March, 1961
FLUORIDE COMPLEXING OF YTTRIUM(III) IN AQTJEOUS SOLUTION
441
THE FLUORIDE COMPLEXING OF YTTRIUM(II1) IN AQUEOUS SOLUTION BY ARMINED. PAUL,LINDAS. GALLOAND JEANB. VANCAMP Department of Chemistry, West Virginia University, Morgantown, W . Vu. Received August 86, 1960
The fluoride complexing of Y +a in perchlorate solution was studied potentiometrically by measuring the effect on the known ferric fluoride equilibria using the ferrous-ferric electrode. At 25" and an ionic strength of 0.5 M the equilibrium quotients for the reactions Y '-3 H F = YF+2 H+, YF+a H F = YF2+ H+, and YF2+ H F = YFY(,,s, H + are 10.0, 2 and 2, respectivcly. Equilibrium quotients are also given at 15 and 35". The stability of the yttrium fluoride complexes is compared with that of other trivalent ions. The decrease in stability of successive fluoride complexes of Y +3 IS less than that for oiher trivalent ions.
+
+
+
During the past decade there has been accumulating a vast amount of quantitative and qualitative data pertaining to the complexing of metal ions with various ligands. The lanthanide ions have been fairly well studied in this respect, but little data is available concerning the complexing tendencies of other Group IIIB metal ions. Complexing of Y+a with organic ligands' has been studied more extensively than with inorganic ligands. Complexing with inorganic ligands has been limited to an estimation2 of the hydrolysis constant, a study of sulfate complexing,a and a report that thew is no evidence of chloride complexing from anion-exchange studies.4 The recent study5 of the fluoride complexing of Sc+3 which reported that scandium fluoride complexes are more stable than those of other trivalent ions of sirnilar radius led to the present study, Besides adding to the sparse complexing data available for Y f 3 , it was felt that the study would answer the question of whether Y+3, whose electronic structure is similar to Scf3, also exhibited the tendency of forming exceptionally stable fluoride complexes. The corriplexing was studied potentiometrically by measuring the effect on the known ferric fluoride equilibria using the ferrous-ferric electrode according to the method first described by Brosset and 0rrinp;'j and later successfully used by other workers.7-!' Experimental Apparatus .-The electrodes, measuring apparatus and general procedure are similar to those described elsewhere.7-9 The half-cells were made from 200-ml. three-necked Pyrex flasks. The center neck contained a rubber stopper fitted with a stirrer. A two-hole rubber stopper accommodating the salt bridges occupied the second opening. The third opening also contained a two-hole rubber stopper, one hole of which provided the entrance for NZto prevent oxidation, and the other provided the inlet for the NaF solution. The temperature of the thermostat was controlled to =!=0.1". (1) J. Bjerrum, G.Schwarzenbach and L. G. Sillen, "Stability Constante of Metal-ion Complexee, with Solubility Products of Inorganic Substances. F'art I: Organic Ligands," The Chemical Society, London, 1957. (2) T. Moeller, J. Phys. Chem., 50, 242 (1946). (3) F. H. Spedding and 9. Jaffe, J . Am. Chem. Soc., 76, 882 (1951). (4) K. A. Kraus, F. Nelson and G . W.Smith, J . Phya. Chem. 56, 11 (1954). (5) J. W. Kury. A. r).Paul, L. G. Hepler and R E. Connick, J . Am. Chem. Soc., 81,4185 (1959). (6) C. Brosset and G . Orring, Svensh. Kern. Tid.. 65, 101 (1943). (7) H.W. Dodgen and G. K. Rollefson, J . Am. Chsm. SOC.,71, 2600 (1949). (8) R. E. Connick and M. Tsao, dbid., 76, 6311 (1954). (9) L. G.Hepler. J . W. Xury and Z. Z. Hugus, Jr., J . Phye. Chem., 68, 26 (1954).
+
+
+
Procedure.-Three half-cells, A, B and C contained the same initial concentrations of Fe( C1O4)*, Fe( C104)3 and HC104. Half-cell A contained in addition a known concentration of Y(ClO4)a. The ionic strength in each halfcell was adjusted to 0.50 M with NaClO, and the volumes in each half-cell were equal. Half-cells A and C and halfcells B and C were connected by sodium perchlorate-agar agar salt bridges. After the initial zero potentials (about 0.2 mv.) became constant, aliauots of 0.5 M NaF were added from a calibrated micropipet to half-cells A and B. The potentials of cells A-C and B-C were measured after each addition. The fluoride additions were continued until the YF, precipitated, a t which point the potential of cell A-C began to drift rapidly downward. Solutions.-A solution of Y( C10r)3-HC104 was prepared by dissolving 99.9% Yzoa obtained from the Fielding Chemical Co. in a known quantity of HC1O4. The yttrium concentration calculated from the weight of YZOSdissolved agreed with that obtained by precipitating Y+3 rn Y(OH)l and weighing as YZO3. The acidity was determined by precipitating Y +3 as YF3 and titrating with standard NaOH to the phenolphthalein end-point. The preparation and standardization of solutions of HC104, NaClO4, PTaF and Fe( C104)2Fe( C104)3-HC104 is described elsewhere.'-9
Data and Calculations Three experiments were performed a t 15 and 25' and two a t 35'. At each temperature the Y +3 concentration was varied approximately fivefold and the acidity a little more than threefold. Table I summarizes the initial conditions for each experiment. Typical data for a titration are presented in Table I1 together with calculated values of 6 and (HF)/(H+). TABLE I INITIAL EXPERIMENTAL CONDITIONS 15',
Y(C104)3, M x 104 HC104, h!f X 10' NaC104, M Fe(C104)n,M X lo4 Fe(C104)3,M X lo4 Highest (ZF-),M X l o 3
2 P , 35'
9 869 1.814 0 4726 3 752 3.663 2 953
15O,25'
I S o , 25'.
19 55 6.262 0.4224 3.725 3 627 7.246
35'
47.47 6 384 0 4045 3 619 3.524 9 346
The equilibria to be considered are
++ + + + + +
++ + + + + +
Y+* H F = YF+2 H + YF+' TIF = YFz+ H + YF2+ H F = YF?(aq, H + Fe+3 H F = FeFC2 H + FeF+2 H F = FeFt+ H+ FeF2+ HF = FeF3(aq) H + Fe+a HzO = FeOH+2 H + H F = H + F-
+
Qi
Qz
Q? (24
Q5
Q6 Qh
QHF
(1) (2) (3) (4) (5) (6)
(7) (8)
Moc!ller's2 estimate of 1 X lo-' for the hydrolysis constant of Y+3 allows one to neglect the species YOH+2 a t the acidities used. The &'a above are
Vol. 65
A. D. PAUL, L. S. GALLOAND J. B. VANCAMP
442
of the experimental points fall very close to the theoretical curve. The uncertainties in Table I11 represent the extent to which the Q's can be varied without moving the curve more than 0.02 n unit. When Q2 and Q3 are of the same order of magnitude a rather wide variation is possible in fitting the curve and hence the uncertainties in Q2 and Q3are larger than in other similar studies. Results Table I11 lists the values obtained a t p = 0.50 M for Q1, Q2 and Q3 a t 15,25 and 35'. 0
0.04 0.06 0.08 0.10 0.12 W"/(H+). Fig. 1.-Fluoride complexing of Y(II1) a t 25': o! initial conditions given in first column of Table I; 0, initial condition given in second column of Table I ; A, initial conditions given in third column of Table I.
TABLE I1 DATAFOR TYPICAL EXPERIMENT AT 25' Initial concentrations are those in first column of Table I. Initial volume = 101.00 ml.; NaF = 0.5000 M M1. N a F added
Ea-c
(mv.)
(mv.)
EB-c
-n
0.0498 .0996 .1494 .1992 .2490 .2988 ,3486 ,3994 4482 ,4980 ,8478 ,5976
17.67 32.08 43.57 53.71 63.21 69.94 76.67 82.93 88.20 93.18 97.86 101.77
22.95 40.38 54.73 66.07 76.41 85.18 92.22 99.19 105.55 110.07 113.78 116.60
0.0709 ,1217 ,2131 .2732 .3250 .3861 .4982 ,5798 .6412 .7128 .7947 .9988
t o cells A and B
I
TABLE I11 EQUILIBRIUM QUOTIENTS AT
0.02
Egp;;, 0.562 1.316 2.181 3.194 4.393 5.417 6 595 7.850 9.036 10.27 11.54 12.70
&I
Temp., O C .
15 25 35
10.5 f 0.3 10.0 f 0 . 3 9.5 f0.3
p
= 0.5 M Qs
Q1
1.0 0.5 2 5 1 5 f 2
I f 1 2 f 1 3 f l
Equilibrium quotients for reactions written in terms of F- rather than HF were obtained by dividing Qll Qz and Q3 by the ionization constant of HF a t p = 0.5.1° These values together with the true equilibrium constants a t p = 0 are presented in Table IV. The latter values were calculated using the empirical relations for activity coefficient corrections given by Rabinowitch and Stockmayer" and Kasanen.*2 TABLEIV EQUII IBRIUM QUOTIENTS AT 25' FOR REACTIONS WRITTEN IN TERMS OF FLUORIDE Ior Reaction yi-8
+ F-
YIP+* YF*+
+ +
Q'
(p
= 0.5)
YF+Z s 5 x 103 F- = YF2+ 1 6 X lo3 F- = Y F ~ ( ~ ~ )1.6 x 1 0 3
K
(p
0)
6 5 x 104 5 4 X lo3 4 o x 103
The free energy, heat and entropy changes for reaction 1 were calculated and found to be -1.36 equilibrium quotients expressed in concentrations =t 0.01 kcal./mole, - 0.93 5 0.55 kcal./mole (not activities). Values of Q4, &5, Qs, Qh and QHF used in these calculations are given by Connick, and 1 f 2 e.u., respectively. The rather large uncertainties in Q2 and Q3 prevent the calculation et ~ 1 . ' ~ of meaningful thermodynamic quantities for reThe method for calculating Q1, QZ and Q 3 parallels that given in ref. 5. The average number actions 2 and 3 . I n addition to the equilibrium quotients preof fluoride ions held by each yttrium ion is represented in Table 111, Q for the reaction YF3(,,, sented by E. HF = YF4- H + was estimated to be about 10. It was necessary to take this reaction into considera(9) tion in order to fit only the last 3 or 4 points on the The quantity (HF)/(H+) can be calculated from curve in Fig. 1. The uncertainty in these points is comparatively large since they represent measurethe equation ments close to the precipitation point of YF3. (1 (Qh) / ( H + ) ( e P E I R T- 1) = Q4(HF)/(Hf) Hence the value of 10 should be considered only Q4&5(HF)'/(H+)* Q4&5Q6(W3/(H+) (10) as a rough approximation. The relation between f z , (HF)/(H+) and Q1, Discussion Qz and Q3 is given by If one assumes fluoride complexes to be purely n = ionic, their stability should be roughly proportional to the ionic radii, provided the ionic charges are equal. The ionic radii of Y+3 and In+3 are both 0.95 Comparison of stability constants at From a plot of 6 us. (HF)/(H+), Q1, Q2 and Q3 p = 0 shows that K for the formation of k'F+2 were obtained by a process of curve fitting. is 6.5 X lo4 while that for InF+2 is 4.3 X 104.9 Figure 1 represents a plot of 7i us. (HF),/(H+) Considering the first fluoride complex only, these for the experiments a t 25'. The solid curve was (11) E. Rabinowitch and W. H. Stockrnayer, zbid., 64, 335 (1942). calculated using the Q's listed in Table 111. All (12) R. Nasanen, Acta Chem. Scand.. 4, 140, 816 (1950).
+
+
+
+
+
(10) R. E. Connick, L. G . Hepler. Z. 2. Hugus, Jr., J. W. Kury, W. M. Latimer and M. Tsao, J . A m . Cfiem. Soe., 1 8 , 1827 (1956).
(13) R. W. G. Wyckoff, "Crystal Structure." Interscience Publishers, Inc., New York, N. Y., 1951.
EXTIXCTION COEFFICIENTS O F KETONES ADSORBED OK MONTMORILLONITES 443 JIarch, 19Ci1 ISFIIARED data are in good agreement with the ionic interpretation. However if one compares the stability of the second and third fluoride complexes of Y+3 as given in Table I11 with corresponding data for other trivalent ions as given in Table V, an unusual TABLE V RQUILIBRIUN~ QUOTIENTSFOR TRIVALENT FLUORIDES AT 25’ A N D p = 0.5 Reaction Q Ref, FI.+S H F = lTeF+2 H + 184 10 Ft F+2 H F = FeFS+ H + 10.3 10 FcFz+ H F = FeFa(aq) H + 1.0 10 A1+3 H F = !k1F+2 H’ 1720 14 &\lF+2 H F = AlFz+ H + 131 14 .\I Fz+ H F = AlFz(aq) H’ 85 14 $lFa(sq) H F = A41F,- H + 0 7 14 St +3 + EtF = ScF+2 H + 1910 5 dl F”2 HF = SCFS+ H + 233 5 SCFz+ HF = S~Fa(eq) H + 14 6 5 St,Fa(aq) H F = ScF4H+ 0 85 5 I r +3 + HF = InF+2 H f 6 9 9 11F+2 HF = InFz+ H + 0 5 9
+ + + + + + + + + + +
+ + + + + + + + + + + + +
trend becomes apparent. From Table V it can be seen that the st*ability of successive complexes of 17e+3,Al+3, Sc+: and In+3 decreases by a factor of approximately 10. For the yttrium fluoride complexes, the decrease in stability of successive complexes is mu(-h less than this. As can be seen from Table 111, ihe second and third fluoride complexes of -5‘+3 ar13 of almost equal stability, and if Q for the reaction YF3(,, H F = YF,H+
+
+
has any significance at all, YF4- is more stable than either YF,(.,, or YF2+. It appears that the tendency exhibited by S C +for ~ forming exceptionally stable fluoride complexes is also shown by Y +s to some extent, particularly after one or two fluoride ions have been complexed. The explanation of these discrepancies is not apparent. The crystal structure of YF316 shows each yttrium atom to hav? 8 fluorine neighbors at 2.3 and another at 2.60 A. The structure of Y( 0 H ) P shows yttrium in the center of a regular trigonal prism with oxygen in the six corners and three more oxygens adjacent to the lateral faces of t&e prism. Assuming the radius of F- to be 1.36 A.,13the radius ratio for Y + 3to F- is 0.70, which is close to the minimum radius ratio of 0.732 proposed by Paulingl’ for coordination numbers 8 or 9. The crystal structures of YF3 and Y (OH)3indicate that a coordination number of 9 exists even though the minimum radius ratio has not been satisfied. However when the larger water molecules surround the yttrium ions, perhaps the coordination sphere may become somewhat strained and hence replacement of water molecules by the smaller fluoride ions should relieve the strain and enhance the stability of the fluoride complexes. For yttrium ion, this apparently does not occur until at least two water molecules have been replaced. (14) W. M. Latimer and W. L. Jolly, J . A n . Chem. Soc., 75, 1548 (1953). (15) A. Zalkin and D. H. Templeton, ibid., 76, 2453 (1953). (16) K. Sahubert and A . Seitz, 2.unorg. Chem., 264, 116 (1947). (17) L. Pauling, “The Nature of the Chemical Bond,” Cornel1 University Press, Ithaca, N. Y., 1945, p. 382.
IXFRARED EXTINCTION COEFFICIENTS OF KETONES ADSORBED ON Ca-MONTMORILLONITE I N RELATION TO SURFACE COVERAGE. CLAY-ORGANIC STUDIES. PART IV1 BY REINHARD W. HOFFMANN AND G. W. BRINDLEY Contribution No. 60-16from the College of Mineral Industries, Department of Ceramic Technology, The Pennsylvania State University, University Park, Pennsylvania Received Auguat 29, 1960
Infrared spectra of one- and two-layer complexes of 2,5-hexanedione and 2,5,gnonanetrione with calcium-montmorillonite were studied. The extinction coefficients for the absorptions of the adsorbed ketones were determined by a differential technique: The weight loss of a sample upon heating was related to the corresponding decrease in absorbance. The organic content and surface coverages of the clay-organic complexes were tthen determined using the extinction coefficients. The relation of the X-ray diffractionpatterns to the surface coverage is discussed. The surface coverage a t which the two-layer complex starts to form can be used as a relative measure of surface mobility. The extinction coefficients of the 1404, 1365 and 1312 (1325) cm.-‘ vibrations and that of the -0 stretching vibration decrease with increasing surface coverage. At values of the surface coverage approximating those a t which the two-layer complex appears, there is R change in the observed rate of decrease. The rate of decrease and its inflection point are discussed in terms of decreasing clay-organic interaction. The use of these observations to arrive at a site energy distribution for the adsorption is discussed.
Introduction In previous s t u d i e ~ , adsorption ~?~ isotherms for neutral aliphatic molecules adsorbed on clay surfaces from aqueous solutions have been determined (1) Part 111: L. G. Tensmeyer, R. W . Hoffmann and G . W. Brindley, J . Phys. Chem., 64, 1655 (1960). (2) G. W. Rrindley and M. Rustom, Am. Minerubpist, 4S, 627 (1958). (3) R. W. Hoffmann and G. W. Brindley, Geochim. et Cosmochim. Acta. 20, 15 (1!200).
by measuring the depletion of the organic solutions. It is difficult to remove the clay from the liquid medium without altering the equilibrium amount of organic material adsorbed on the clay surface, since any washing or drying treatment of the clay causes a change in the organic content. One of the objectives of these investigations has been to study the properties of clays in relation to the amount of organic material adsorbed, such as
