SOLUBILITY OF METALSULFATES IN WATER
Jan. 20, 1957
267
+ +
diffraction pattern showed that It k I = 2n, the requirement for a body centered lattice. This is consistent with the space group possibilities for KzNiFd.8 By analogy with this structure the z parameters for oxygen and strontium were assigned the values zo = 0.151 and zsr = 0.347. The calculated relative intensities are compared with the observed intensities in Table I. No correction was made for absorption. The agreement is good and appears to establish the structure of SrzIr04 as the K2NiF4 type.
TABLE I OBSERVEDAND CALCULATED INTERPLANAR SPACINGS AND INTENSITIES (FORWARD REFLECTIONS) hkl
d/n(obs.)
d/n(calcd.)
002 101 004 103 110 112 105 006 200 202 211 116 204 107 213 215 206 118 220 222 109 301 217 224 303 208 310 312 226 305 1.0,11
6.48 3.72 3.23 2.88 2.74 2.53 2.143
6.42 3.72 3.23 2.89 2.75 2.53 2.148 2.150 1.944 1.861 1.722 1.692 1.665 1.667 1.612 1.442 1.443 1.392 1.374 1.344 1.346 1.289 1.265 1.268 1.241 1.242 1.229 1.208 1.158 1.158 1.124
1.942 1.859 1.720 1.689 1.667 1.612 1.441 1,392 1.374 1.346 1.288 1.264 1.242 1.229 1.207 1,158 1.124
I(obsd.)
m+ mw-
41.8 22.7 1.2
S+
100.0
S-
55.5
4.6
W
m+
13.1 16.6 27.6 5.7 4.7 22.2 0.7 0.1 51.5 6.9 19.0 3.3 3.1 0.2 8.5
m w W
m' W S-
m
w W W+
0.9
W'
0.4 0.1 8.4 5.2 10.9 3.9 8.7 1.7 3.5
WS
w
IO *
,\b2-?
I(ca1cd.)
++
w++ WW++
W
Discussion The formation of S r d r 0 4 as a pure phase by the reaction of mixtures containing excess strontium [CONTRIBUTION FROM T H E O A K
3 000
io00
4000
5000
2000 VOLUME OF FE'* ADDED.
Fig. 1.-Typical
photometric titration curve for (IrClc)-4,
indicates that all of the excess strontium oxide is volatilized a t 1200" since no change in lattice constant was observed. With less than stoichiometric quantity of strontium, the excess iridium is liberated and the pattern of SrzIrO* again shows no change in spacing. The compound SrzIrOd is therefore assumed to be stoichiometric. Other transition metals which have been shown to give AzB04 phases are titanium, manganese and molybdenum.8 Each of these elements also gives a perovskite-type oxide with strontium but all attempts to prepare a phase of this type in the strontium-iridium-oxygen system have failed. Lower temperatures and reactions in which only stoichiometric amounts of oxygen were allowed to react have been tried without success. This cannot be explained on the basis of relative ionic radii since the Goldschmidt radius of I r + 4is 0.66, which lies close to the range covered by the other elements (Ti+4= 0.64, Mn+4 = 0.52, M o + ~= 0.62). The tolerance factor for the perovskite structure t = [(RA R o ) / d (RB Ro)]for iridium is 0.95, well within the common range. Acknowledgment.-The authors wish to thank Dr. David Hume for his suggestion of the photometric method for iridium and Mr. Donald Klein for his advice concerning the experimental procedure.
+
+
STORRS, CONN.
RIDGENATIOiVAL LABORATORY, CHEMISTRY DIVISION]
The Solubility of Several Metal Sulfates at High Temperature and Pressure in Water and in Aqueous Uranyl Sulfate Solution1 BY ERNEST V.
JONES,
M. H. LIETZKEAND WILLIAHL. MARSHALL RECEIVED AUGUST 18, 1956
The solubility of several metal sulfates a t high temperature and pressure has been determined in HzO and in aqueous U02S04. I n most of the systems investigated a large increase in solubility of the metal sulfate is observed a t a given temperature as a function of increasing UOZSO~ concentration. This type of behavior points to a considerable degree of complexing in the saturated solutions. Several heats of solution have been determined and a possible relationship between these values and the amount of complexing is suggested.
Introduction Two general methods are available for the determination of the solubility of salts a t elevated tem(1) T h i s d o c u m e n t is based on work performed for t h e U.S. Atomic Energy Commission a t t h e Oak Ridge N a t i o n a l L a b o r a t o r y .
perature. I n the analytical method saturated solutions are Prepared and separated from the solid Phase a t the equilibrium temperature, and In the synthetic method known ratios of salt and water are heated and the temperature
ERNESTv. JONES, M. H.LIETZKE AND
2GS
is observed a t which the disappearance or appearance of a solid phase occurs. I n the present investigation the synthetic method was used to determine the solubilities of the sulfates of lanthanum, yttrium, cadmium, cesium and silver a t high temperatures in water and in aqueous U02S04. I n addition, the solubility of zinc sulfate in water alone a t high temperatures was investigated while the solubility of Bas04 in aqueous uo&04a t 250' was studied by the analytical method. Since uranyl sulfate is known to complex the sulfate ion,2it seemed of interest to determine how much the solubilities of these salts increase in UO2SO4 solutions. The solutions used were 0.1272, 1.145 xnd 1.348 m in U02S04. Experimental Methods \Vith the exception of the determinations below 100" and the filter technique for BaSO4, all solubility measurements were made by the silica tube method described by Etard3 ant1 A. B e ~ x a t h . ~The , ~ silica tube method depends upon the visual observation of the appearance or disappearance of crystals in synthetic "mixtures" of known concentration. In the case of CSZSOC and Ag,SOa, both of which showed a positive temperature coefficient of solubility, the water and crystals were weighed directly into the silica tubes for each determination. In all other cases the solutions were prepared by dilution of a more concentrated stock solution. In place of the 2 to 3 mm. i.d. heavy wall tubes used in most of the previous work, both 4 and 6 mm. i.d., conventional wall silica tubes were employed in the present experiinents. These tubes did not permit as high temperatures to be reached, but made easier the visual observation of the appearance or disappearance of crystals. The tubes were heated in aluminum block furnaces which were mounted on motor driven rockers so that thorough mixing of the solution and crystals could be accomplished. T h e temperature, which was controlled t o about &lo, was measured by an irori-constantan thermocouple placed directly above the silica tube in the furnace and recorded on a calibrated Brown recording potentiometer. A temperature correction curve was obtained by the use of two thermocouples. One thermocouple was placed within a silica tube in the position t o be occupied by a tube containing the synthetic mixture. The other thermocouple was placed in its true position to be occupied during a run. which was used as a correction factor, was inThe A ( t " , significant a t 100" but increased t o approximately -3" a t 300O. The lowest solubility measured was limited in each case by the ability t o detect the first appearance of crystals in the tube. In some cases the crystals were visible in much more dilute solutions than in the case of other salts. For example, the measurement of the solubility of La*(S04)3 in UOZSO, solution could be extended t o 0.047, LaZ(SO4)3while the cadmium and zinc sulfate limits were much higher (2 t o 37,). In all cases the measurements were extended to as dilute solutions as possible using the experimental methods available. The temperatures a t which crystals dissolved or appeared in the tubes could be reproduced to about & l o . The solubility of BaS04a t 250' was determined by the use of a two-cell filtration-type apparatus similar to that of Booth and Bidwell.6 The solutions and excess radioactive Ba1"S04 w-ere equilibrated a t 250" in one of the two compartments, which were separated by a sintered platinum filter obtained from the American Platinum Works. After equilibrium had been established the apparatus was inverted and the empty end cooled somewhat below 250'. The saturated solution was forced through the filter by means of the vapor pressure differential. The entire apparatus was then cooled to room temperature, the solution sample removed and the barium concentration in the solution determined by
e.),
('2) R . 11. U e t t s and R. K . hlickels, J . Chenz. Suc., Suppl., 2 , 286 ( I $119). (3) Etard, Ann. chiin., [71 603, 2 (1894). (-1) A . Benrath, 2. anorg. ailgem. Chem., 231, 285 (1937). ( . 5 ) A. Benrath, ibid., 247, 147 (1441). ((;) K. S. Booth and R. W. Bidwell, THISJ O U R N A L , 74, 2567 (1950).
%'ILLIA>I
L.
MARSIIALL
Vol. :!I
radioactive analysis. Equilibrium was checked by making a series of runs as a function of time. It was found that equilibrium was attained within 24 hours.
Experimental Results Lanthanum Sulfate.-The high temperature solubility of La2(S04)3 in HzO, in 0.1271 m, and in 1.348 m uo~so4solution is given in Table I. An experimental point is given a t which temperature two liquid phases appeared. The formation of two liquid phases a t elevated temperatures is characteristic of the system U02S04-Hz0.' I n all three cases the solubility of LaZ(S04)3 showed a negative temperature coefficient. Since the Laz(S04)~solutions all showed a strong tendency to supersaturate it was necessary to heat them 50 to 70" above the solubility equilibrium temperature before crystals appeared. The temperatures given represent the points a t which the last crystals redissolved upon slow cooling. The L a ~ ( S 0 4 )crystals ~ did not adhere to the sides of the tubes. Several solutions of La2(SO4)3 in water and in 0.1272 m U0&304 were heated to 300" to check on the possible existence of double solubility points a t a particular concentration. No such points were observed. However, a metastable solubility point was observed a t 149' for 3.87 mt. yo La2(S04) in 1.348 m U02SOd solution. TABLE I THESOLUBILITY OF LASTHANUM SULFATE IN WATERA N D IN URANYL SCLFATE SOLUTIONS LadSOds,
t
("C.)
108 113 115 117
97 105 120 145 169 199 215
Wt.
yo
I ("C.)
1,az(SOd~, w t .
7%
In water In 1.348 m UOzSO4 4.39 0.61 7s 3.97 .46 91 and 149 3.80 .37 148 120 3.75 .31 I n 0.1272 nz 150.5 3.70 UOzSOa 151 3.60 2.04 152 3.46 1.74 154 3.23 0.87 160 2.69 .29 168 2.30 .15 178 1.85 .083 202 1.22 042 222 0.88 0.57 Two liquid layers at 291'
To check on the stability of La~(S04)zsolutions a t high temperature a 0.31% La2(S04)3solution in water was held a t 260' for 48 hours and an 0.87% La2(S04)3 solution in 0.1272 m uozso4 was held a t 250' for 72 hours. KOchange was observed in either system in the amount or form of the Laz(S04)~crystals. When the tubes were cooled the crystals redissolved a t the expected temperatures. The solutions of L a ~ ( s 0 in ~ )water ~ and in dilute U02S04 solutions seemed to be stable a t 250'. The increased solubility of La~(S04)3in UOzs04 solution is probably due to strong sulfate complexing by uozso4. Cadmium Sulfate.-The solubility of CdSO4 in HzO, in 0.1272 and in 1.348 m UOzS04 solution is shown in Table 11. CdSO4, like Laz(S04)3, ( 7 ) C . H. Secoy, i b i d . , 7 2 , 3343 (1930).
SOLUBILITY OF METALSULFATES IN WATER
Jan. 20, 1957
shows a negative temperature coefficient of solubility; and the presence of uozs04 in the solution increases the solubility of CdSO4 a t a given temperature. The CdS04 solutions also show a tendency to supersaturate ; hence the temperatures given represent the points a t which the last crystals redissolved as the systems cooled. The crystals obtained a t the higher temperatures and lower concentrations either of CdS04 or of uo~so4adhered to the walls of the tube, while the crystals from the more concentrated solutions were non-adherent. CdS04, like La2(S04)3,shows no double solubility points up to 300' in water and in 0.1272 m UOSOt solution. Two double solubility points were observed in CdS04 solutions in 1.348 m uo2so4. A 31.33 wt. % CdS04 solution shows a metastable point a t 209.0' and a 30.12% CdS04 solution shows a similar point a t 215.8'. This system also shows a transition point a t 216' for 29.37% CdS04 solution as can be shown by plotting the data in Table 11. TABLE I1 THESOLUBILITY O F CADMIUM S U L F A T E I N W A T E R URANYL SULFATE SOLUTIONS CdSO4, I ("C.)
124
160
wt.
%
In water 29.6 14.8 7.4 3.0 In0.1272m
AND I N
CdSOb, 1
("C.)
21 182 47.8 (209)" 190 84.2 102.9 (215.8)" u02so4 215.8 175 28.5 216.8 200 14.2 220.8 219 7.1 228.7 251 2.8 230.6 244.3 Metastable solubility points.
wt.
%
In 1.348 m
UOZSO~ 31.62 31.33 30.50 30.12 29.37 29.69 28.68 26.54 25.83 22.51
269
a t 225'. The crystals redissolved on cooling. When the solution was heated again no crystals appeared a t 225' but the "normal" adherent crystals were observed a t 257'. The non-adherent crystals could not be produced in the tube again upon several subsequent heatings. This phenomenon was also observed in the 5.1% and in the 11.3% solutions. No explanation is given for this behavior although the points may indicate an extension of the lower half of the zinc sulfate solubility curve. The 2.1% ZnS04solution contained 0.39 mg. silica/ ml. after 3 heatings to 260' while a similar solution that had been heated only once contained 0.03 mg. silica/ml. When the 11.3% ZnSO4 solution was held a t 250' (2' above the point a t which crystals appeared in the tube) for 24 hours a very large amount of crystalline material was deposited in the tube. The X-ray analysis of these crystals showed them to be a mixture of ZnS04eH20 and ZnSO4 which seems t o preclude a hydrolytic reaction as the cause of the increased number of crystals. When the 17.0% ZnS04 solution was held overnight a t 210' (28' below the point a t which crystals normally appear) crystals appeared in the solution. These crystals redissolved readily when the solution was cooled. Their composition was not determined. TABLE I11 THESOLUBILITY
ZnSO4, wt. %
1 ("C.)
In 178 203 219 238
OF Z I N C S U L F A T E I N W A T E R
water 34.0 25.5 22.7 17.0
t ("C.)
248 253 255 257
ZnSO4, wt.
%
11.3 8.5 5.1 2.1
Cesium Sulfate.-The solubility of CsZSO4 in water, in 0.1272 and 1.348 m U o ~ s o 4 is , shown in The CdS04-UOzS04-Hz0 system did not appear Table IV. Throughout the range investigated the solubility to be as stable as the corresponding Laz(S04)8U02S04-H20 system. When the 2.8% CdS04 of CszS04showed a positive temperature coefficient. solution, in 0.1272 m uO&ot, was held for 24 hours No crystals reappeared in any of the solutions when a t 230' or for one hour a t 240' (below the point a t they were heated t o ca. 310'. The cesium sulfate which crystals normally appeared in the solution) a solutions showed no tendency to supersaturate. In brownish-orange precipitate appeared in the tube. all cases a t the elevated temperatures the crystals The precipitate redissolved when the solution was dissolved and reappeared a t the same temperature. cooled. Analysis revealed that the precipitate did The point a t 23' was obtained merely by stirring exnot contain any uranium. When the brownish cess CszS04 crystals with water in a beaker. No crystals were dissolved in dilute sulfuric acid a col- attempt was made to thermostat the solution; and orless solution resulted. When NaOH was added an ordinary mercury thermometer was used to to the solution a white precipitate, probably Cd- measure the temperature. The agreement with (OH)z, was obtained. Hence i t appears that the the values obtained by Berkeley8 was satisfactory. This system is the only system of the particular precipitate may have been CdO. Zinc Sulfate.-The solubility of ZnS04 in water ones studied in which the presence of dissolved U02also decreased as the temperature increased. SO, decreased rather than increased the solubility of The data obtained are shown in Table 111. The the saturating species. The solid phase in equilibZnSO4 solutions, unlike the L a ~ ( s 0 and ~ ) ~CdS04 rium with 1.348 m uozso4 is more than likely a solutions, did not exhibit the strong tendency t o double salt of uozs04 and CsLSO4. Yttrium Sulfate.-The solubility of Yz(S04)3 in supersaturate. No double solubility points at a given concentration were found below 300'. The water, in 0.1272 m, in 1.143 nz, and in 1.348 m crystals of ZnS04 that appeared adhered to the UOzSO4, is shown in Table V. Also included in the walls of the tube. Occasionally the ZnS04 solu- table are experimental data showing the temperations showed very peculiar behavior. For exam- tures a t which the two liquid phase region appears.' ple, when the 2.1% solution was heated for the first ( 8 ) Earl of Berkeley, Phil. Trans. Roy. SOC.( L o n d o n ) , SOSA, 210 time, non-adherent crystals appeared in the solution (1904).
ERXESTV.
2 70
JONES,
M. H. L I E ~ Z KAND E WILLIAML. MARSHALL
Higher temperatures are shown in the two liquid phase region a t which precipitation had not yet occurred. I n all cases the solubility of Y~(so4)3 shows a negative temperature coefficient. Since the YZ(SO4), solutions all showed a strong tendency to supersaturate i t was necessary to heat them 50 to 70' above the equilibrium solubility temperature before crystals appeared. The temperatures given represent the points a t which the last crystals redissolved upon slow cooling. The Y2(S04)3crystals did not adhere t o the sides of the tubes. Just as in the case of La:(SO,), no double solubility points were observed a t any given concentration below 300'. The Y,(SO~)~-UO&~O~-HZO system seemed t o be stable a t temperatures as high as 320'. TABLE 11THE SOLUBILITY
O F C E S I C M SULFATE I N W'ATER
AND I N
URANYLSULFATE SOLUTIOSS I ("C.)
CszSOa, w t .
%
t
("C.)
CSZsok,
Wt.
%
I n water In 1.348 m UO&Od 23 63.5 167.7 0.30 17j 71.8 178.5 .33 211 73.5 187.3 .37 270 75.0 214.8 .44 292 75.5 245.3 .49 In 0.1272 m UOZSOJ 285.6 .54 53.8 64.19 288, 6 . 78" 61.6 66.00 77.3 67.12 84.2 67.55 104.8 6 8 . .58 1~10.2 70.07 17i,ti 71.66 238, 6 73.48 Two liquid layers a t 293", much undissolved. (1
TABLE V THE SOLUBILITY O F YTTRIUM SULFATE I N W A T E R URAXYL SULFATE SOLUTIONS t
(OC )
\'z(so4)1, w t % In water 4.76 2.38 1.59 1.19 0.95 .64 .48 .23 .12 I n 0.1272 m UOzSOi 4.47 2.23 0.81 0.22
1
("c.)
Yz(So4)3,
Wt.
AND IS
%
In 1,145m U02SOi 208 4.32 108 215 3.41 150 217 3.44 153 230 2.78 160 245 2.12 161 1.43 163 1.08 176 I n 1.338 m UO~SOJ 186 212 4.31 195 221 3.62 224 3.43 153 23 1 2.99 162 237 2.74 189 b 2.34 229 b 2.14 No crystals up t o 'z T w o liquid layers at about 277". 320'. Held a t 320' for 2 hr. b T w o liquid layers at about 370'. No crystals after several hours above 280 to 310" for 2.14 wt. c;h Y2(SOa)3solution.
Vol. 79
there is a positive temperature coefficient of solubility. However, the solubility is increased in uo2so4 solution above the solubility in pure water,g whereas in the case of Cs2S04there is a decrease. The phase equilibria of this system were sufficiently reversible to permit precise data to be obtained. Barium Sulfate.-The solubility of BaS04 a t 250" as a function of uranyl sulfate concentration is shown in Table VII. I t is interesting to note that a plot of log Bas04 molality against the square root of the UO2SO4 molality is linear. Also, if the log of the Bas04 molality is plotted against 25' p H data obtained by Orbanlo for the particular concentrations of 'cTozs04 in water alone a linear plot is obtained with a slope of unity. 9 s yet this observation has not been explained. These data point out, a t least a t 250', the rather large increase in solubility of Bas04 in UO;SOk-H*O solution. The amount of increase is far in excess of that expected on the basis of Debye-Hiickel effects. Heat of Solution.-If the solubility data for any particular solid phase in a given solution are plotted as the log solubility, expressed as mole fraction or TABLE
VI
THESOLUBILITY OF SILVER SULFATE I X URANYL SULFATE SOLUTIOXS t ("C.)
.4ggSOa, Wt. % I n 0.1272 m UO2SOd
("c.)
AgzSok, Wt.
70
1111.348 vzUOzSOI 42.5 1.54 36.0 1.04 1.46 52.0 2.36 54,O 2.18 84.6 3.63 83.3 2.72 106.5 5.38 107.8 3.01 131.4 127.4 i.43 161.6 9.89 140.0 A red product formed" 174.3 11.01 197.1 12.82 214.9 14.14 237.5 15,7.5 259.1 17,25 Initial composition = 3.01% AgZSOa.
THE SOLUBILITY
A N D IN
UO2S04, m
Q
TABLE VI1 250' OF BARIUMSULFATE IS URANYLSULFATE SOLUTIONS
AT
U O I S O ~ , "M
M o l a l solubility B a s 0 4 X lo5
O 53 0.000 0.000 ,125 1.47 ,126 ,313 3.38 ,319 ,625 9.26 ,647 ,775 10.8 ,809 18.4 1.00 1.06 23.8 1.15 1.22 40.0 1.25 1.34 Molarities are for volumes a t 25'.
WATEK
BaSOda X 105, 'V1
0.53 1.46 3.32 8.95 10.3 17.4 22.4 37.3
molality against l / T K " . ,then, from thermodynamic considerations, the tangent to the curve at given temperature will yield a heat of solution. For the ideal system, in which heat capacities of the individual constituents remain constant and the interSilver Sulfate.-The solubility of silver sulfateg ionic interaction changes are negligible Over the Same in 0,1272 and in 1.348 m uoZs04 solution is temperature range, the tangent will also remain conshown in Table VI. In this case as with Cs2SO4 (10) E. O r b a n , "The pH M e a s u r e m e n t of U r a n y l S u l f a t e Solutions (9)
If.B a r r e , .4)2n.c h i m . p h y s . , [ 8 ] 24, 211 (1911).
from 2 5 to 60' C.," AEC Declassified R e p o r t 3583 (1952).
271 CHLORINE EXCHANGE OF ALUMINUM CHLORIDE AND HEXACHLOROCYCLOHEXANE
Jan. 20, 1957
portions of the curves corresponding to the higher temperature range except in the case of the solubility of Ag2SOI in 1.35 m UO2SO4 where the converse is true. Solid phase transitions are clearly shown by discontinuities in the solubility curves. The large decrease in the heat of solution of the TABLE VI11 solid phase as a function of increasing uranyl sulfate THE AHBoln FOR SEVERAL METAL SULFATES IX U02S04 in the solution appears indicative of a large degree SOLUTION of interaction and/or complexing of the solute AHsain (kcal./mole) with either U02S04 or the ionic species in solution. In 1.348 m I n 0.1272 m Aqueous salt systems which do not complex the Salt I n H20 UOiSO4 uo.lso4 solute exhibit negligible change in the heat of soluyz(s0,)a - 130 - 80 -48 tion as the aqueous salt concentration is increased. Laz(SO& - 120 -60 -43 Perhaps some theoretical quantitative relationships CdSOa ..... -33 - 13 can be derived to correlate the change in the heat of AgzS04 3.8 6.6 11 solution per unit change of concentration of the cszso4 1.6 1.6 4.9 electrolyte species in solution with complexing conof solution given in Table VIII. In the case of the stants of the reacting species. plots showing two slopes the values are for the OAKRIDGE,TENS.
stant, thus producing a linear function of log solubility against 1/T OK. When the data for the various systems described in this paper are plotted in this way i t is found that the majority of the slopes do remain constant and yield the approximate heats
[CONTRIBUTION FROM TECHNICAL DEVELOPMENT LABORATORIES, CONMUNICABLE DISEASE CENTER,PUBLICHEALTH SERVICE]
Chlorine Exchange between Aluminum Chloride and y- 1,2,3,4,5,6-Hexachlorocyclohexane BY GEORGEW. PEARCE AND SONJA KRAUSE RECEIVED JCLY 23, 1966 Chlorine exchanges readily between aluminum chloride and y-1,2,3,4,5,6-hexachlorocyclohexane( y-HCC) a t temperatures above 100" with some decomposition and isomerization, primarily to a-HCC and 8-HCC. Mechanisms for both exchange and isomerization are proposed.
Chlorinated hydrocarbon insecticides labeled with chlorine-36 would be very helpful in studies of the fate and toxicology of this class of compounds in resistant strains of arthropods as well as in higher animals. The use of exchange reactions appeared to offer a simpler way of preparing C1-36 labeled compounds than their complete synthesis. In addition, study of exchange reactions can provide knowledge of the structural nature of compounds in relation t o their properties in general. Using lindane, the 7-isomer of 1,2,3,4,5,6-hexachlorocyclohexane (r-HCC) , as a chlorinated hydrocarbon insecticide of importance chlorine exchange reactions between it and XlC13 labeled with C1-36 have been attempted. 1,2,3,4,5,6-Hexachlorocyclohexane has previously been labeled by direct chlorination of benzene with chlorine-36.' Experimental Both the apparatus and the method of preparing labeled aluminum chloride were based on the work of Willard and colleague^^^^ in which the reactions HCl* AgNO3 +AgC1* 4- HNOa AgC1* A1 +Ag AlC13* are employed. The latter reaction is carried out under rigorously anhydrous conditions in an evacuated system. Two solutions of chlorine-364 one containing 41.6 mg. C1/
+
+
+
(1) J. T.Craig, P. F. Tryon and W. G. Brown, AltaZ. Chem., 26,1661 (1953). (2) C. H . Wallace and J. E. Willard, THIS J O U R N A L , 72, 5275 (1950). (3) M. Blau and J. E. Willard, i b i d . , 73, 442 (1951). (4) C1-36-P Processed, Oak Ridge National Laboratory, AEC, Isotope C at alowe, July, 1952.
microcurie and the other 3.62mg . Cl/microcurie, were available for this work. The silver nitrate was Mallinckrodt Reagent grade and the aluminum was Baker and Adamson, 16 to 60 mesh. The AgCl was prepared in a tube protected from light and fused to a flask containing the y-HCC6 which was used without further purification. After the AgCl had been melted under vacuum to remove the last traces of water, A1 was added, still under vacuum. When the mixture was heated to 460-500", A1Cl3was evolved as a gas and could be driven over into the flask containing the y-HCC if this was cooled in liquid nitrogen. Heat had to be applied for one t o two hours for complete conversion of the AgCl t o AlCh. The flask containing A1C13 plus y-HCC was then sealed off completely from the rest of the apparatus and could be treated as desired with the contents still under vacuum. A number of runs were made in which the heat treatment of the AlCl3 and lindane and/or the shape of the reaction container were varied. Products from the reaction mixture were recovered by dissolving in ether and water and extracting each layer with small portions of the other solvent. Aliquots of the water phase and products obtained from the ether phase were assayed for C1-36 activity after conversion t o AgC1. A shielded end-window Geiger tube was used to count samples. The approximate disintegration rate was estimated by standardizing this count against a Q-gas flow counter which was assumed 507, efficient. Percentage exchange was calculated as percentage of total possible exchange a t equilibrium.
Experimental Results Run 1: 600 mg. lindane in 100-ml. reaction flask, 42.6 mg. C1 containing 1 microcurie. Lindane and AlCls36 heated about 7 minutes at 115-130'; no discoloration m.p. 111.5-112.5' (recrystallized product); exchange, 16%. Run 2: 500 mg. lindane in Y-shaped tube, 41.84 mg. C1 Lindane and .k1Claa6driven containing 11.5 microcuries. back and forth several times by heating alternate arms of ( 5 ) Pure samples of (I., 8 - , y-, 6- and e-isomers were obtained from t h e Ethyl Corporation.
