1106
ALBERTW. JACHEAND GEORGEW. CADY
up to the moment of reaction. The large reaction distance and small electrostatic contribution in system 3 may thus find an explanation. Of the halogens it appears that only the F- ion has any substantial solvation2*but data for F- substitution are not available to check this interesting correlation between electrolytic and kinetic behavior of ions. The thiosulfate ion is strictly not spherically symmetrical but the double charge may be regarded as “smearing any such deviations. This double charge incremes r by a factor of 1.2. The effective diameter of the Szoa--ion is not known but is certainly large so that the order of magnitude of the values recorded for system 4 is not unexpected. From the magnitude of r , the distance between ion and &pole center, we call arrive at a figure for d, the distance between the approaching ion and the C-atom at the moment ofreaction. The vagueness of the concept “center of the dipole” makes a definite calculation diffcult. We may, however, make the reasonable assumption (cf. ref, 10) that the center of the dipole lies close to the point of contact of the covalent radii of the atoms responsible for the dipole, ice., at 0.77 A. from the carbon center. Neglecting the slight variation of d/r for small angles, we may subtract this figure from the values of and obtain the true kinetic reaction distance. Systems and only are given since for reasons discussed above it is not possible to give a precise structural significance to the values in the other systems. (23) Bernal and Fowler, J . Chsm. Phys., 1, 515 (1933).
Vol. 56 TABLE111 Br.. .C
I . . .c
d (HzO) d (CH,OH) d ((CHs)&O)
1.42 5.79 5.26
Transition state Electron diffraction Microwave
2.34 1.91 1.939
2.27 2.53 5.08 2.02 2.55 2.10 2.144
Distance
The distances so calculated may be compared with the average internuclear separation of the final products from electron diffraction and from microwave values26,2u as well as with the Values Of the transition State distance estimated on theoretical grounds from non-kinetic data. lo The divergence of the d values is considerable and the uncertainties of these experimentally difficult reactions makes it unprofitable to discuss the recorded differences. The advantage of the present method lies not in the values SO obtained but in the fact that it yields such a value without the rigid assumption of a single direction of approach. I t also yields without further assumptions a value for the deviation term. The variation of reaction distance and activation energy with the total solid may produce Prefererltial reaction in Orle direction. Kinetic analysis will detect this by a low probability factor. In other cases no serious directional differences arise with a consequent Of p r (24) L. Pauling, “Nature of the Chemical Bond,” Cornell University Press. Ithaca, N. Y.. 1945, p. 167. (25) Simmons, Phgs. Rm., 76, 686 (1949). (26) Simmons and Swan, ibid., 80, 289 (1950).
s o L u m I w OF FLUORIDES OF METALS IN LIQUID HYDROGEN FLUORIDE BY ALBERTW. J A C H EAND ~ GEORGE H.CADY Departtnent of Chemistry and Cheinical Engineering, Universily of Washinglon, Sealtlc, Washington Received May d?. 1966
Solubilities of the fluorides of a number of metals in liquid Iiydrogen fluoride have been measured over a kmperature range from about -25O to approximately 12’. In general the values found are in accord with the qualitative information available in the past. Many of the solubilities are of the same order of magnitude as those of the corresponding hydroxides when dissolved in water; however, it appears that in general the fluorides are more soluble. Some fluorides are much more soluble in hydrogen fluoride than the corresponding bases are in water. For example, the value for silver(1) fluoride is 3 X IO‘ times that for silver oxide.
I n spite of the fact that it is well known that liquid hydrogen fluoride resembles water by being a highly polar polymerized solvent for electrolytes, it is still true that only a limited amount of quantitative information about its solutions is to befound in the literature. For example, in recent reviews both Simon+ and JltnderZbgive qualitative information regarding the solubilities of many substances but report only a few quantitative values. The (1) This paper represents part of a thesis submitted by A. W. Jache to the Graduate School. Univereity of Washington, in partial fulfillment of the requirements for the Ph.D. degree. (2) (a) J. H. Simons, “Fluorine Chemistry,” Academic Press, New York, N. Y., 1950, pp. 225-292; (b) Jander, “Die Chemie in Wasserahnlichen Losungsmitteln,” Springer-Verlag, Berlin, 1949. pp. 1-37.
purpose of the present paper is to increase the number of quantitative data and to compare the solubilities of fluorides in hydrogen fluoride with the solubilities of hydroxides in water. Experimental Preparation of Fluorides. -- Hydrogen fluorirk- was prcpared by a mehhod similar to that of Simons.3 A solution containing hydrogen fluoride and potassium fluoride in a ratio of about three moles t,o one was obtained by absorbing commercial hydrogen fluoride in molten KHF2. Traces of water and other impurities were removed by an electrolytic decomposition which was continued until fluorine had been liberated at the anode for two houri or more. Hydro(3) J. H. Simons, “Inorganic Syntheses,” Yol. 1. hlcGraw-Hill Book Co., Inc., New York, N. Y., 1939, p. 134.
Dec., 1952
SOLUBILITY OF FLUORIDES OF METALS IN LIQUID HYDROGEN FLUORIDE
gen fluoride was then distilled from t'he mixture through a vertical 60-cm. length of copper tubing packed with small copper ribbon. The condensed liquid was collected in a polyethylene bot,tle. Later it was dist,illed from this bottle through a system, madc of polyethylcne, into t,hc bottle to be used in a solubility mcasurcmcnt. Liquid prcparcd in this manner was found t;o have a specific olcctrical conductivity of about 7 X 10-4 ohm-' em.-'. This indicates a rather high degree of purity but not as high as that at,tained by Fredenhagcn and Krefft.' In thc table, the second column indicates by letters the means used to obtain the metal fluoridcs. The letters a, b, c, etc., stand for the methods now to be described. (a) These fluorides were commercial chemicals or were obtained as research samples from chemical manufacturers. Each fluoride was placed in the bottle to be used for the solubility nieasurcnicnts arid was then washed three to five times by decantcation with liquid hydrogen fluoride. This washing procedure served t,o remove water from the sample, to convert, impurities such as carbonates, chlorides, oxides, etc., to fluorides and to dissolve any traces of impurities morc soluble t.han the fluoride being studied. I t was uscd with all of tbe fluorides discussed in this papcr except silver ( I )fluoride and t,halliuni ( I ) fluoride. (b) These fluorides were prepared by the action of liquid hydrogen fluoride upon carbonates of thc nwt.als. (c) Fluorides indicat,cd by this letter were produced by the action of liquid hydrogen fluoride upon nitrat.cs of t.he metals. (d) Cerium fluoride was precipitat,ed from a n aqueous solution of the nitmtc. The solid was washed with water and then with hydrogen fluoride. (e) Fluorides of lithium and calcium were formed by thc act8ion of hydrogen fluoride upon the metals. After the mctal had all dissolved its salt was washed by decantation several times with hydrogen fluoride. (f) .Aluminum fluoride was obtained by dissolving thc metal in aqueous hydrofluoric acid. The salt was washed in the usual manner with hydrogen fluoride. (g) Mercury(1) fluoride was formcd by the action of hydrogen fluoridc upon mcrcury(1) osidc. (h) Bismuth(II1) fluoride was prepared by reaction of bismuth hydroxide with hydrogen fluoride. (i) Iron(I1) fluoride was obtained from iron( 11) chloride and hydrogen fluoride. ( j ) Thallium( I) fluoride was prepared by dissolving met,allic thallium in aqueous hydrofluoric acid. This process required a long t.ime even though the sticks of thallium wcre wrapped wit,h platinum wire t.0 serve as a cat,alyst. The resulting solution was evaporatcd t,o dryness. Hydrogen fluorido was then added and was later cvaporated away. The salt obtained by this process was found to contain 8.48% fluorinc (theor., 8.51o/b). ( k ) Thallium(II1) fluoride wa.7 formed by passing fluorinc, diluted with nitrogen, t,hrough a solution of thallium(1) fluoride in hydrogen fluoride. As the fluorine passed, thalliuni(II1) fluoride precipilat,ed. Even t>hough a large excess of fluorine was used, much of the t~halliuni(1)fluoride failed t,o react,. The prccipitatcd thallium( 111) fluoride was washed twi.ce by derantmattionwith liquid hydrogen fluoride. Finally it was again suspciitlctl in hydrogen fluoride in the solubility bott,lc and was again Imated with dilut,c fluorinc. (1) Silvcr(1) fluoridc was niado by thc action of liquid hytlrogcii fluoridc ul~onsilver carboiiatc and upon silvcr oxide. The solvent was cvaporat,cd away. The solid was again dissolved in hydrogni fluoride and the solvent was again cvaporated away; fluorine contci!t found 15.2% (theor. = 15.0%). (m) Silvcr(I1) fluoride was prcparcd from silvcr(1) fluoride using 1 lit! same type of procedure as tpat described for thallium(111) fluoride. Preparation, Sampling and Analysis of Solutions .-VCSscls madc from polyethylene plastic werc cniploycd in t.hcsc studics. This material is resistant to att>ack and is conveiiicnt 1.0 use bwausc tubing can be fused toget,hor and apparatus can be :tsscniblctl in inuch (,he sanic manlier as glass a p p a r x t . ~ ~The . suitabili1.y of Ilic plast.ic was shown by aI1owiiig a 2i0-1nl. 1101 t.lc t.o st.aiid tilled wilh licluid hytlroErn fluorido for four clays. Tlic dry l)ot,t,l(!wcighvd about, 93 g . and 6hc t.r(.lttnicwl, !villi I i y t l r o ~ ~Iliioritl(: ~n c-ausc.cl it 10 Increase in weiglil. 16 nig. 'l'li(w wts a slight. (lerkrnirlg o f rolor. A srconcl t.reatnirnt, wit.11 hytlrogcw fluoritlc crrtuscbti ihe bott.lc 1.0 incrcase iii ncbiglit by only OIIC ~nilligraii~.All bottles uscd in this research were prot.rcbatcd for several days with liquid hydrogen fluoridc. When a t.cn-gran1 sample of hydrogen fluoridr which had stood in a bottle \
___
~-
(4) K. Fredenhagen and 0.T.Ifretlt, 2.Elekirochenr., 35,670 (1929).
1107
for R week was cvaporated to dryness in a platinum crucible, t'hc residue remaining weighed not over 0.1 mg. As the rcsult or continued use, t,hc bot,tlcs darkened in color somewhat. Solut)ilit,y bottles were madc from 270-rnl. bot>t,lcs. Tw-o picccs of 12 mm. ( i d . ) tubing were sralcd to each bottle to siorve as inlot and out,lct for gases. These tubes could be closed with Teflon (polytct,rafluorocthylcnc) stoppers. The bottle was puncturcd near the top with a nredlc which left a hole of about 0.5 mm. diamctcr. This hole was opened when a sample was to be taken; at ot.her times it was kept closed wit,h a piece of polyethylene cemented in place with a waxy form of poly-(clilorot.rifluoroethylenc). Sampling bot,tles had volumes of about 70 ml. Each bot& was attached to a no. 20 hypodermic nccdle by m a n s of a Teflon adapter. The only opening in t.hc sampling apparatus was the holc in t,he needle. A solution was prcparcd by washing an excess of the mctal fluoride to be studied n4t.h hydrogen fluoride. Several decantations were made as indicated undcr method (a) for obt.aining fluorides. Finally, the solubi1it.y bottle was about two-thirds filled with liquid hydrogen fluoridr. In most. cases dilutcd fluorinc was but)blcd through the solut,ion for an hour or more t,o react with m y water which might have been prcscnt. This st,c!p was omittrd when one of t.he following fluorides was prescmt : nicrcury( I), silver(I), t.hallium(I), cerium(III), lcad(II), antiniony(III), bisinut'h( I I I ) , iron(II), cobalt(I1) and nickcl(I1). The bottlc was then closed with Teflon stoppers and it was placed in a low frequency shaker in an air t,hcrmostat. At approximat.ely daily intervals saniplcs of the clear supcrnatant solut,ion wore wit,hdrawn into sampling bottles. First, the solid was allowed t,o sett.lc in tfhc solution; then the sampling bottle, which had been cooled t.o thc temperature of t.hc solution, was squccacd, and the hypodermic ncc!dlc was inserted t.hi,ough t.hc tiny holc in t,he wall of the solubi1it.y bott,lc. As the sampling bottle expanded to its original size a sample of liquid flowcd in through the needle. The holc in t.he solubilit,y bottlo was then covered at once, and thc sampler needle was capped wit,h a piece of polyet,hylcnc plastic. The siao of th(. saniplc was determined by weighing thr. sampling bottle bdorc and after discharging the sample for analysis. Usually the sample was analyxod simply by evaporating it, to dryness in a platinuni crucible and weighing the residual mctal fluoridc. Normally the fluoridc was dried in the at.mosphere a t 110" bcfore weighing. Sodium fluoride was heat,ed t.o redness, however, 1.0 drive off all hydrogcn fluorid(!. I n thc cases which follow othw ana1ytic:al mothods were usc!cl: ( 1 ) Silvcr(1) fluoridr: \vas det.ermined by discharging the solut.ion int,o hydrochloric acid. The silver chloride so formed was weighed. (2) Silvcr(I1) fluoride was usually determined by discharging the solution int,o a inixturc of aqueous pot,assium iodide and ice. The iodine so formed was titrat,ed with sodium thiosulfate. h comparison of t.his method wit,h t,hc ono usually uscd for other salts indicat.ed that the two wcrc in agrc!c!mcnt within Ihc limits of accuracy. (3) hl(:rcury(I ) fluorido solut,ions wwc anelyatd 11y the pivccdurc uscd for silvcr(1) fluoridv. ( 4 ) hlcrcury( 11) fluoride obtairicd by cvaporat,ion of the solvc!rit, was collvcrtcd to the oxide and wcighed as such. (5) Solulaionscontltiiiiiig tjhalliuiii(I ) fluoride wcrt! added 1.0 a mixt>uroof ice and watcr. Potassium iodide was then added and t'he weight, of tjlie precipit,atc!d thallous iodide was dct'crmincd. (6) Thalliurn(II1) fluoride was detcrmincd by the method uscd for silvcr(11) fluoride. In t,hc onc case testcd, t8hismethod agreed substantially wi tb the method usually cmploycd for other salt,s. (7) Solutions of manganese( 111) fluoride were analyzed by prc!cipitat.ing mangancsc dioxide. This was dried and then weighed. (8)Sinec solid iron(I1.I) fluoridc was vnry hygrosco1)ic it was converted to ferric oxide and wcighcd as such. Analysis of. the Solid Phase.-Afl.c!r l h e solubility deterniii~af ioris for a salt had all 1 1 ~ 1 iliado, 1 (he cap was removed froin lhc! solut)ilit,y bot xntl soniv of t.lic solid pliasc was an 80-mesh romovrcl oil LL spatiila. Tliis slurry was plttrtrtlo~~ c!oppcr gauze filler and liquid was quickly rvriioved by ceiit r i f u g a t h . Thc solid was t.lic!n wrighrd in a plat.inuni crucible and in iiiosL casos it was dried a t 110' 1.0 remove hydrogen fluoride. Loss 111 weighL upon drying corrcspondcd t.o hydrogen fluoride driven off. For sonic of thc salts thc special analytical procedures described in t)hc preceding paragraph were used.
ALBEIWW.JACIIEAND GICOI~GE W.CADY
1108
Vol. 56
TABLE I 80LUBIL[TY OR' PLIJORIDES I N 1 ~ Y U l t O U J 3 NFLUORIDE
Salt
Preparation method
LiF
a,
NaF KF RbF CsF NHdF
a
e
Ref. 6 Ref.7 Ref.8 Ref.5
(hFn b AgF AgFz BeF2 MgFi CaF2 SrFI BaFt ZnFZ CdF, HgFt HgzFs AIFa TlFa TIF CeFI CeF, ZrF, ThFd
1 m c
b b,e a a
b b a
g f
k j d a c
c
PhFp b SbFg a BiFI h MnF, FeFI FeFI CoFI
a a
i a COF~ b NiFz h
QC.
12.2 11.0 8 20 10 17 12.4 11.9 11.5 11.2 12.2 12.2 12.2 12.2 14.2 14.2 11.9 11.8 11.2 11.5 11.9 11.9 11.9 12.4 11.8 12.4 11.9 12.4 11.5 11.9 11.8 11.9 14.2 11.9
Solubility in g. unsolvatod salt per 100 Solubility OC. Solubility
10.310.1 30.1f0.1 (36.5) (110) (199) (32.6) 0.010f0.005 83.2f0.8 0.048f0.006 0.015 1 0 . 0 0 4 0.025f0.003 0.817 f 0 . 0 1 5 14.83 f 0 . 0 9 5.60f0.12 0.024 1 0 . 0 0 2 0.201fO.009 0.54f0.01 0.877 h 0 . 0 0 5 s0.002 0.081fO.003 580 f 4 5 50.043 O.lOfO.O1 O.009fO0.002 S0.006 2.62f0.09 0.536 f 0 . 0 0 5 0.010*0.002 0.16410.004 0.008f0.002 0.006f0.002 0.257 1 0 . 0 0 3 0.03Gf0.002 0.037 f 0 . 0 0 2
g. of
-
3 . 3 10.3 f 0 . 1 9.8 2 5 . 1 f 0 . 1 -45 (27.2)
-
-16
(177)
-
0 . 0 0 8 l t 0 004 43.8f0.1 0.030 f 0 . 0 0 2 0.013 f 0 . 0 0 3 0.025 1 0 . 0 0 4 1.061f0.005 14.63 f O . O 1 4.74320.04 0.019 l t 0 . 0 0 2 0.198f0.001 0.62 f O . O 1 0.81 f 0 . 0 2 50.003 0.029 rtO.005 450f50 10.037 0.1010.01 0.015f0.003 10.002 3.56f0.04 0.285 f O . 0 1 0 0.011 f 0 . 0 0 3 0.147rt0.010 0.003 f 0 . 0 0 2 0.005f0.002 0.264 f 0 . 0 1 5 0.033&0.002 0.040 f 0 . 0 0 2
8.3 -9.8 8.9 5.1 - 3.3 3.3 3.3 3.3 4.4 - 4.4 7.8 4.5 5.1 - 8.9 7.8 7.8 9.8 - 8.3 4.5 8.3 9.8 8.3 -7.8 7.8 - 4.5 9.8 4.4 9.7
-
-
-
Results For each salt the solubility was determined a t three different temperatures. A run was continued a t one temperature until the concentration of solute in each bottle remained constant for about three days or more. Daily analyses were made using samples of solution weighing about 10 g. The solubility values are listed in Table I together with an indication of the apparent precision of the measurement. The meaning of the table may be clear by considering the row for lithium fluoride. This salt was obtained in two ways. A commercial sample was used and another sample prepared from lithium metal and hydrogen fluoride was also tested. Both samples had the same solubility. At 12.2" and also at -3.3' and -23.0' the experimental values for the solubility ranged between 10.2 and 10.4 g. of LiF per 100 g. of hydrogen fluoride. This uncertainty in the analysis is represented in the table by reporting the solubility t o be 10.3 f 0.1. The 10.1 is not the probable error; it represents the precision of the analyses or, in the case of low solubilities, the limit of accuracy which could be reached with the balance. The larger of these two limits is the one recorded in the table
bxdrogen fluoridc C. Solubility
Compositiorr of aolid (not dried), Former molea HE' Solirldity iolubility per mole salt ratio rating
-23.0 -24.3
10.3 f 0 . 1 22.1f0.1
1.09 4.10 4
-23.1 -25.0 -25.2 -24.2 -23.0 -23.0 -23.0 -23.0 -23.0 -23.2 -25.2 -22.5 -24.2 -25.2 -25.2 -25.2 -23.8 -23.1 -22.5 -23.1 -23.8 -23.1 -25.2 -25.2 -22.5 -23.8 -23.2 -25.0
3 5 0.010f0.004 0.28 27.2f0.7 1.94 0.024 f O . O O 1 0 . 0 0.014 f 0 . 0 0 2 0.03 0.033 f 0 . 0 0 2 0 . 2 3 1 . 4 4 4 f 0 . 0 0 4 1.86 14.43 f O . O 1 3.22 3.61 f O . O 1 5.95 0.016f0.002 0.05 0 . 1 8 9 f 0 . 0 0 1 0.07 0.61f0.01 0.01 0.79f0.02 0.00 10.004 0.22 0.027 f 0 . 0 0 3 1.04 305 i 1 5 1.92 S0.041 0.53 0 . 1 0 6 f 0 . 0 0 6 0.03 0 . 0 2 3 f 0 . 0 0 2 1.00 10.001 0.84 3.67f0.01 2.23 0.191 f0.003 0.13 0.01OzkO.002 0.16 0 . 1 3 4 f 0 . 0 0 3 0.11 50.001 0.60 0.005 f 0 . 0 0 2 0.32 0.272 f 0 . 0 1 6 0.35 0.040f0.003 0.38 0.035 f 0 . 0 0 2 0.25
0.74 0.25 0.30 0.5 0.G4
32 3 X 10' 1 . 6 X IO' 8.7 6.8 16.5 1.15 1.1 X 102 75 10
5 5 5 5 5 5 1 S
SS
ss ss ss ss I
1 1
..
1.2
1
8.5
8 1
..
.. ..
.. 2 . 3 X102
I
6.5
1
5 X loz 80 7.7 X 10'
I
..
27
*. 1
.. 1 1
for each solubility value. In the column markctl "Composition of solid" the value 1.09 means that the centrifuged sample of lithium fluoride as removed from the solubility bottle lost weight upon drying corresponding to 1.09 moles of H F per mole of LiF. This is an indication, but not a proof, that the solid in the bottle was LiFeHF. Values in this column mag fail to represent the true composition of the solid phase for two reasons: (1) an appreciable weight of liquid might remain on the solid even after centrifugation, thereby giving too high a value; (2) hydrogen fluoride might be lost from the solid before the sample coiild hc weighed, thereby giving too low a value. In spite of these experimental difficulties, many of tlic values appear to be clear in meaning regarding thc composition of the solid phase, For example: some of the solid phases apparently were LiFeHF, NaF.4HF, AgF.BHF, AgFz, BeF2, MgFz, CaF2. 2HF, SrF2.3HF,BaFAHF, ZnFz, etc. Data in the table for NHaF,SKF,6 RbF? and CsF8 have been selected from the literature and the refer( 5 ) 0. Ruff and L. Staub, 2. WID^^. dbem. Cham., 212. 399 (1933). (8) G . H. Cady, J . Am. Chsm. Soc., 56,1431 (1934).
(7) K. R. Webb and E. B. R. Prideaux, J . Cham. SOC.,1 1 1 (I!J39). ( 8 ) R. V. Winsor and G . H. Cady, J . Am. Chem. Sue.. 70. I500 (1948).
Dcc., 1952
SOLUBILITY OF FLUORIDES OF METALSIN I,IQUID HYDROGEN FLUORIDE
ences are listed in the column labeled "Preparative method." These solubilities were not determined in this research, but they are included to permit the tabulation-of these fluorides.
Discussion
1109
the "solubility ratio" column. This quantity is defined by the equation solubility ratio = ( C M ) H B / ( ~ Y ) H ~ O
in which (CM)HFis the concentration of the metal in gram atoms per 100 g. of hydrogen fluoride in a
In the table, the column giving former solu- solution saturated with the fluoride of the metal is the concentration of the hility ratings uses the system of notation employed at loo, and (CM)H*O by Simons.' The letter S means soluble probably metal in gram atoms per 100 g. of water in a soluover 1%, while SS means slightly soluble, and I tion saturated with the hydroxide of the metal at stands for insoluble. Much of the information 25". When calculating values for this ratio, used by Simons in the preparation of his table solubilities of the hydroxides were taken from was obtained from a paper by F r e d e n h a g e ~ ~Seidell.l1 Solubilities for the oxides rather than Although there is a general agreement between the the hydroxides were used for the elements silver(I), present research and these qualitative ratings, mercury(I1) and lead(II), and a few of the soluthere are some fluorides which are actually more bilities were taken a t 18 or 20' rather than at 25". It is apparent that the fluorides, in their solusoluble than the ratings indicate. A majority of the few quantitative solubilities, now reported in bilities, are much like the corresponding hythe literature, for salts studied in this research are droxides. Thus, the fluorides of the alkali and lower than the values in the table. Lithium fluo- alkaline earth elements are soluble, and the soluride has, according to Fredenhagen,s a solubility bility increases as one goes to elements of higher of 2.B g. per 100 cc. of solution a t 18" while its atomic weight in one of the families. Strontium solubility as given by Bond and Stowelois close to fluoride is out of line, but its high solubility as 0.044 mole of LiF per mole of H F over a tempera- compared to that of barium fluoride is probably ture range from 0 to 40". This corresponds to due to the difference in degree of solvation of the 5.7 g. of LiF per 100 g. of hydrogen fluoride and is solid phases of the two substances. One here comconsiderably less than the va.lue, 10.3, found in this pares the solubilities of SrF2.3HF and BaFz.6HF. research. Bond and Stowelo found the solubilities Probably BaF2.3HF would be more soluble than of ca.lcium fluoride, zinc fluoride and magnesium SrF2.3HF. Fluorides of the subgroup elements of fluoride at 0" to be less than 0.01%. Freden- group two are less soluble than the alkaline earth hagen'@ solubility of silver(1) fluoride a t - 15' fluorides. There is a decrease in solubility in the is 33 g. of the salt per 100 cc. of solution. While order: sodium fluoride, magnesium fluoride, alumithis may be in rather good agreement with data in num fluoride. The expected type of behavior is the table, one cannot be certain, because the shown by silver, mercury, thallium and iron in that the lower valence fluorides are more soluble density of the solution is not known. A sample of commercial hydrated ferric fluoride than those in which the metal has its higher valence. By contrast, the higher valence fluorides of was found to be very soluble. Repeated washing of the solid by liquid hydrogen fluoride yielded cerium and cobalt are the more soluble. Several of the fluorides, particularly silver(1) fluoanhydrous ferric fluoride of low solubility. Perhaps the influence of a little water upon the solubility ride, are relatively much more soluble than the corof this salt is the result of the formation of fluoride responding hydroxides and none are strikingly less ions by the reaction HzO H F = HsO+ F-. soluble. Perhaps this means that on the whole, Fluoride ions so formed may dissolve ferric fluoride fluorides are more basic in hydrogen fluoride than due to the formation of complex ions such as hydroxides are in water. FeFa---. Hydrated chromic fluoride was also Acknowledgment.-This research was performed found to be very soluble, but the action of hydrogen under contract with the Office of Naval Research, fluoride did not yield a satisfactory anhydrous salt. U. S. Navy Department. The authors are grateful Since one may regard the metal fluorides as bases to the Whitemarsh Research Laboratories of the in the hydrogen fluoride system of compounds, it is Pennsylvania Salt Manufacturing Co. for analyzed of interest to compare their solubilities in hydrogen samples of mercury(I1) fluoride, cerium(1V) fluoride with the solubilities of the corresponding fluoride, manganese(II1) fluoride, cobalt(II1) fluohydroxides in water. This is done in the table in ride and iron(II1) fluoride.
+
+
K. Fredenhagen, et aZ., 2. physik. Chem., 8164, 176 (1933). (IO) P. A. Bond and V. M. Stowe, J . Am. Chcm. Soc., 53,30 (1931).
(9)
(11) Seidell, "Solubilities of Inorganic and Metal Organic Compounds," D. Van Nostrand Co., Inc., New York, N. Y., 1940, 3rd ed.
