Anal. Chern. 1983. 55. 1089-1094
in need of additional investigations. RTP Polarization Study. The degree of polarization of R T P and LTP from compounds adsorbed on PAA-salt mixtures and filter papeir was measured. The purpose of Ibis study was to rellate the degree of phosphorescence polarization to the rigidity of the phosphor adsorbed in the solid matrix. With 4-phenylphenol adsorbed on the solid surfaces, zero polarization of the R T P and L T P was; obtained within experimental error. Thig result is not surprising since radiation scattered by a !solid surface causes loss of orientation of both exciting and emitted light (27). Scattering, therefore, causes depolarization of luminescence to occur. Teale (28)has described a method by which the polarization of luminescence emitted from turbid solutions could be measured by observing a thin layer of the luminescent volume in the plane of incidence and observation. Faucon and Lussan (29) demonstrated Teale's method of correcting for depolarization by scattering by using horizontal slits which were 1 mm high in the excitation and emission beams. We used the above method of placing 1mm horizontal slits in the exciting and emitted beams to correct for depolarization of R T P and L T P emitted from solid supports. However, within experimental error zero polarization was measured. Apparently the depolarization of RTI? from solid supports is too extensive to be corrected1 for by the method deslcribed above. Registry No. 4-Phenylphenol,92-69-3;4,4'-biphenol, 92-88-6; biphenyl, 92-52-4;2-naphthol,135-19-3;2,3-dihydroxynaphthalene, 92-44-4; a-naphthoflavone, 604-59-1; 9-anthracenemethanol, 1468-95-7;NaCI, 7647-1.4-5; NaBr, 7647-15-6;poly(acry1ic acid), 9003-01-4.
1089
(2) Parker, R. T.; Freedlander, R. S.; Dunlap, R. B. Anal. Chim. Acta 1980, 179, 189. (3) Hurtublse, R. J. "Solid Surface Lumlnescence Analysis: Theory, Instrumentation, Applications"; Marcel Dekker: New York, 1981; Chapters 5 and 7. (4) Dalterio, R A.; Hurtublse, R. J. Anal. Chem. 1982, 5 4 , 224. (5) Ford, C. D.;Hurtubise, R. J. Anal. Chem. 1979, 5 1 , 659. (6) Chen, R. F.; Bowman, R. L. Science 1985, 147, 729. (7) Ford, C. D.; Hurtubise, R. J. Anal. Chem. 1980, 52, 656. (8) Seybold, P. G.; White, W. Anal. Chem. 1975, 4 7 , 1199. (9) Jakovljevic, I. M. Anal. Chem. 1977, 4 9 , 2048. (10) White, W.; Seybold, P. G. J . Phys. Chem. 1977, 8 1 , 2035. (11) Bower, E. L.;Winefordner, J. D. Anal. Chim. Acta 1978, 102, 1. (12) Meyers, M. L.; Seybold, P. G. Anal. Chem. 1979, 5 1 , 1069. (13) Ramasamy, S. M.; Hurtubise, R. J. Anal. Chem. 1982, 5 4 , 2477. (14) Scientific Polymer Products, Catalog 801, Ontarlo, NY; p 20. (15) Ostrowska, J.; Narebska, A. Colloid Po/ym. Sci. 1979, 257, 128. (16) Leyte, J. C ; Zuiderweg, L. H.; Vledder, H. J. Spectrochim. Acta, Part A 1987, 23A, 1397. (17) Pimentel, '3. C.; McClellan, A. L. "The Hydrogen Bond"; W. H. Freeman and Co.: San Francisco and London, 1960; pp 212-213. (18) Thompson, W. K ; Hall, D. G. Trans. Faraday SOC. 1987, 6 3 , 1553. (19) Schulman, E. M.; Parker, R. T. J . Phys. Chem. 1977, 8 1 , 1932. (20) McAleese, D. L.; Freedlander, R. S.; Dunlap, R. B. Anal. Chem. 1980, 52, 2443. (21) Von Wandruszka, R. M. A.; Hurtubise, R. J. Anal. Chem. 1977, 4 9 , 2164. (22) Von Wandruszka, R. M. A.; Hurtublse, R. J. Anal. Chim. Acta 1977, 9 3 , 331. (23) Aaron, J. J.; Kaleel, E.; Wlnefordner, J. D. J . Agric. Food Chem. 1979, 2 7 , 1233. (24) Pesce, A. J.; RosBn, C. G.; Pasby, T. L. "Fluorescence Spectroscopy"; Marcel Dekker: New York, 1971; p 96. (25) Nlshijlma, Y.; Teramoto, A.; Yamamoto, M.; Hiratsuka, S. J . Po/ym. Scl., Polym. P h p . Ed. 1987, 5 , 23. (26) Bayllss, N. S.;McRae, E. G. J . Phys. Chem. 1954, 58, 1002. (27) Reference 24, p 91. (28) Teale, F. W. J. Photochem. Photobiol. 1989, 10, 363. (29) Faucon, J. F.; Lussan, C. Biochim. Slophys. Acta 1973, 307, 459.
RECEIVED for review December 7,1982. Accepted February 10, 1983. Financial. support for this project was provided by the Department of Energy, Division of Basic Energy Sciences, Contract No. DE-AC02-80ER10624.
LITERATURE CITED (1) Parker, R. 1.; Freedlander, R. S.; Dunlap, R. B. Anal. Chim. Acta 1980, 720, 1.
Determination of Formation Constants of Calcium Complexes of Difluorornethylenediphosphonic Acid and Related Diphosphonates Tekum Fonong, Donlald J. Burton, end Donald J. Pietrzyk" Chemistry Depan'ment, The University of Iowa, Iowa City, Iowa 52242
Formation constants flor the 1:l and 2:l complexes formed between Ca2+and difluoromethylenedlptiosphonate(F,MDR), dichloromethylenediphiosphonate(CI,MDP), and i-hydroxyethane-1,l-diphosphonate(EHDP) were determlned by a potentlometric titratlon procedure where the lndlcator electrode Is a Ca Ion selective electrode and the titrant is a CaCI, solutlon. Measurements were made in the absence of lonlc strength controll and In 0.10 M NaCl and 0.10 M NH,CI solution. Logarithm formatlon constants, although similar for the 1:i complex, change In the order EHDP > CI,MDP > F,MDP which Is opposite to the1 acld strength exhlbited by the ligands. For the 2:l Ca*":llgand complex the EHDP complex formatlon constant is similar to the l:icomplex and the complex is signiflcantly more stable than either of the 2:l Ca2+:CI,MDP or Ca2+:F,MDP complexes. This Is probably the result of coordlnatlon Involving tlhe hydroxyl group In EHDP. Weak 1:i Na+:llgand complexes were also detected. 0003-2700/83/0355-1089$01 SO/O
Biological interest in pyrophosphates, compounds characterized by a P-0-P structure, increased when it was found that the presence of these substances in plasma and urine was capable of inhibiting the precipitation of calcium phosphate. Subsequent studies (1-4) have shown that diphosphonates of the type P-C-P 'we also valuable agents in the regulation of calcium metabolism in experimental laboratory animals and are potentially useful therapeutic agents for diseases involving abnormal calcification and excessive bone resorption. Success in clinical trials, for example, for osteoporosis, Paget's disease, and others, have been reported (1-6). Recent investigations have indicated high activities for the diphosphonates in inhibiting bone lysis of mice calvaria in vitro either cultivated in the presence of human tumors (7) or in tumor conditioned media (8),in rat tumor modeling studies (9),and in chemical trials for the treatment of malignant hypercalcemia (6). The two diphosphonates that have been studied extensively 0 1983 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 55, NO. 7, JUNE 1983
1090
Table I. Ionization Constants for the Diphosphonate Ligands acid
PKa,
F,MDPu C1,MDP EHDP,~ F,EDP
25 'C, 1.44 i. 0.15 1.70 f 0.19 1.87 f 0.14 1.07 f 0.96
PKa2 PKa, = 0.10 M (NaCl), NaOH Titrant 2.11 f 0.04 5.66 i. 0.02 2.13 f 0.10 5.66 f 0.04 2.76 f 0.05 6.78 f 0.02 1.97 k 0.09 4.49 f 0.04
7.63 k 0.02 8.30 i. 0.04 10.20 f 0.02 6.30 f 0.03
F,MDP C1,MDj' EHDP
25 "C, p = 0.10 M (Me4NC1),Me,OH Titrant 5.78 2 0.05 1.46 i 0.15 2.14 f 0.05 2.18 f 0.09 5.91 f 0.04 1.69 f 0.20 1.73 f 0.06 7.03 i 0.02 2.78 f 0.02
8.16 f 0.02 9.51 f 0.03 11.19 f 0.02
p
PKa4
a S e e r e f 1 7 ; < 2 . 6 , 5 . 8 0 , 8 . 0 0 r e p o r t e d f o r p K a 1 , p K a 3 , p K a 4 a t ~ = O ( 1 9 ) .1 . 7 f 0 . 2 , 2 . 4 7 f 0 . 0 4 , 7 . 2 8 i : 0 . 0 6 , 1 0 . 2 9 f
0.07 reported for 0.10 M KC1 solution (13). 5.89, 9.50 (20) and 6.11, 9.78 (19) reported for pKa3and pKa4 in 0.10 M Me4NC1and c1 = 0, respectively. 2.31, 6.99, and 10.93 ( 1 5 ) , 2.80, 7.00, and 11.16 (ZO), and 2.54 i. 0.05, 6.97 f 0.05, and 11.41 ?: 0.05 (11)reported for pKa2,pK,,, and pKa4 at c1 = 0.10 M Me,NCl (15, 20) and p = 0.50 M Me,NCl (11). are 1-hydroxyethane-1,l-diphosphonic acid (EHDP) and dichloromethylenediphosphonic acid (C1,MDP). The di-
H
/L
c
L
\
P(O)(OH)2 MDP
HO'
\P(O)(OH)2 EHDP
CI'
/
\P(O)(OH)2 C12MDP
F
\
P(0)(OH)2 F2MDP
phosphonates, which are structurally similar to the pyrophosphates, have the added advantages of increased stability toward chemical hydrolysis and enzymatic degradation and increased potency in the regulation of abnormal calcification in vivo (1-5). It has been suggested that complex formation between Ca2+ and the diphosphonates is one of several major factors that influence their biological activity. Several investigators have reported formation constants for complexes between EHDP and Ca2+as well as with other metal ions (10-16) determined via a potentiometric titration of EHDP with base in the presence of the metal ion. Apparently, no formation constant studies have been reported for C1,MDP. We have recently synthesized and characterized the compound difluoromethylenediphosphonic acid (F,MDP) (17). Properties of this acid suggest that it should also be a therapeutically useful diphosphonate and perhaps even be more efficient than the previously reported diphosphonates since the ClzMDP is often found to be more effective than EHDP (1-6). This report focuses on the determination of the formation constants for complexes formed between Ca2+and FzMDP and is part of a continuing study on the biological activity of F2MDP and other diphosphonates. The formation constants were determined via a potentiometric titration method employing a Ca ion selective electrode as the indicator electrode. Because of the interest in other diphosphonates, formation constants of Ca2+complexes of EHDP and ClzMDP were also determined by this method. EXPERIMENTAL SECTION Reagents. The procedure for the synthesis and characterization of FzMDP is provided elsewhere (17). EHDP and ClzMDP were obtained as disodium d t s from the Department of Dentistry, University of Iowa. Procedures for their purification and conversion to the free acid were similar to that used for FzMDP (17). All other salts used in the study were analytical reagent grade whenever possible. Tetramethylammonium (Me4N+)hydroxide and chloride were purchased from Curtin Matheson Scientific. Instrumentation. The Orion Model 901 Ionalyzer was used to determine pH (Orion 91-01 and 90-01glass and SCE electrodes) and pCa (Orion 93-20 and 90-01 calcium sensitive liquid ion exchange membrane and SCE electrodes). An Orion 96-11 sodium combination electrode was used for the determination of pNa. Conductance measurements were obtained with an Industrial
Instruments RC conductivity bridge and immersion type cell. Temperature control (25 0.1 "C) was maintained by pumping water from a Wac0 constant temperature bath through a water-jacket type titration vessel. A 10-mLDigipet obtained from Manostat was used as the buret for the titrations. Procedures. The K , values were determined by titration of known millimolar quantities of the diphosphonates as free acids and Na salts at known ionic strengths as described elsewhere (17). Calibration of the pH meter was with Clark and Lubs buffers. This titration procedure was also used to determine the percent purity of hydrated diphosphonates. Titration of 0.05-0.15 mmol of the diphosphonate in 150 mL of C02-freewater was carried out with a CaC12titrant of about 0.07 M. Its exact concentration was determined by a gravimetric procedure where Ca2+is precipitated as the oxalate salt. The pH, when required, was adjusted within the range of 8.5 to 10 with an NH3/NH4C1 buffer (about 0.006 M). Ionic strength was controlled where indicated by addition of either NaCl (0.10 M) or NH,Cl(O.lO M); since the buffer concentration is low the added electrolyte determined the ionic strength. The titrant was added in small aliquob and the potential was recorded when equilibrium was attained. About 5 min was required to reach equilibrium in the early stages of the titration; in later stages 30 s was sufficient. The equilibrium Ca2+ concentration was established from a carefully prepared calibration curve of potential vs. Ca2+concentration; the Ca2+standards used contained the same ionic strength and buffer pH and concentration as used in the titration. Reverse titrations, where Ca2+was titrated with the diphosphonate (ion-selective electrode) or where the diphosphonate-metal ion mixtures were titrated with NaOH (pH titration) and the conductometric Ca2+-diphosphonate titrations, involved a similar procedure. Formation constants and K , values were calculated by a computer program which provides a nonlinear least-squares fit of the data; the program (18)and its details are available on request. All constants reported here are the average of usually four to six individual determinations following many preliminary experiments which were used t o establish the optimum experimental conditions.
*
R E S U L T S AND DISCUSSION The dissociation of F2MDP as the free acid is shown in eq 1. Similar protonic equilibria can be written for the other
diphosphonic acids; a summary of the K, values for these acids is given in Table I.
ANALYTICAL CHEMISTRY, VOL. 55, NO. 7, JUNE 1983
L O
2 0
4.0 m l CaCt,
6 C
8.0
r l No04
Flgure 1. Titration of FJADP with a CaCI, titrant. I n (a) a 0.07400 M CaCI, tltrant was added to 53.0 mL of a p = 0.10 M (NH4CI), 25.0 f 0.1 OC, pH 9.50 solution while in (b) the solution also contained 0.1090 mmol of (NH,),F,IblDP.
The potential coordination sites of FzMDP are the hydroxyl groups and to a much lesser extent the phosphoryl oxygens. If the solution pH is maintained above pK,, then protonated complexes between Ca2+ and FzMDP should be negligible. Initial experiments were intended to establish the stoichiometry of detectable compdexes. Conductance titrations of CaC1, M FzMDP as the free (0.06 mmol in 50.6 mL) with 9.75 X acid provided detectable end points that corresponded to first the formation of a 2:1 Ca:L complex followed by a 1:l Ca:L complex as the titration with FzMDP was continued. For the reverse process, where 0.24 mmol of F2MDP in 50.0 mL was titrated with 0.1 M CaC12,only a 2:l Ca:L stoichiometry was clearly indicated. Potentiometric titration of Ca2+with FzMDP and the reverse titration were carried out with a Ca ion selective electrode to monitor Ca2+concentration and ia basic pH (pH >9.0) to ensure that F,MDP was in the L4- form. In the first case the equilibrium concentration of Ca2+steadily decreases as the Ca:L ratio changes from 2:l to 1:l when the FzMDP titrant, used as the (NH4+)4salt, is added. Shortly after exceeding the 1:l ratio, the Ca2+concentration levels off. For the reverse titration, where F2MDP as the (NH4+)4salt is titrated with CuC12, a gradual rise in Ca2+equilibrium concentration, which is significantly less than in the absence of the F2MDP, is observed. Depending on concentration, poorly defined end points appear in the 1:l and 2:l Ca:L ratios. A typical titration curve for the titration of FzMDP with a Ca2+ titrant is shown in Figure 1. Also shown in Figure 1 is a titration curve in the absence of F2MDP. Comparison of these two curves illusitrates the extent that FzMDP reduces the equilibrium Ca2+Concentration. The latter curve is also typical of a Ca2+calibration curve for the Ca ion selective electrode. Potentiometric titrations of Ca2+-Fz1MDP mixtures with NaOH indicated that protonated complexes also formed at lower pH values. In these experiments both pCa and pH were monitored simultaneously during the course of the titration. Figure 2 illustrates a typical titration curve where Ca:L is present in a 1:l ratio. Two end points are detected when monitoring the Ca2+concentration. The first Ca end point also corresponds to the end point for the titration of the first two acidic protons (see Table I for pK, values) of the ligand. This behavior in conskitent with the formation of a CaHzL complex for a mixed CaH3L+-CaH2L species initially and during the early stages of the titration. As NaOH is added and the pH approaches pK,,, or the point where the two most acidic protons are neutralized, a shift toward CaHzL in the equilibrium shown in eq 2 occurs. Continuation of the tiCa2+
+ H4L + CaH,L + 2H+
1091
(2)
Flgure 2. Titration of a 1:l Ca2+:H4F,MDP solution with a NaOH titrant. The analyte was a 173-mL solution that was 0.100 M in Ca2+ and H4F,MDP and the titrant was 0.06880 M NaOH; Ca2+ Is monitored in (a) while H+ is monitored In (b).
tration produces a second break in the Ca curve which corresponds to the neutralization of the third and fourth acidic protons (see Table I for pK,, and pK,). This indicates a gradual transition from the CaHzL to the CaHL- and finally CaL2- species. If the starting point of the titration is a 1:l mixture of Cazt:Na2HzL,only the latter part of the titration as shown in Figure 2 is observed. These two titrations also indicate a modest stability for a 1:l Ca:L complex since the Ca2+ can compete favorably with the protons for the basic sites. This combined pH-pCa titration with NaOH was repeated for the case where the F2MDPas the free acid was in the ratio of 2 1 Ca:L. As the NaOH titrant is added, the equilibrium Caz+concentration gradually decreases without any indication of a titration break until the end point for the fourth proton is reached. At this point the equilibrium Ca2+concentration levels off. This behavior, which is in contrast to Figure 2a where the Ca:L ratio is 1:1,indicates the presence of a major equilibrium of the type 2Ca2+
+ HzL + CazL + 4H+
(3)
A gradual neutralization of the protons as NaOH titrant is added causes a shift to the right, further decreasing the Ca2+ concentration in solution. These data also indicate a reasonable stability for the 2:l Ca:L complex. Previous workers (10,13-16) have used data obtained by titration of a mixture of Mn+ with MDP and EHDP with strong base to determine formation constants. Since pH changes during the titration, the constants obtained are for protonated species. Thus, we chose to titrate F,MDP as the Lp form in basic solution with Ca2+in order to avoid formation of the protonation complexes. The two-step equilibria illustrating the complex formation between FzMDP and Ca2+at pH >9 are shown in eq 4 and 5. Since the equilibrium Ca2+concentration can be obtained
from potentials recorded during the titration and a Ca2+ calibration curve for the ion-selective electrode, it is necessary to derive expressions relating the equilibrium Ca2+concentration with the formation constants and ligand concentration. Analytical concentrations for FzMDP (at pH >9, only L4is present for F,MDP) and Ca2+in terms of a mass balance are given by
CLb = [L4-] + [CaL2-] + [Ca2L]
(6)
1092
ANALYTICAL CHEMISTRY, VOL. 55, NO. 7, JUNE 1983
CCa2+= [Ca2+]+ [Cali2-]
+ 2[Ca2L]
(7)
The mean number of Ca2+bound per L" is defined by ii or
+ 2[Ca2L] [L4-] + [CaL2-] + [Ca2L] [CaL2-]
fi=
(8)
Combining eq 4 to 7 with 8 and eliminating Lp yields an equation in terms of ii, [Ca2+],and the formation constants K for eq 4 and 5 , or
+ 2KbLKg&[Ca2+] 1 + Kf;lL[Ca2+]+ KkLKd&[Ca2+] KbL[Ca2+]
fi=
(9)
Kijk
The quantity ii can also be defined in terms of measurable experimental parameters or
+
CoVo - [Ca2+](Vo V) fi=
cv
A summary of the formation constants determined for the 1:l and 2:l Ca:L complexes for the three ligands is listed in Table I1 for uncontrolled ionic strength and for an ionic strength of 0.10 M NaC1. For the latter case the buffer concentration was purposely kept small enough to provide reasonable buffer capacity while at the same time have little effect on the ionic strength; similarly the concentration of Ca2+and L4- has little effect on the ionic strength. Because of the differences in pK,, and pK,, between the three ligands (see Table I) protonated complexes are still present even though the pH is maintained at pH >8.5. Thus, designation of the formation constants in Table I1 follows the form
(10)
where Co is the concentration and V, the volume of the added Ca2+titrant and C and V correspond to the ligand concentration and volume, respectively. ClzMDP and EHDP will require a more basic solution (see Table I for pK, values) than used for FzMDP to ensure the L4- state, particularly for EHDP. Thus, the above equations must be modified for these two ligands in order to account for the contribution of protonated complexes. By use of eq 9 and 10, the formation constants can be calculated providing the quilibrium Ca2+concentration can be determined during the course of the titration. This was accomplished by using a computer program (18) that provides a nonlinear least-squares fit of the data. The program required input of the equilibrium Ca2+concentration and the corresponding volume of added Ca2+titrant to yield guesstimates for the formation constants which were subsequently refined by iteration. The equilibrium Ca2+ concentration was determined by potentiometric titration with a Ca ion selective electrode. The working range of the electrode was verified in the following way. Cell potentials were determined for a series of standard solutions of CaClz containing controlled ionic strength of 0.10 M (NaCl or NH4Cl) using a buffer a t the same concentration and pH as that used in the titration. Although not shown here the calibration curve was similar to that shown as curve a in Figure 1. The linear portion for a plot of cell potential vs. Ca2+concentration extended from 4.5 x 10-5 M to 1.00 X 10-1 M with a Nernstian slope of 30 mV/pCa. Below this concentration the slope decreases rapidly with decreasing Ca2+ concentration. Because of the curvature, Ca2+solutions below 10-5 M were not used in the calculations. The electrode was also shown to provide constant, reproducible potentials between the pH range of 4.50 to 11.00, At lower pH values, potentials increased abruptly due to electrode response to H+, while above pH 11 potentials slowly decreased as pH increased. Preliminary experiments demonstrated the concentration ranges for the Ca2+titrant concentration, ligand concentration, buffer concentration, and pH that would provide an optimum reproducibility for the measurable quantities. Furthermore, these experiments demonstrated the specific portions of the titration curve that provide the best data for the determination of the formation constants at an optimum accuracy. Although small variations are apparent, which depend on the ligand and the ligand-Ca titrant concentration, the titration curves obtained were similar to the ligand-Ca titration curve shown in Figure 2b. For calculation of the formation constants, the greatest number of potential measurements were recorded a t volumes surrounding the 1:l and 2:l Ca:L stoichiometric points.
= [M,(HjL)k][M]-i[HjL]-k
(11)
In the uncontrolled experiment the ionic strength varies from about 0.02 M to 0.04 M. Under these conditions formation constants are larger in comparison to the constants when using NaCl(O.10 M) to adjust the ionic strength. Several reports have suggested that Na+ forms weak complexes with EHDP (11, 16). Determination of the formation constants in the absence of Na+, where ionic strength and buffer are provided by NH4C1and NH4C1/NH3, respectively, are larger than for the Na+ solution and suggest a weak Na-L interaction. Constants for the NH4Cl solution are also shown in Table 11. Constants for the Na-L interaction were determined by a potentiometric titration procedure using a Na ion selective electrode as the indicator electrode. Ionic strength and pH were controlled by NH4Cl (0.10 M) and NH4C1/NH3buffer. Only a 1:l Na:L complex was detected for each of the three ligands. Furthermore, the formation constants are very small and appear to be identical for the three ligands; these data are summarized in Table 111. In previous studies on the acidity of F,MDP the titrant was NaOH and the ionic strength was controlled with NaCl (17). Since Na+ forms a weak complex with the ligand, its presence, particularly at the ionic strength used, could influence the K, values for the ligand. For this reason the K , values for F2MDP, C12MDP,and EHDP were determined in the absence of Na+ by using (Me)4NOHas titrant and (Me)4NC1for ionic strength control. The latter two ligands were also examined because the pK,, and pK,, for these ligands have not been reported before. The pK, values found for the three ligands in the absence of Na+ are also listed in Table I where they can be compared to pK, values found in the presence of Na+; previously reported pK, data from other laboratories are also included. The presence of the Na+ begins to exert its influence via mass action and its own competition with protons for the ionization sites as the acid strength of the acidic sites in the ligands decreases. Thus, the effect is the greatest on KQ and with EHDP, which is the weakest acid of the three ligands; the acid strength of the ligands follows the order FzMDP > ClzMDP > EHDP which is consistent with the electronegative influence of the groups attached to the methylene carbon. The diphosphonates will form complexes with other metal ions. This complexation was demonstrated by a NaOH titration of a solution containing a known concentration of the ligand in the H4L form and a known concentration of metal ion such that the M:H4Lis 1:l. Several typical titration curves with FzMDP as ligand are shown in Figure 3. In Figure 3a the titration curve is for F2MDP as the free acid in the absence of the metal ion. When the metal ion is present (Figure 3b-e), the coordination is indicated by a shift and/or more welldefined end points in the pH region corresponding to the pK,, and pK,, for FpMDP. A favorable competition of Mnf with the protons for the acidic sites and an apparent modest decrease in the pK, values for the protonated meta1:L species suggest complex stabilities similar or greater in magnitude to that found for Ca2+. Although not shown in Figure 3, FzMDP
ANALYTICAL CHEMISTRY, VOL. 55, NO. 7, JUNE 1983
1093
Table 11. Logarithm Formation Constants for CaZ+-DiphosphonateComplexesf liganda (NH,),F,MDP Na,F,MT)PC Na,H,F,MDPC ( NH, ), F,MDP Na,(NH,)HCl,MiDP Na,Cl,MDP Na,HCl,MDPC (NH,),CI,MDP~ (NH,),HC~,MDP~ Na,(NH,)HEHDP Na, HEHIDP (NH, ),HEHD;~
(",),F,EDP Na,F,EDPCpe (NH,),F,EDP~
PH 9.50 9.86 5.66 9.40 9.00 9.92 6.60 9.70 9.25 8.50 10.21 9.10 9.00 9.00 9.00
1% K1,lf
log KZOl 2.74 f. 0.08
log K,,,
4.86 i 0.06 4.23 i 0.04
1% KZl,
log Kl,, 2.98
4.36
t
0.03
5.05 f. 0.08
1.56
i: 0.03
2.89
f.
i
0.02
0.07
4.49 f 0.05 3.74 i 0.03 4.99
f
0.02
2.35 f 0.02 4.75 f 5.44 t 5.30 f 5.37
2.58
0.03 0.14 0.06 0.20
2.28 f. 0.04 4.43 i 0.33 2.46 i 0.06 4.26 i 0.30
0.02
1.26 f 0.10
1.76 f 0.03
1.35 f. 0.10
f
a Ligand concentration is about 1.00 X M , CaC1, titrant is about 8.87 X l o w 2M, MIL ratio covers the range of about 1:lto 3:1, and temperature is 25 "C. Uncontrolled ionic strength. Ionic strength = 0.10 M (NaC1). Ionic strength = 0.10 M (NH,Cl). e Complexes not detected. Previously reported formation constants for CaZ+-EHDPcomplexes are K,,, = 3.53 and K,,, = 6.40 at p = 0.10 M (Me,NCl) ( 1 6 ) , K,,,= 5.39 and K,,, = 4.48 at p = 0.10 M (Me,NCl) ( 1 4 ) ,K,,, = 6.04 and K,,, := 3.63 at p = 0.10 M (KCl) ( 1 3 ) ,and K,,, = 3.58 f 0.14 and K,,, :: 5.74 f. 0.10 at p = 0.50 M (Me,NCl) (12).
Table 111. Logarithm Formation Constants for NaZ+-DiphosphonateComplexes ligand (NH,),F,MDP" (NH,),C1,MDPa (NH,),HEHDP" (NH,),F,EDP"
PH 9.50 9.70 9.60 9.00
logK,,,, 1.29 f 0.01 1.28 i 0.03
1% Klll
1.28 i 0.03 0.85 i 0.03
a Ionic strength is 0.10 M (NH,Cl), temperature is 2.50 * 0.1 "C, ligand concentration is about 8.00 X lo-, M, NaCl titrant is 0,1000 M and Na+/L range covers a range of 2 : l to about 1 5 : l .
0.0
1.5
3.0
4 5
6.0
m L NoOH
Figure 3. pH titration of 1l:l M"+:H,F,MDP solution with a NaOH titrant. Curve (a)Is for H,F,MDP in the absence of MI'+ whlle (b) to (e) contain M"'; the anaiytes were 50.0-mL solutions that were 8.20 X M in M"+ and H,F,MDP and the tltrant was 0.03960 M NaOH.
coordination with Co2', Cd2+,Cu2+,Zn2+,Sc3+,and UO?+ was also demonstrated via this titration procedure. Although formatioin constants for complexes formed between EHDP and Ca2+and with other metal ions have been reported previously (IO-l6),little information about the exact structure of the com~plexesis available. Crystal structure studies on crystals of EHDP as the disodium, dihydrogen, tetrahydrate salt have been reported (21). These data indicate that no discrete molecules exist and the structure is one that comprises an intricate network of hydrogen bonds and ligand-bridged coordination complexes. Of particular interest is the location of the Na+ in the crystal. The coordination for one Na+ is nearly octahedral while the other appears to be nearly square pyramidal. Furthermore, the ligand acts as
a bidentate ligand for one Na+ and a tridentate for the other Na+ where coordination involving the C-OH group is indicated. In solution the stability of the 1:l and 2:l Ca:L complexes follows the order EHDP > C1,MDP > F2MDP. This order is consistent with the electronegative trend exhibited by the groups attached to the carbon. The differences are small except when comparing the 2:l Ca:L complexes. In this case the 2:l Ca:EHDP complex formation constant is much larger than for either the ClzMDP or FzMDP complexes and is similar in magnitude to the 1:l complex. This difference may be the result of the greater electronegativity of the halogens or, more likely, that the hydroxyl group is also participating in the coordination in solution. Evidence to suggest that Ca2+ coordination is via chelation involving both terminal groups rather than a t one terminal group was provided by studying (F4EDP) (22). For the ligand, (H0)2(0)PCFzCFzP(O)(OH)2 F4EDP only a 1:l Ca:L complex was detected and its formation constant is much smaller than that for the other diphosphonate ligands. As expected, the additional fluorines increase the acidity of F4EDP. Complexes between Na+ and F4EDP are barely detectable; however, when using NaCl for ionic strength control, the mass action influence of Na+ is large enough so that the Ca2+complexes with F4EDP are not detectable and it is only in uncontrolled ionic strength or in NH4Cl solution that very weak Ca2+complexes are detected. Summaries of the ionization and complex formation constants for F4EDP are listed in Tables I and 11, respectively. Registry Nu. EHDP, 2809-21-4; C12MDP,10596-23-3;F2MDP, 10596-32-4. LITERATURE CITED Francis, M D.; Centner, R. L. J . Chem. Educ. 1978, 55, 760-766. Russell, R. G. G.; Flelsch, H. Clln. Orthop. Relat. Res. 1975, 106, 24 1-263. Flelsch, H ; Russell, R. 0. G. J . Dent. Res. Suppl. 1972, 5 1 , 324-332. Khalrl, M. R. A.; Johnston, C. C . ; Altman, R. D.; Wellman, H. N.; Serafinl. A. N.; Sanksy, R. R. J . Am. Med. Assoc. 1974, 230,562-567. Stover, S. I-.; Hahn, H. R.; Mlller, J. M. Paraplegia 1976, 1 4 , 146-156. Jung, A. Am. J . Med. 1982, 72,221-226. Galasko, C. S.6.; Samuel, A. W.; Ruston, S.; Lacey, E. B r . J . Surg. 1980, 6 7 , 493-496. Jung, A.; Mermlllod, 8.; Barras, C.; Baud, M.; Courvoisier, B. Cancer Res. 1981, 4 1 , 3233-3237. Slrls, E. S.; Sherman, W. H.; Baquiran, D. C.; Schlatterer, J. P.; Osserman, E. F.; Canfieid, R. E. New Engl. J . Med. 1980, 302, 3 10-3 15. Irani, R. R.; Moedrltzer, K. J . Phys. Chem. 1962, 66, 1349-1353. Carroll, R. L.; Iranl, R. R. Inorg. Chem. 1967, 6 , , 1994-1998. Carroll, R. L.; Iranl, R. R. J . Inorg. Nucl. Chem. 1968, 30, 2971-2976.
Anal. Chem. 1983, 55, 1094-1098
1094
(13) Kabachnlk, M. I.; Lastovskll, R. P.; Medved, T. Y.; Medyntsev, V. V.; Kolpakova, I . D.; Dyatiova, N. M.; Dokl. Chem. (Engl. Trans/.) 1987, 777, 1060-1063. (14) Gravenstetter, R. J.; Cliley, W. A. J. Phys. Chem. 1971, 75, 676-682. (15) Wada, H.; Fernando, Q . Anal. Chem. 1971, 4 3 , 751-755. (16) Wada, H.; Fernando, Q . Anal. Chem. 1972, 4 4 , 1640-1643. (17) Burton, D. J.; Pletrzyk, D. J.; Ishihara, T.; Fonong, T.; Flynn, R. M. J . FlUOfhe Chem. 1982, 20, 617-626. (18) Gordon, G. “Executive Computer Programs as Applied to Handling Chemical Data” (Modified by K. Sando); University of Iowa: Iowa City, IA, 1972. (19) Blackburn, G. M.; England, D. A.; Kolkman, F. J . Chem. Soc., Chem. Comm~n.1961, 930-932. (20) Grabenstetter, R. J.; Quimby, 0. T.; Fiautt, T. J. J. Phys. Chem. 1967, 7 7 , 4194-4202.
(21) Barnett, B. L.; Strlckland, L. C. Acta Crystallogr.,Sect. 6 1979, 835, 1212-1 214. (22) Maruta, M., unpublished work, University of Iowa, 1982.
RECEIVED for review November 22, 1982. Accepted March 3, 1983. This work was presented a t The 184th American Chemical Society Meeting, Kansas City, MO, Sept 1982. This investigation was supported by Grant AM 28077 awarded by the National Institute of Arthritis, Diabetes, Digestive, and Kidney Diseases. D.J.B. would like to thank NSF and AFOSR for support of organofluorine research a t the University of Iowa.
Extraction of Platinum and Its Separation from Palladium by Polyurethane Foam Sargon J. AI-Bazl and Arthur Chow* Department of Chemistty, University of Manitoba, Winnipeg, Manitoba, Canada R3T 2N2
The Influence of thiocyanate, hydrochloric acld, and pH of the solutlon on the extraction of platinum( I I ) and platinum( I V ) by polyether-type polyurethane foam was demonstrated. Dlstrlbution coefflclents of more than I O 4 were obtalned. The effect of the chlorlde salts of dlfferent cations on the extractlon of platinum( I I ) Increased In the order of LI’ < Na+ < K+ 5 NH4+ lndlcatlng that platinum( I I )-thlocyanate complex [most probably Pt(SCN):-] Is extracted through the “catlonchelation” mechanlsm. The separatlon of platinum( I V ) and palladium(11) was also studied and the resuns Indicated that 20-fold excess of platlnum has no effect on the extractlon of palladlum.
Several solvent extraction methods have been used for the preconcentration of platinum from aqueous solution including the use of methyl isobutyl ketone (I),tri-n-butyl phosphate (2), tri-n-octylamine hydrochloride in toluene (3), 4-(5nony1)pyridine in benzene ( 4 ) , and di-n-octyl sulfite in cyclohexane ( 5 ) . Two different mechanisms have been suggested, the “ion-association” mechanism (6-9) in which platinum is extracted as an ion pair of the type 2C+-PtX82or C+.HPtX6- (C is a bulky cation or a protonated organic solvent, X is an anion) and the extraction of platinum as a neutral species of the type PtX,L, (X is a halide or pseudohalide, L is a neutral ligand) by several organic solvents (10-1 2). For the separation of platinum and palladium, Forsythe et al. (13)used solvent extraction with methyl isobutyl ketone and Gregoire and Chow (14) employed silicone-rubber foam treated with dimethylglyoxime. The anion exchanger Amberlite IRA-400 (15)and the cation exchanger Amberlite IR400 (16) have both been used to separate Pd(I1) and Pt(1V). The purpose of the present work was to find the optimum conditions for the preconcentration of platinum [as Pt(I1) and Pt(IV)] from thiocyanate solutions by polyether-type polyurethane foam. The study also involved the separation of platinum and palladium and the mechanism by which platinum(I1)-thiocyanate complex is distributed between the foam and the aqueous phase.
EXPERIMENTAL SECTION Apparatus and Reagents. A Model 306 Perkin-Elmer atomic-absorption spectrometer was used for platinum and palladium determinations, a Fisher Accumet Model 520 for pH measurements, and a Varian 634 UV-visible spectrophotometer for absorbance measurements. Platinum and palladium solutions were made from K2PtC1, and Na2PtC&.6H20(Johnson-MattheyChemicals, Ltd.) and PdClz (Matheson, Coleman and Bell) in 0.1 M hydrochloric acid. All other chemicals used were of analytical grade. Polyether-type polyurethane foam (no. 1338 M) was obtained from G. N. Jackson Ltd. (Winnipeg, Manitoba) and washed by using the procedure previously reported ( 1 7 ) . A 5.0 M stock solution of potassium thiocyanate was prepared in doubly distilled and deionized water. Extraction Procedure. Aliquots of the stock solutions of Pt(I1) or Pt(1V) chloride and potassium thiocyanate were diluted to 100 mL with distilled water. Any addition of acid was made after diluting the thiocyanate solution in order to minimize the decomposition of thiocyanate caused by direct contact with concentrated acid (18). The extraction of 3 X lod M platinum by 99 1 mg foam and the calculations of percentage extracted (%E) and distribution coefficient (D)were done as described previously (17).
RESULTS AND DISCUSSION The Extraction of Platinum(I1). T o establish the optimum conditions for the extraction of Pt(I1) from thiocyanate solutions with polyurethane foam, the effect of various parameters was investigated. The extraction of platinum from solutions 0.5 M in thiocyanate and 0.4 M in hydrochloric acid was studied from 5 min to 30 h. The extraction increased sharply up t o 15 min and then slowly until 45 min, beyond which it was almost time independent. A minimum time of 45 min was regarded convenient for further studies. T o determine the optimum range of thiocyanate for the extraction of Pt(I1) from 0.8 M hydrochloric acid solutions, the effect of thiocyanate concentrations varying from zero to 0.9 M was studied. At zero thiocyanate, less than 1% of platinum was extracted and the UV spectrum showed a characteristic absorbance a t 230 nm, which has been assigned to PtC1:- by several workers (19-21). For solutions containing thiocyanate, the UV spectra showed a broad absorbance below 330 nm partially overlapped with the intense absorbance of
0003-2700/83/0355-1094$01.50/00 1983 American Chemical Society
