO~]o,
12 M
I'~. 7.--Equifibriurn distribution of nitric acid between water and benzene solution of the amine.
1426
A. S. KERTESand I. T. PLATZNER
It was found that 18 M nitric acid is completely miscible with undiluted amine nitrate; this might be taken as evidence for the extraction of associated nitric acid molecules rather than ionized nitric acid. Similar observations and interpretations were made by SHULER(25)and by COLLOPYand co-workers(26) in nitric acid extraction by tributyl phosphate. The amine nitrate forms a molecular addition compound with associated nitric acid and the amine nitrate-nitric acid species formed may be an amine analogue of the tributyl phosphate-nitric acid species TBP.HNO3 or even TBP.(HNO3)2. The comparable shapes of the isotherms for extraction of nitric acid by TBP and by amine nitrate emphasise this similarity (compare Figs. 6 and 7 with Fig. 2 of reference 25). The above reasoning implies the existence of the equilibrium (R2HN HNO3)o + (HNOa)a ~ (R2HNHNO3.HNO3)o (4) with the equilibrium concentration product quotient K2 = [R2HNHNO3.HNO3]o (5) [R2HNHNO3]o[HNO3Ja where [HNO3]a is the concentration of undissociated, or ion=paired, nitric acid in the equilibrium aqueous phase, which can be calculated from the data of KRAWETZ.(27) Knowing both the total organic acid concentration and the initial concentration of amine in the organic phase, we may write [R2HNHNO3.HNO3]o = [HNO3Jo--[R2HN]o and [R2HNHNO3]o = [R2HN]o--[R2HNHNO3.HNO3]o -----2[R2HN]o--[HNO3]o Substituting these values in Equation (5), values for the quotient K2 have been calculated (Table 2). The constancy of the values of K2c ° ' = 1.85 and K2c'n' = 2.24 is considered remarkable in view of the fact that both ionic strength and the media change markedly as the aqueous nitric acid content increases, and considering that the activity coefficients of the solutes were taken as unity. The mass-action equilibrium (4) obviously fails to explain the extraction mechanism when the ratio [HNOa]o/[R2HN]o exceeds two. There are two possible ways in which further nitric acid, exceeding the above ratio of two, can be extracted into the organic phase. The first is by partition of associated nitric acid between its aqueous solution and the amine nitrate-nitric acid complex, and the second by a further complex formation reaction leading to the formation of molecular addition compounds of the type R2HNHNO3.(HNO3)~. with x ,> 2. It is not possible at present to decide which process is responsible for the phenomenon. This fact is not surprising, in view of the widely contradicting interpretations of solute-solvent interaction in the more thoroughly studied system tributyl phosphate-nitric acid (reference 25, and references therein). However, since the physico-chemical measurements on the loaded organic phase (Figs. 2 and 3) do not indicate any further compound formation beyond R2HNHNOa.HNO3, we believe that nitric acid is extracted by amine nitrate by two (25) W. E. SHULER,Report DP-513 (1960). (26) T. J. COLLOPYand J. F. BLUM,J. Phys. Chem. 64, 1324 (1960); T. J. COLLOPYand J. H. CAVENDISH,3". Phys. Chem. 64, 1328 (1960). (27) W. J. HAMER(Editor), The Structure of Electrolytic Solutions, p. 43. J. Wiley, New York (I 959).
Distribution of nitric acid between water and Amberlite LA-1 liquid anion exchanger 1427 TABLE 2.--CALCULATEDEQUILIBRIUMCONCENTRATIONPRODUCTQUOTIENTEOR THE REACTIONOF ASSOCIATEDNITRICACIDWITH AMINENITRATEUSINGEQUATION(5) HNO3 (M)
HNO3 (M)
[R2HN]0 Initial Equil. Equil. (M) aqueous aqueous org.
/(2
[R2HN]0 Initial Equil. Equil. (M) aqueous aqueous org.
K2
Carbon tetrachloride diluent 0-89
0-45
2.15 2.90 3.46 4.19 5-16 6.52 8.53 9.24 2-15 2-90
1-20 1.87 2-36 2.99 3.86 5.06 6.88 7.47 1.65 2.35
0"96 1.04 1.09 1.19 1-31 1.46 1.64 1.76 0.51 0.55
2.14 2.18 1.93 1.88 1.77 1.82 2.27 1.81 1 "93 1.90
0.45
0.25
3.46 4-19 5-16 6.52 2.15 2-90 3.46 4.19 5-16 6.52
2.88 3 "56 4 "47 5.75 1.88 2.58 3-12 3.83 4.76 6.08
0.59 0-63 0-70 0.77 0.28 0.31 0.34 0.37 0.40 0.43
1.75 1.78 1.72 1-89 1.71 1.68 1.81 1.82 1.77 1-76
4.19 5.16 6.52 2.15 2.90 3.46 4.19 5-16 6.52
3-55 4.44 5-71 1-87 2.58 3.12 3.82 4.74 6.06
0.66 0-72 0.79 0.29 0.32 0.35 0.38 0.41 0.45
2-18 2.21 2-36 2.12 2.05 2.14 2-08 2-12 2.63
Benzene diluent* 0.89
0.45
2-15 2.90 3-46 4.19 5-16 6.52 8.53 2.15 2.90 3.46
1.16 1.86 2.31 2-96 3.79 4.99 6.75 1.65 2.34 2.84
0.99 1.06 1.13 1.23 1.36 1.50 1-76 0.51 0.56 0.60
3.16 2-36 2-30 2-29 2-24 2.31 2.31 2.20 2.22 2.07
0.45 0.25
* Corrected for the solubility of the acid in pure benzene. mechanisms only: the first by the equilibrium (4) and the second by physical distribution without further complex formation. The first mechanism is chiefly operative when the amine is equilibrated with aqueous nitric acid solutions up to about 6-5 M, and the second one takes place virtually only at the highest acid concentrations. Consequently, the equilibrium conditions can be described simultaneously by equation (4) and by the distribution constant of associated nitric acid [HNO3]oM Kd -- [HNO3].
(6)
where [HNO~]o M stands for the amount of associated nitric acid in the organic phase which does not enter into compound formation with the amine and [HNO3] a stands for the amount of associated nitric acid in the equilibrium aqueous phase. Assuming ideally immiscible phases, we can write [R2HN]o = [R2HNHNO3]o+[R2HNHNOa.HNO3]o and [HNO3]o = [ R 2 H N H N O 3 ] o + 2 [ R 2 H N H N O a . H N O a ] c + [ H N O 3 ] o M
1428
A.S. IfamTmand I. T. PLATZl~m
Solving Equations (4) and (6) for IG we finally obtain [HNO3]o - [HNO3],
[R2HN]o(1-t-2[I-INO3]aK2) [HNO3].(I+[I-INO3].
(7)
2)
Calculations according to Equation (7) showed that Kd is proportional to the amine concentration in the diluent, and the distribution constant has a value of Kd ----0.10-4-0.02 for 1 M amine in both the diluents employed. The values computed obviously reflect the inaccuracy of 42 and the non-ideality of the organic phase at high acid loadings. Nevertheless, the facts that Kd is constant for a given amine concentration and proportional to the amine concentration confirm the mechanism postulated for the extraction of nitric acid by amine nitrate.
Acknowledgement~The authors wish to express their appreciation to Dr. ANNABECKfor her participation in rnnny valuable discussions concerned with this investigation.