THEHYDROGEN FLUORIDE SOLVENT SYSTEM
3241
The Hydrogen Fluoride Solvent System.
I X . 1 9 2
Potentiometric Study of the
Systems: (1) Cu(s),CuF,(s)ITlF(HF)ITlF,(s) (Pt); (2)
Ag(s)lAgF,TlF(HF)IT'F,(s)
(Pt); (3) Ag(s)lAgF(HF)IAgFz(s) (Pt)
by A. F. Clifford,% W. D. Pardieck, and M. W. Wadley Department of Chemistry, Purdue University, Lafayette, Indiana (Received April 14, 1966)
The emf's of the systems given in the title, measured potentiometrically, have the values for of 0.9269 f 0.0019,1.3816 f 0.0071, and 0.5654 f 0.0069v and the values for AGOzT3 of -42.7 0.1, -31.8 f 0.2, and -26.1 f 0.3 kcal/mole, respectively. From the same were calculated data, the thermodynamic ionization constants Kb(apparent) and for the reactions AgF $ Ag+ F- and TlF $ TI+ F- in HF a t 0'. The apparent and molar activity coefficients for the electrolytes above were also calculated.
EO273
*
+
Introduction Data on cell potentials in liquid hydrogen fluoride have been reported in only a few instances. They are of considerable interest, however, because of the information which they can give, not only on the free energies of the cell reactions but also on activity coefficients, equilibrium constants, and the like, and also for the contributions which they can make toward determination of other thermodynamic properties. In a continuation of the potentiometric studies in liquid HF reported earlier,4 the standard electrode potentials were determined for the systems mentioned in the title. TJsing these data, the Gibbs free energies for the cell reactions were calculated and also the basic ionization constants and apparent and molar activity coefficients of the electrolytes AgF and TIF in H F at 0". The standard electrode potential of the single-compartment cell
1
+
neither of the electrolytes in solution could be oxidized by T1F3 or reduced by silver metal. The standard electrode potential of the cell at 0" was found to be 0.5654 0.0069 v. Similarly, a single compartment sufficed for the cell
*
Ag(s) /AgF(HF)lAgFz(s>(Pt) for which E O 2 7 3 was found to be 1.3816 f 0.0071 v. Great difficulty was experienced in obtaining stable and consistent potentials in this last case until methods were developed for the purification of the silver fluorides. In all three cases the oxidizing fluorides (TlF8 and AgF2) were prepared or purified externally and were transferred in an anhydrous atmosphere to a platinum-cup electrode, as described in ref 4.
Experimental Section Apparatus. The experimental apparatus was essentially as previously de~cribed.~
I
CU(S), CUF~(S) T1F (HF) T1F3(S) (Pt) was redetermined and found to be 0.9269 i 0.0019 V, which differs by 1.3 mv from the previously reported value.4 The potentials for the cell Ag (s) IAgF,TIF(HF) 1 TlF,(s) (Pt) could also be determined in a single compartment, since
(1) Work supported by U. S. Atomic Energy Commission under Contract No. AT(l1-1)-620; paper No. COO-620-2. (2) Taken from the thesis submitted to the faculty of Purdue University in partial fulfillment of the requirements for the degree of Doctor of Philosophy of M. W. Wadley, and for the degree of Master of Science of W. D. Pardieck. (3) Department of Chemistry, Virginia Polytechnic Institute, Blacksburg, Va. 24061. (4) A. F. Clifford and E. R. Zamora, Tram. Faraday SOC., 57, 1963 (1961).
Volume 70, Number 10 October 1966
3242
Reagents. Thallium(1) fluoride, purchased from K & K Laboratories (Jamaica, N. Y.), was analyzed for thallium content using the method of Smith and Wilcox.6 The fluoride content was determined by titration with an aqueous solution of thorium(1V) nitrate previously standardized against a solution of accurately weighed K F using alizarin red S indicator (Found: T1, 90.6%; F, 8.4%). The thallium(II1) fluoride used was that prepared by Z a m ~ r a . The ~ silver(1) fluoride, purchased from Harshaw Chemical Co., Cleveland, Ohio, was purified by recrystallization from liquid hydrogen fluoride to obtain the white salt AgHF2,which was either used as such or was heated in a platinum evaporating dish to obtain the yellow-orange AgF. This had a silver fluoride content of 99.66%. The silver(I1) fluoride, also from Harshaw Chemical Co., was purified by stirring it in liquid H F for 8 hr, filtering off the bulk of the liquid, and repeating the process twice more on the residue. This procedure removed any silver(1) fluoride and any oxide as water. General Procedure. The system was first rinsed with liquid hydrogen fluoride and this liquid was discarded. I n a drybox the electrolyte was weighed into the vessel containing a Teflon-coated magnetic stirring bar. The vessel was removed from the drybox in a polyethylene bag containing a desiccant and was quickly put in place after the T1F3or AgFz was placed in the platinum-cup electrode, while being flooded with dry nitrogen. The cell was flushed with dry nitrogen for at least 0.5 hr and finally with H F vapors. Liquid hydrogen fluoride, doubly distilled into the reservoir and having a conohm-’ cm-’ ductivity varying from 1.54 to 5.72 X corresponding to water Concentrations of 0.01 M or less, was allowed to flow into the cell with the magnetic stirrer in operation. After thorough stirring to dissolve the electrolyte, the potentials were measured by a digital voltmeter which was frequently compared with a K-2 potentiometer (Leeds and Northrup). The largest deviation was 0.07 mv. Samples were removed from the cell by forcing liquid into the buret using nitrogen pressure after the system had attained the “steady-state potential,” that is, when the potential changed less than 1 mv/hr. These were evaporated and analyzed for concentration of electrolyte. By a method of successive dilutions, further potentials at various concentrations were obtained. Treatment of Data. All calculations were performed on the IBAI 7090 computer. I n the treatment of data, two assumptions were made which seemed justified by the results. They are (1) that the electrolyte was completely dissociated at all concentrations and ( 2 ) that molarity rather than molality may be used for expressing concentrations. If the first of these assumptions The Journal of Physical Chemistry
A. F. CLIFFORD,W. D. PARDIECK, AND M. W. WADLEY
is not correct, the slope of the line obtained by plotting the square root of the ionic strength us. calculated potentials will differ from that which would be obtained if it is correct, but the intercept a t infinite dilution should not change, unless dissociation at infinite dilution is incomplete. However, judging from the observed trends of the activity coefficients, this is not the case. As to the latter assumption, in solutions for which the density is very nearly the same as that of the pure solvent, and the density of pure solvent is near unity-as it is for HF-the terms may be used interchangeably without appreciable error. The calculated potentials were obtained using the equation 2RT E”’ = E o b a d nF 1nm* (1)
+
which for the cells having n = 2 becomes
E”’
=
Eobsd
+ O.O542[’/z(lOg
[$I+]-k log [F-I)]
(2)
The apparent activity coefficients were calculated from the cell potentials according to the method used in ref 4 using the equation (3) ‘lMo1ar” activity coefficients were also calculated using the equation
where S is a function of valence and number of ions, dielectric constant of the solvent, and absolute temperature; QI = Zcfzf2( i e . , 2k) ; A is a function of ionic size, dielectric constant of the solvent, and absolute temperature; p is the number of ionic species produced by the dissociation of one molecule of electrolyte; and MI is the molecular weight of the solvent. These, however, showed no greater internal consistency than the apparent activity coefficients and were not used in the final interpretation. All fluoride ion concentrations included the contribution from the solvolysis of water (not exceeding 0.01 M ) according to the equation HF
+ HzO
H30+
+ F-
(5)
The basic ionization constant for water in HF has been found by Kongpricha and Clifford6 to have the value 0.55, for the equilibrium given in eq 6. ( 5 ) G. F. Smith 49 (1942).
and C. 9.Wilcox, Ind. Eng. Chem., Anal. Ed., 14,
(6) S. Kongpricha and A. F. Clifford, J . Inorg. Nucl. Chem., 18, 270 (1961).
THEHYDROGEN FLUORIDE SOLVENT SYSTEM
3243
I n this equation [F-] represents total fluoride ion concentration from all sources, but [HzO]represents undissociated water only. Equation 6 may be expressed for this work as
Kb
[H30+]([sum of other cations] (0.01 - [H30+])
+ [H30+]) --
1.04 f 0.03, several orders of magnitude greater t,han reported in ref 4. The system Ag(s)]AgF(HF)/ AgFz(s) (Pt) was found to follow the equation
E"'
=
[H30+],this becomes the quadratic
X 2+ ([sum of other cations]
+ 0.09526p"'
(13)
The dissociation constant of AgF in H F was calculated to be 0.087 f 0.0025. The system Ag(s) IAgF,TlF(HF) IT1F3(s)(Pt) followed the equation
0.55 (7) Letting X
= 1.3816
E"'
= 0.5654 i 0.047p"'
(14)
The data are summarized in Table I and Figure 1.
+
0.55)X - 0.0055 = 0 (8) This total fluoride ion concentration is the sum of the concentrations of all cations in the solution including oxonium ion concentration obtained by solving eq 8 for each set of data. The basic ionization constant of the electrolyte may be calculated using the equations AIF(HF)
M+
+ F-
(9)
and
where the symbols have their usual meanings and where [MF] is the concentration of un-ionized metal fluoride, i.e., the stoichiometric concentration of metal fluoride minus [AI+]. To take the effect of the ionization of water into account
+
[IT-] = [&'I+] [H30+]
(11)
the water content of the H F was always determined conductometrically immediately before each run. If now the assumption is made that the activity coefficients actually represent the degree of ionization, an apparent KB' can be calculated for each electrolyte concentration (taken as the ionic strength). Extrapolation by the IBM 7090 of --/log KB' us. p'/' to zero ionic strength gave the thermodynamic K B for the metal fluoride.
Results The E O 2 7 3 reported by Clifford and Zamora4 for the Cu(s),CuFz(s)ITIF(HF)/TIF3(s)(Pt) system was corroborated. The cell potentials were found to obey the equation E"' = 0.9269
+ 0.1002~/2
(12)
The dissociation constant of T1F in H F was found to be
0.0
I
I
0.5
I
1
I
P"'
I
I
I
I
2 .o
Figure 1. Oxidation potential vs. (ionic strength)'/2. A: Ag(s)jAgF(HF)lAgF&) (Pt); 0. run 1, A, run 2 ; curve drawn for run 2 only. B: Cn(s),CuF,(s)l TIF(HF)lTIFa(s)(Pt); V, this work; 0, ref 4. C: Ag(s)/AgF,TlF(HF)]TlF3(s)(Pt); curve omits highest point.
Discussion Dissociation Constants. The value of 1.04 found for the K B of T1F in H F at 0" may be compared with that of TlOH in water, which was measured by Bell, et al., as 0.151,' 0.141,8 or 0.1j19 at 25") the last, two values being corrected to zero ionic strength. Bell and George' found very little variation with temperature, the value measured at 0" being 0.155. Thallium fluoride, TlF, was found by these same workers to have a dissociation constant of 0.79 in water at 25". The value of 0.087 found for the K B of AgF in H F at, 0" may be compared with that of AgOH in mater at (7) R. P. Bell and J. H. B. George, Trans. Faraday SOC.,49, 619 (1953).
(8) R. P. Bell and M. H. Panckhurst, J. Chem. SOC.,2836 (1956). (9) R. P. Bell and M. H. Panckhurst, Ree. Traa. Chim., 75, 725 (1956).
Volume 70. Number 10 October 1966
3244
A. F. CLIFFORD, W. D. PARDIECK, AND M. W. WADLEY
Table I : Data for the Cell Reactions
[TlFl
YibPP)
0.0032 0.0061 0.0082 0.0277 0.0348 0.0450 0.1038 0.1518 0.3560 0.4134 0.9760 0.0450 1.8610 1.2102
1.1117 1,2399 0.8892 0,6275 0.6244 0.4708 0.1384 0.0976 0.0620 0.0146 0.0079 0.0080 0.0043 0.0003
0.7527 0.7384 0.7300 0.6850 0.6751 0.6635 0.6237 0.6048 0.5589 0.5500 0.4892 0.4594 0.4269 0.3298
0.0130 0.0158 0.0179 0.0371 0.0441 0.0541 0.1121 0.1596 0.3620 0.4191 0.9796 1.3620 1.8633 4.0639 -l/log KB(TIF) =
0.0899 1.0425 0.1047 1.0299 0.1142 1.0329 0.1800 1.0182 0.1986 1.0136 0.2226 1.0148 1.0252 0.3286 0.3946 1.0248 0.5992 1.0158 0.6452 1.0464 0.9888 1.0408 1.0728 1.1664 1.0399 1.3646 1.0985 2.0158 1.0415 - 0.77037p’/1
...
0.9238 0.9212 0.9290 0.9372 0.9442 0.9440 0.9746 0.9819 0.9917 0.9987 1.0403 1.0402 1.0545 1.1172
... 1.2774 x 3.7977 x 4.5776 x 2.2659 x 2.4921 X 1.6847 X 1.4835 X 9.065 x 6.162 X 8.787 x 5.317 X 3.6 x
10-l 10-2 lo-*
’
lo-, ’ lowa 10-55
10-75
B. The Syst,em Ag (s )/AgF ( H F)/AgF& ) (Pt )’ [AfiFI
y*(app)
1.7580 0.9622 0.5152 0.3495 0.2384 0.1440 0.8687 2.0094 1.2829 0.8001 0.5219 0.3254 0.2136 0.1291 0.07970 0.06092
0.07244 0.2504 0.4587 0.7211 0.4867 0.3444 0.3971 0.06167 0.08997 0.1420 0.2252 0.2802 0.3870 0.5553 0.7603 0.3928
[F-ltota~
p1/2
Eobsd
1.7628 1,3268 0.9694 0.9828 0.5254 0.7213 0.3616 0.5963 0.2521 0.4592 0.1595 0.3896 0.1037 0.3087 2.0116 1.4179 1.2859 1.1333 0.8042 0.8957 0.5270 0.7242 0.3317 0.5732 0.2208 0.4660 0.1371 0.3648 0.08830 0.2898 0.06977 0.2556 -l/log KB(AgF) = 0.94335
KB~PP)
EO’
1.4506 1.4487 1,4784 1.4930 1.5462 1.6080 1.6454 1,4455 1.4700 1.4929 1.5112 1.5449 1.5687 1.5978 1.6264 1.6813 ’ 0.34653p”l
1,5038 1.4454 1.4169 1.3956 1.4141 1.4304 1.4237 1.5113 1.4936 1.4721 1.4504 1.4401 1.4249 1.4079 1.3931 1.4242
C C
C
C C
C C
8.840 1.2451 2.0762 3.8117 4.0686 6.1241 1.09381 2.56566 2.2776
X X X X X X X lo-’ X lo-’ X
-
C. The System Ag (s)/AgF,TlF (HF)/TlF,(s) (Pt) [XgFl
rdapp)
WFI
r**TlF
0.017 0.032 0.064 0.117 0.139 0.219 0.464 0.781
1.7590 0.3780 0.6718 0.3472 0.1926 0.3807 0.2780
0.029 0.065 0.128 0.250 0.306 0.498 1.027 1.692
0.685 0.645 0.615 0.580 0.575 0.500 0.485 0.435
See ref 4.
‘ W.D. Pardieck.
’ Run 1.
[F-lt~ai
0.055 0.105 0.199 0.373 0.450 0.721 1.494 2.475
#l/Z
Bobsd
0.2248 0.3181 0,4423 0.6083 0.6691 0.8480 1.2216 1.5729
Data not used for calculation of E o and KB.
0.7854 0.7924 0.7253 0.7044 0.7121 0.6550 0.6142 0.7883
EO’
0.5455 0.5998 0.5795 0.6028 0.6236 0.5995 0.6106 0.8207
Not used for calculation of KB.
2 5 O , which was determined by Johnston, et aZ.,1° and Silver fluoride, again by Beck’l to be 2.0 X
more highly dissociated than the corresponding hydroxides in water. (They are stronger “solvo bases”
AgF, was found by PaulI2 to have a dissociation constant of 0.44 in water at 25”. The fluorides in HF are thus seen to be considerably
(10) H.L.Johnston, F. Cuta, and A. B. Garrett, J. Am. Cham. SOC., 55,2311 (1933). (11) M. T.Beck, Acta Chim. Acad. sci. HUW., 4,227 (1954).
The Journal of PhVsical Chemistry
THEHYDROGEN FLUORIDE SOLVENT SYSTEl
3245
Table 11: Electromotive Series a t 0" in Liquid H F In HF Couple
-In water Ref
Eana
Cd(s)-CdFi(s) Pb(s)-PbF* .b/zHF(s) Pb(s)-PbFz(s) Hz(g)-HF (1) Cu(s)-CuF*(s) FeF,(s)-FeFs (s)' Hg(l)-HgzFz(S) Ag(s)-AgF(HF) Ag(s)-AgF (HF,satd) TlF (HF)-TlFI (S) AgF(HF)-.4gFz(s) F -(HF )-F?(g)
0.29 0.26 0.17 0 -0.52 -0.58 -0.80 -0.88 -0.94 -1.45 -2.27
14 14 14 15 14 15 Assumed This work 14 This work This work
-2.708
C
Couple
Cd-Cdz+
CU(S)
+ TlF,(s) + CUF,(S)
Cu(s)
+ TlF(HF)
+ 2AgF(HF) +
+
CUF~(S) 2Ag(s) CU(S)
+ 2AgFz(s) + CUFZ(S)+ 2AgF(HF)
E" = 0.9269 v E" = 0.3615 v E"
1.7431 v
Koerber and De VriesI3 have studied the potential of the cell
+
CU(S> HgzFz(s)
+CuFz(s)
+ 2Hg(l)
and found its potential, E0273 = 0.277 v. From the very scanty evidenceI4 on the cell system Hz(g) Hg,FZ(s) + 2HF(1) 2Hg(l), the potential for this reaction appears to be about 0.8 v. For purposes of
+
+
0.41
Pb-Pbz+ H2-H cu-cu2+ FeZ+-Fe3+ Hg(l)-Hg?+ Ag-Ag
0.12 0.00 -0.34 -0.77 -0.79 -0.80
~1 +-TP +
-1.25 -1.96 -2.87
+
+
Ag +-Ag2 F --Fz(g)
'W. M. Latimer, "Oxidation Potentials," 2nd ed, Prentice-Hall, Inc., New York, N. Y., 1952. and 0. Kreft, 2. Elektrochem., 35, 670 (1929). than the hydroxides are.) Since the dielectric constants of the two solvents are nearly the same, this probably reflects the much greater solvating power of H F for anions compared with water. Since, however, there is as yet no information on the enthalpies and entropies of solvation of these fluorides in HF, the true cause of the greater dissociation can only be surmised. Oxidation Potentials. The cell potentials available through this and other work now make possible the beginnings of an electromotive series for the hydrogen fluoride solvent system. From the three cells studied, two derived potentials can be obtained such that all of the thallium and silver couples can be compared to the Cu-CuFn couple. Thus
7
Eonsa
+
' In 2.76 M KF.
K. Fredenhagen
comparison, therefore] we shall define the potential of the Hg-HgzFz couple to be -0.80 v. From the available cell potentials the electromotive series in Table I1 can be generated. Potentials for the couples of ref 13 at different temperatures and for the corresponding reactions involving the amalgamated metals are also given in ref 13. It can be seen from Table I1 that, although the actual values vary somewhat, the trend in potentials is the same in H F and in water. The spread in HF from the cadmium couple to the Ag(1)-Ag(I1) couple is slightly, but not significantly, greater than in water. If the series for H F involved ions in solution for the higher oxidation states rather than solids, the potentials would all shift to somewhat more negative values relative to hydrogen, but there is no reason to believe that the order or the spread would change significantly. This points out the great similarity between water and H F as solvents and contrasts strongly with the considerable difference between water and liquid ammonia. l6 The more negative potential obtained by Koerber and deVries14for the Ag(s)-AgF couple is to be expected from their higher AgF concentration. (12) A. D. Paul, Thesis, University of California, Berkeley, Calif., 1955; UCRL 2926. (13) G. G. Koerber and T. De Vries, J. Am. Chem. SOC.,74, 5008 (1952). (14) A.'F. Clifford and G. Balog, U . S. At. Energy Comm., Nucl. Sci. Abstr., 5 , 694 (1951); AECU 1491 (nd); and A. F. Clifford and E. M. Jeram, unpublished work. (15) V. A. Pleskov and A. M. Monosson, Acta Physicochim. U.R.S.S., 2 , 615 (1935).
Volume 70, Number 10
October 1966
