216
J . hf. P. J. l
rAND J. A. ~ A.
ICETEL.4AR ~
~
~Vol. 66
THE DISTRIBUTION OF SULFURIC ACID BETWEEN WATER ASD KEROSENE SOLUTIONS OF TRI-n-OCTYLAMINE AND TRI-n-HEXYLAlliIINE' B Y J. M. P. J. V E R S T E G E N 2 AND J. h.h.K E T E L A A R 3 Laboratory .for General and Inorganic Chentisti-y of the University of Amsterdam, Amsterdam, Holland Reeeinsd Mat, $l$1961
In this paper Q description is given of the distribution of sulfuric acid betwcen water and a lterosene solution with tri-noctylamine or tri-n-hexylamine. Assuming ideality in the organic phaw no evidence could he obtained that the law of mass action underlies the fundamental processes. Assuming aggregation of the amine salt to a micelle of constant activity a qualitative agreemmt is found and the empirical formula K A = uIirso,[arninejnseems to be of general interest, the power n changing a t inflection points in the titration curves of the weak base anion-exchanger. In the range of acidities where no free amine exists a treatment based on mixed crystal equilibrium betweeen amine sulfate and aminc bisulfate leads to rewith X as the sulta which are much the same for both amines. An empirical formula KR = ( l / a ~ ~ o , ) ( P / (-l mole fraction of the amine bisulfate describes the system.
x))"
Introduction I n a previous article4 it was shown that the distribution of sulfuric acid between water and benzene solutions of tri-n-octylamine (TOA) or tri-n-hexylamine (THA) could be based on the formation of the sulfates (TOAH)2S04or (TI~AH)T
SOr. If it was assumed that the concentrations in the organic phases mere equal to the activities t'he law of mass action gave a satisfactory description of the results, until a certain amine sulfate concentration was reached and deviations occurred. The results were in reasonable agreement with those given by Allen6 who explained the deviations as being due to aggregation of the amine salt to micelles.6d6 I n the present study it is shown that the processes underlying the distribution show remarkable differences when the benzene is replaced by kerosene.
Experimental
The commercial .products TOA aad THA are distilled $t 1-2 mm. The boiling points are !82-186' and 118-121 , respectively. During the distillation of TOA the forerun solidifies in the collector. This is probably di-n-octylamine. The molecular weight of the main raction of TOA (boiling range < 2,) is determined by means of otentiometric titration with standard perchloric acid in a &enzene-wateralcohol medium. It is found to be 351.6. Assuming that the impurity is di-n-octylamine, this value is in agreement with n percentage of 98.6yo T0.4. The molecular weight of the main fraction of THA is determined as being 278.5. No attempts were made to investigate if the iso-derivatives were present. In all distribution measurements once-distilled amines were used. Samples of both THA and TOA were exposed to the laborat o g atmosphere for three days. No change in might could be observed and otentiometric titration with standard perchloric acid in a genzene-alcohol-water medium before and after exposure did not show any significant difference. Ten ml. of a 0.098 M amine solution in kerosene,7 modified
with 4 vol. o/o n-octyl alcohol to prevent third phase formation, were brought into contact with the same volume of a ueous phases containing known amounts of sulfuric acid
(&I.
We did not find the anomalous behavior which was reported by Allen and McDowells for the case of uranium extraction. To the aqueous phases, which were brought into contact with the kerosene solution of T H 1,0.33 M Nu2S04 was added to prevent dissolution of the formcd (TIIA€I)r SO, in the aqueous phase. Equilibrium was achicved by means of rcntrifugal stirring for 1 minute. Thirtv minutes after equilibration the layers were separated and the acid content of the water layer (Cf) was determined by means of potentiometric titration using either 0.1 $1 or 0.02 M NaOTI. To verify the material balance in some cases the organic loading C, was determined by stripping the acid from the kerosene phases with sodium hydrovidc and titrating back the excess alkali. It was found that Ca =
Results If it is assumed that the reaction bctwccn amine and sulfuric acid taking place a t the interface is analogous to the reaction between TT2S04 and KHt, we map write and 2THX
+ H&04
(TIIAIT)2SO,
(la)
At higher widities the equilibria could be rcprcsented as (TOAH)&O,
(1) The work, described in this article. forma part of the program of RCN (Reactor Centrum Nederland), The Hague. Thn experimental
results have been obtained in the Laboratory for General and I n o r ganio Chemistry of the University of Amsterdam. (2) Institutt for Atomenergi, Postboks 175, Zillestriim, Norway. (3) Laboratory for Electrocliemistry, University of Arristerdam. (4) J. PI. P. J. Verstegen and J. A. A. Ketelaar, Trana. Faradaw Soc.. 57, 1627 (1901). ( 5 ) IC. A . Allen. J . Phys. Chcrn., 60,239 (1956). (0) K . A. Allen, ibid.. 60, 943 (1956). (7) Tho kerosene used in our experiments was furnished b y the Anisterdamsche Chininefahriek. Some physical and chernical properties are: s . ~ ' . (at 2 4 O ) = 0.778, refractory index n% 1.431, boiling range 195-240°, aromates I 2 3 4 6 8 10-2 10-1
10-3
4 i im o W . Fig. 2.-Free
L L , ,
1
2 34579
amine concentration [TOA] us.
+&&
1 -
s8
910-1 8 7 1 6 5 t 4 t
4.5 f 3.5 3 '* 2.5 2 ' 1.5 -
4.5 I 3.5
4
5
' 9
2 3 2.5 2 1.5 10-2
8
1
3.5 2.5
p 2
1.5
1
e
0.05 0.07 Fig. 3.--X2/( 1 -
10-2
_
1
.
/
.
L-L..
ii
0.01 0.15 0.2 0.3 moW. X ) us. .2/aH,so, for TOA.
-
1-
2
lo-'
Fig. 2a.-Free
3 456789 2 3 456789 10-3 10-2 ;/aaiso, mole/l. amine concentration [THAI us.
3.5
$ 2.51 1.:1
9
It
corresponding to an empirical constant Kz
= ~ H Z S O[TOAI ,
(7)
with the numerical value K 2 = 75 X
(mole/l.)12
In the range of high acidities the slope is found to be equal to - 3/4, according to an empirical constant K, =
~ ~ 2 8 [TOA]' 0 ,
=
24 X 10-14 (mole/l.)'
(8)
In Fig. 2a the slope a t low acidities is difficult to estimate and only can be given as
with 24 < n < 36. At high aciditirs it is f o l d that the slope is equal to -1,12, or Kp
=
~H&O~[THA]"
(8a)
with K ~ .= 15 x 10-9 (m01e/1.)9
The changes in n in the formula K A = U H ~ S O , . [amine]" take place a t the activity a t which the neutralization step in the titration curves (Fig. 1,la) starts. The readers attention might be drawn to the fact that the empirical formulas given in 7, 7a, 8 and 8a show a qualitative agreement with the results obtained from the distribution of €1804 be-
0 0.040 0.045 0.050 0.055 0.060 mW1. Fig. 3a.-X*/(l X ) us. i/Gfor TFIA.
-
tween water and a 0.100 M solution of TOA in benzene a t 65°.4 The deviations from the Figs. 2 and 2a must be ascribed to the first formation of bisulfate. These deviations start a t [THAI < 0.030 M and a t [TOA] < 0.020 M , in agreement with the greater basicity of TIIA, which was found p r e v i ~ u s l y . ~In the range of acidities where the equilibria mentioned under 2 and 2a are predominant another treatment must be chosen.
Feb., 1962
HEATOF VAI~OR~ZATION AND HEATOF FUSION OF I'J!XRIC CHLORIDE
Allen6 already pointed out that it might be assumed that the normal amine sulfate and the amine bisulfate form a comp1ctc:ly miscible ideal solution, which is analogous to the ideal solid solution of the components of a solid ion exchanger. I n that case the activities of the resin species can be represented by their mole fraction X . ' The equilibrium constants K and K, of the equilibria 2 and 2a can be given as
1
K~ = 30.5 (m01~/1.)-3
a t high acidities we find a slope X2
d log -1-x = I d log U I I , R O ~
Ks
with
=
__ 1 a1i1s04
(TOhH)BSO4 2( TOAH)~SOI
= (TOAH)HSO4
+
X(TOAH)BO~ = 1 - X
(11)
and the same relations exist for THA. Our K , can be written now as
(-)I x* -x
a
with the numerical value K~
(10)
and
x*
and
K , = X 2 ( ~ ~ ~ ~ ) ~ ~(9)~ C corresponding to UH~SO,X(TOAH)~~O
X(ToAH)nso4=
(i+
K4 =
219
=
11.4 (moie/1.)-3
In the THA case (Fig. 3a) only the slope 6 is found corresponding to
with the numerical value Kla = 5570 (mole/l.)-a
Plotting u'/*H~so, us. X*/(l - X) for the TOA and the TI-IA system, respectively, we find the Figs. 3 and 3a. There is again a qualitative agreement, but the form of the TOA Fig. 3 is somewhat more complicated, in agreement with the more complicated titration curve (Fig. 1). In Fig. 3 the part a t low acidities corresponds to a slope d log-
X2
1-x = 6
d log U1/*"2S0r
in agreement with an empirical constant
Discussion From the experiments and the results described above no quantitative conclusions can be drawn. The relations obtained are only empirical and will certainly have no general thermodynamic validity. The only conclusion which might be drawn is that the transition points in Fig. 2, 2a and 3 correspond to inflection points in the titration curves 1 and la. A surprising point of interest is furthermore t,he great difference in the processes underlying the distribution when benzene is replaced by kerosene as the solvent.
THE HEAT OF VAPORIZATION AND TKE HEAT OF FUSION OF FERRIC CHLORIDE BY CHARLESM. COOK,JR. PigmPnts Department, E . I . du Pont de Nemours & Co., Inc., Wilmingh, Dehurare Received June 9 , 1961
Pressures of Fe&&(g) above solutions of FeC12 in FesCls(1) were measured in the range 300-470". These data indicate vap.) = 14.6 i that above liquid ferric chloride log PFoCle(mm.) = 48.57 - 12.55 log T - 6373/T, where AHa~.~(FenC16 0.5 kcal./mole FerCle, ASsm.T(FenCle vnp.) = 24.8 f 1.0 e.u./mole FerC16. The boiling point is estimated to be 315'. The heat of fusion was measured by a drop calorimeter and found to be AHSm.,(FeCl$fusion) = 9.0 i 0.4 kcal./mole FeCla.
Introduction The currently acceptedla!b value of the heat of vaporization, 12.04 kcal./mole Fe2C16, at the accepted boiling point of liquid ferric chloride, 3 1 9 O , derives from measurements by Stirnemann2 of the total pressure within a closed bulb containing ferric chloride a t temperatures up to 493O. This thermodynamic value was calculated without correcting the observed total pressures for the sig-
nificant partial pressure of Cls present along with the FesC16(g) in the vapor space above the ferric chloride.8 The magnitude of the chlorine pressure correction to Stirnemann's data is discussed below, and the heat of vaporization of FezCls(l) is recalculated. Experimental
(1) (P) "Selected Values of Chemical Thermodynamic Properties," Circular 500,National Bureau of Standards (1952); (b) 0. Kubaschewski and E. Evans, "iMetallurgica1 Thermochemistry," 2nd Ed., John Wiley and Sons, New York, N. Y.,1956. (2) E. Stirnemann, Neues Jahrb. Mineral. Geol. u. PaZdontoZ.. 62A, 334 (1925).
(3) W. Kangro and E. Petersen, Z . anorg. Chem., 261, 157 (1950), corrected Stirnemann's presaures for P F ~ C and I ~for P C I ~The latter calculation, however, contained the assumption that aFacil = 1, which is invalid for Stirnemann's experimenta above the FeCkFeCla eutectic temperature because of the solubility of FeClz in molten FezCls. See H. Schllfer, Z. anoru. u. a2Zeem. Chem., 266, 269 (1951).
Vapor Pressure.-Ferric chloridoferrous chloride mixtures were contained in a ca. 13-cc. cylindrical Pyrex sample
