2440
DONALD
L. HORROCKS AND A. F. VOIGT
iodide, since this implies that H I is a product of the radiolysis. The effect seems to be greater than can be attributed to experimental error. According to the discussions above the rate of iodine production in ethyl iodide and presumably in the other alkyl iodides must depend on: (1) the net rate of 1 2 production in the spurs; (2) the rate of production of thermal radicals in the spurs; (3) the ratio of H I to 1 2 production. The question lilny be raised as to which of these factors is responsible for the fact that GI, with no added scavengers increases with increasing number of P-hydrogvns i n the i n ~ l e c u l e . From ~ ~ ~ ~Table I1 i t may be seen that the effect is not attributable to the rate of thermal radical production (columns 2 and S), that the net rate of I 2 production in the spurs does increase with increase in the P-hydrogen content, except in going from C2HJ to sec-CdHJ, (column 5 ) and that the ratio of thermal radical reaction with H I to reaction with 1 2 also increases (column 7). The G value for reaction of thermal radicals with H I shows the smoothest increase as a function of ( 2 3 ) E. I.. Cochran, W. H. Hamill and R. R . Williams, J r . ,
NO. 503 FROM
the number of P-hydrogens in the iriolecule (column 8 ) . Column 7 gives the total rate a t which iodine is produced in the spurs by reactions 5 and 6 , including the iodine which back reacts via t h e exchange reaction 9. -4s the carbon chain letigt?l of the alkyl iodide increases the number of bonds in the molecule other than the C-I bond increases. A s a result of this the rate for C-I bond rupture might be expected to decrease; this was found experimentally. I t can also he seen that the sc.comiary iodides (isopropyl and sec-butyl) show slightly higher rates of C-I bond rupture than the corresponding primary compounds, which might be c\pected on the basis of bond strengths. Acknowledgment.-The work of this paper was supported in part by the University Research C o ~ n inittee with funds made available by the Wisconsin Alumni Research Foundation and in part by the United States Atomic Energy Commission. During part of the work one of the authors ( R .J.11.) held a fellowship given by the Monsarito Chemic~l Company.
THIS
MADISON, ~'ISCOSSIN
I O ~ ~ R N 76, A L 214,; , (1034).
[CONTRIBUTION COLLEGE.
VOl. 7!J
T H E I N S T I T U T E FOR
ATOMICRESEARCH AND DEPARTMEST OF
CHBMISIRY,
IOWA!STATf.
u. s. ATOMIC E N E R G Y COMMISSION] The Formation of Thallium Chloride Complexes and their Extraction into Ether' W O R K W A S PERFORMED I N THE h M E S
LABORATORY OF T H E
'BY DONALD L. HORROCKS AND A. F. VOIGT RECEIVED KOVEMBER 12, 1956 The extraction of thalliuni(II1) chloride into isopropyl ether has been studied using tracer tecliiiiqui.5 w i t h Ti2". ,II? preciable extraction occurs under conditions of high acidity, and the extraction is markedly dependent upon the ionic strengt I! o f the solution. I n order t o study the effect of varying acidity and chloride concentratiou over a wide range, the study w l i i run a t rather high ionic strength. The temperature coefficient and heat of exlractiou were determitied in the range of 20-3)'. The empirical formula of the extracted species was round to be HTlC4 and a higher complex, H,TICIS, was found not t r ~ extractable. At low chloride concentrations and low acidity the extraction cannot be explained completely by this mcch:iriism and it is postulated either that TlClr extracts or t h a t HTlCld ionizes in ether under these conditions.
The extraction of thallium(II1) from hydrohalic acids into ether has been reported by several investigators.2-6 I t has been developed into an analytical method for separating thallium from most elements, except iron(III), gallium(III), gold(III), and some others. Aside from noting the optimum conditions for complete extraction of the thallium, no investigations of the nature of the extraction process have been made. This research was undertaken to obtain information about the thallium compound present in the ether phase, the dependence of the extraction upon the concentrations of acid and chloride, the temperature coefficient of the extraction and the extraction equilibrium constant. These should be of interest in comparison (1) Based in part on a thesis presented by Donald L. Horrocks to Iowa State College in partial fulfillment of the requirements for the Ph.D. degree, h-iovember, lS55. ( 2 ) Kothe, Sfokl und Eisen, 12, 1052 (1892); 13, 333 (1893). ( 3 ) (:I) A. A. Koyes, W. C. Bray and E. 13. Spear, TIIISJ O U R N A L , SO, 489. 615,559 (1908); (b) E. H. Swift, :bid., 46, 2375 (1924). (4) 1 . Wada and R. Ishii, Bull. I n s ! . P h y s . Chem. Res. T o k y o , 13, 2634 (1034): [C. A . , 28, 3334 (1934)l. ( 5 ) 13 hl. Irving, Qwnrl. Reu., 6, 200 (1951). ( 6 ) H. M. Irving and P. J. C. Roasotti, Analyst, 17, 801 (1952).
with the analogous Fe(II1) which has been showt.11 to extract into ether as HFeCL7 Preliminary investigations of the extraction of thallium(II1) from HClOd-LiCl-LiC104 solutions into isopropyl ether showed appreciable extractiot.11 only a t relatively high acid and chloride conceiitrations and also showed that the extraction depends very markedly U ~ T the I ionic strength. I t was thus necessary to perform the investigations in solutions of constant, high ionic strength. On the other hand the sensitivity of the tracer method made it possible to do the experiments at low COP centration, about M, of thallium. The results of these Conditions was that the determination of the activity coefficients of the v u ious ionic species was difficult, but not as difficult as it would have been a t higher thallium concentrations. I n estimating activity coefficients it was necessary to make some assumptions which nlay not be completely valid. The distribution of thallium between phases w a y measured by tracer techniques over a range of ( 7 ) N. H. Nachtrieb aud J. C. Conway, Tars (1948).
J O U R N A L , 7 0 , :3,5~17
FORMATION O F THALLIUM CHLORIDE COMPLEXES
May 20, 1957
244 1
TABLE I EMPIRICAL FORMULA OF SPECIES IN ETHERPHASE Initial conditions
[TI]X 104,
[Cl-1 X 10:. moles/l.
[ H + ] X 104. moles/l.
C1-/TI
H +/TI
6.66 8.63 75
2.75
5.7
3.8
8.6 5.9
4.13 4.39 4.69
0.86 1.oo 1.02
moles/l.
3 . 0 M total concn., 0.10 MLiCl, 0 . 4 M HCIOd 3 . 0 M total concn., 0 . 4 M LiC1, 1 . 6 M HCIOd
. /
chloride ion, acid and total salt concentrations. In the interpretation of the data, it was found advantageous to use formation constants determined by Benoits for the various thallium(II1) chloride complexes up to TIC&-. Experimentally, this species, in the form of the acid, was essentially the only thallium compound found in the ether phase. Hence the distribution could be represented by a constant
in which ( ) represents activity, [ ] concentration, = (T1C14-)/(T1+3)(C1-)4and ye HTlC14 is the activity coefficient of HTlCl4 in ether. This extraction coefficient, K, can be expressed in terms of the measured distribution ratio, KT = [Tile/ [TlIwby considering all of the aqueous chloride complexes of thallium, their formation constants of the form
Kd
2.70
Technique.-Aqueous solutions of the desired thallium(111), C1-- and H + concentrations and ionic strength were extracted with equal volumes of isopropyl ether by mixing in a constant temperature bath with an automatic shaker. The temperature was 30' except for the study of temperature dependence. The solutions were shaken for 40 minutes and then allowed t o separate for 30 minutes before the samples were removed with pipets. The ether samples were evaporated into 2 t o 3 ml. of water. Analyses were made for acid, chloride and thallium. The acid and chloride analyses were done volumetrically, the thallium analyses radiochemically, with the addition of known amounts of carrier T1( N03)J in.HXOs to facilitate precipitation and filtration, T h e thallium, precipitated as T12CrO4, was filtered onto small paper discs, washed with alcohol and dried a t 110". The samples were covered with 3.3 mg./cm.2 cellophane discs and mounted on cardboard backing for counting.
Discussion
The temperature dependence of the extraction was determined in two series of extractions a t temperatures of 20, 25 and 30'. I n each series the chloride concentration was varied a t constant acidity ;and ionic strength. The distribution ratio, and the appropriate activity coefficients. The K T , decreased markedly as the temperature was equation obtained is increased, for example from a solution 0.2 iM in LiCl, 0.6 M in HClO4 and 1.53 M in LiC104, KT was 4.0 a t 20°, 2.5 a t 25" and 1.5 a t 30". Values of AH of extraction were calculated by the Van't Hoff equation t o be 9.5 0.7 and 9.9 f 0.3 in which KOis set equal to 1. Hydrolyzed forms kcal./mole a t 0.4 and 0.6 M HC104, respectively, of thallium(II1) were assumed to be unimportant a t an ionic strength of 2.33. Each of these values a t the acid concentrations used. Several series of is the average of six or more calculations a t different extractions were made, each one a t constant acidity chloride concentrations. The empirical formula of the compound in the and ionic strength, but with varying chloride concentration. Since the concentration of the ethereal ether phase was determined by analyzing this thallium species was always small its activity co- phase for chloride, ionizable hydrogen and thalefficient was assumed constant. Thus a new ex- lium. Results in Table I have been corrected for the extraction of chloride and acid under identical traction constant could be defined, as conditions except for the absence of thallium. The high values observed for the Cl-/TI ratios are due principally to the small amounts present. From these data, it was assumed that the empirical formula. of the ethereal species was HT1Cl4. The observed distribution ratio, KT, decreased as Materials.-Thallium( I ) nitrate, obtained by dissolving the chloride concentration was increased. Data thallium metal in "03, was oxidized to thallium(II1) by aqua regia. The thallium(II1) chloride was extracted into a t an ionic strength of 3.0 are shown in Fig. 1. ether from a F to 7 N HC1 solution and precipitated as T1- Similar results were obtained a t an ionic strength (OH), by shaking with a dilute XaOH solution. The T1- of 2.0. A reasonable explanation of this effect is (OH)3 was purified by repeated solution in 6 t o 7 N HCl, extraction into ether and reprecipitation as T1(OH)3. The the formation of a non-extractable species, such as TlZo4P\;o3 obtained from the Oak Ridge National Laboratory, TlC15=. Since it was not possible to obtain suffiOak Ridge, Tennessee, was purified in this same manner. cient data a t low chloride concentrations to evaluA desired amount of the purified active and inactive TlT03 ate the constants for the mono to tetra-chloro comwas dissolved in excess HC104 and the resulting solution standardized, the thallium determined as TlZCr04 as re- plexes independently, the values of Benoits were ported by Kleinbergg and the excess HC104 by titration with used. His determinations of these constants, destandard base, allowing for the titration of the thallium. fined as in equation 2 are for K I , 1.38 X lo8; K2, The solution was 0.0106 M Tl(C104)3 in 4.05 M HC104. KB,6 X and K4, Equation 4 The chloride and hydrogen ion concentrations were varied 4 X and the ionic strength was maintained constant by the use was used with these constants in successive apof standard solutions of YX1, HClOd and LiC104. proximations to determine a value of Kg which would give the best evaluation of K'. This treat(8) R. Benoit, Ball. SOC. chiin. France, 518 (1949). ment gave 2.96 X 10'' for Kg. In the region of (9) J. Kleinberg, "Collected Radiochemical Procedures," A.E.C. very high ionic strength, about 5, the formation of Document LA-1721, LOP Alamos Scientific Laboratory, 1954.
*
DONALD L. HORROCKS AND A. F. VOIGT
2442
Vol. 79
The calculated values of K' are shown in Figs. 2 and 3 . The constant values of k" a t all but the lowest chloride concentrations indicate that the description of the extraction in terms of extractable HTICll and non-extractable H2T1Cljis valid. The sharp increase in the values of K' a t low chloride
I
16
-
14
-
12
-
u '
F
z
$10
-
v)
z O'O
I
1
.5 1.0 INYTIAL [LICI]
I
I
1.5 IN MOLES / L?ER,
*5
8
-3
8 -
a
Fig 1.-Extraction a t total salt and acid concentration of 3 0 AW:0 , 2.2 X HClOn; 0, 1 6 hf HClOd; E,1 0 ?if HCI04; 2 , 0 4 3IHC10,
three phases was observed which has been described in the cases of FeC&and GaC13extraction.10-12 The three phases are a "light ether," a "heavy ether" and an aqueous phase with 90-9670 of the thallium 0.0 05 1.0 in the heavy ether phase. Repetition of these exI N I T I ~[Lici) IN YQES/LITER. periments in the absence of thallium gave rise to the same phenomenon so that under these condiFig. 2. -L)epeiideilce of K' upon ILiCl] at coiistimt acid tions it is caused by the high salt and acid concen- coiicentration for ionic strength equal t o 2 . 0 : 0, I.(i JI tration and not by the heavy metal complexes. HClOi; 3, 1.0 M HCIO,: A, 0.4 .lI HClOj In the case of solutions 4 Jf in HClOl and 1 M in LiCl plus LiC104, large amounts of the acid were found in the heavy ether layer. The explanation 20( by h'achtrieb and Fryxell for the similar phenomenon in FeC13 extraction was that the acid HzFeC15 was forming ethereal polymers by multiple hydrq18gen bonds. In the present case HClO4 may be replacing the H2FeC15in such a polymer with ether 16molecules. The difficulties of interpretation which were brought about by the three phase behavior a t very - '14 high ionic strength made it necessary to conduct the more detailed study of the extraction a t lower ionic strength. The total salt plus acid concentrations were kept a t 2.0 and 3.0 M while the chloride and acid concentrations were varied. Equation 4 was used to calculate K' using the Debye-Hiickel e q ~ a t i o n 'for ~ the estimation of the activity CO-
-
eficierits of the thallium species. As the various forms of thallium(III) are large ions, the product B X a i was taken as 3 for all species. The activity coefficient of the chloride ion was assumed to be equal to the mean activity coefficient of HC1 und r the same conditions; values of 0.993 a t 2.0 31 HCl and 1.288 a t 3.0 -If HCl a t 30" were used. ( I O ) R.
W.Dodson, G.
J. Forney and E. 13. Swift, THISJ O U R N A L ,
6 8 , 3573 (1936).
(11) R . J. Myers and D. E. Metzler, ibid., 79, 3772 (1950). (12) S . H. h-achtrieb and R. E. Fryxell, ibid.. 71, 4036 (1949); 74, 897 (1952). (13) I. RI. Klotz, "Chemical 'Thermodynamics," Prentice-Hall. Nerv York, S. Y.,IYRO,11. 329.
00
05
INITIAL
1.0
(LtCI 1 IN
I5
2.0
25
MOLESILITER.
Fig. 3.-Dependence of K ' upon [LiCI] at constant acid cviicentration for ionic strength equal t o 3.0: 0, 2 , 2 AI H C I O ~ ; o,1.6 M HCIO,; A, 1.0 lir HCIO,; -, n.4 ;li HCIO,.
concentrations is attributed to one or both of two factors, the co-extraction of TlC13 or the ionization of HT1C14in highly dilute ether solutions. The directly measured extraction coefficient, KT, showed a first-order dependence on the concentration of hydrogen ion. The quantity K', involving formation constants and activity coefficients, also showed such a relationship when plotted against the HClOb concentration as in Fig. 4. The linear acid dependence agrees with the idea that HTlC14 is the extracted species and in addition indicates that the activity coefficient of the hydrogen ion is constant a t a given ionic strength. The acid dependence is expressed in equations 1 and 4 and if these truly represent the data one would expect the curves of Fig. 4 to pass through the origin. The fact that they do not, indicates that a t lower acid concentrations some other process is taking place which could be either the ionization of HT1Cl4in ether or the co-extraction of TIC&. Conclusions The following results and conclusions can be drawn from this study. The extraction was found to have a negative temperature coefficient in the range of 20 to 30'. An average value of AH of 9.7 f 0.5 kcal. per mole was found under the conditions of the experiments. In order to explain the reduced extraction a t high acidities a non-extractable species with formula HST1Cl6was postulated. The data were used to calculate its formation constant as 3.0 X 10". The empirical formula of the extracted species was found to be close to HTlC14. At a given ionic strength and acid concentration the extraction equilibrium constant was found to be essentially constant, independent of the chloride concentration except when this was quite low.
[ COXTRIBUTION FROM
2443
SOLIDARSENICHYDRIDES
May 20, 1957
THE
10
-
8 -
_ a
x
6
-
4 -
,
2 -
0.4
0.8
12
1.6
2.0
[HClO*]. of K' upon
Fig. 4.-Dependence HC104 concentration : 0 , 3 . 0 M total salt and acid concentration ; 3 2.0 A 1 total salt and acid concentration.
The lack of constancy was possibly due to ionization of HTlCla in the ether phase or to coextraction of the TlCl, species. The extraction was found to have a linear dependence on the hydrogen ion concentration in agreement with the supposition that HTlC14 was the species extracted. The fact that the extraction did not extrapolate to zero a t low acidity agrees with the idea that TlC13 is extracting or that HTlC14 is ionizing in the ether phase. Values of the extraction coefficient a t 30" are about 2.3 a t an ionic strength of 2 and 4.3 a t an ionic strength of 3 and relatively high acidities. XMES, IOWA
DEPARTMENT OF CHEMISTRY AND RADIATION LABORATORY, UNIVERSITYOF CALIFORMA, BERKELEY]
Solid Arsenic Hydrides BY W. L.
JOLLY,
L. B. ANDERSON AND R. T. BELTRAMI
RECEIVED OCTOBER 3, 1956 I'arious methods for preparing solid arsenic hydrides have been evaluated as to yield and product composition. Atom ratios of hydrogen/arsenic as high as 0.52 have been observed for solid hydrides. A method has been developed for preparing a n unstable, volatile arsenic hydride (probably biarsine, A s ~ H from ~ ) arsine by use of a silent electrical discharge.
I. Introduction The literature contains many conflicting statements regarding the solid arsenic hydrides. Not only the compositions of these hydrides, but also their methods of preparation, have been the subject of dispute.' We have evaluated various preparative methods in terms of both yield and hydrogen conten6 of the product. Several new preparative methods have been developed, including a method for the preparation of an unstable, volatile hydride (probably biarsine) . (1) See, for example, "Gmelins Handbuch der Anorganischen Chemie," System-Nummer 17, Verlag Chemie, G. m. b. H . , Weinheim/Bergstrasse, 1952, pp. 195-200.
Some studies of the adsorption of arsine on arsenic have been made, in an attempt to elucidate the nature of the solid arsenic hydrides.
11. Results and Discussion A. Reduction of Arsenic Solutions. 1. Reduction by SnC12.-Weeks and Druce2 reported that a brown, amorphous arsenic hydride may be prepared by the action of an ether solution of stannous chloride on a solution of arsenic trichloride in aqueous hydrochloric acid 2AsQ + 4 ~ + 2 ~ ~ ++~2 ~ 1 1-= A~ ~ +~ . & H I , ~ ] r~ -( 2 ) E. J. Weeks and J. G. F. Druce, Chem. N e w s , 129, 31 (1924)
