Anal. Chem. 3904, 56,749-752
may be adjusted by adding acid or base. Such increased handling is less desirable; however, use of a greater buffer capacity solution in such samples will only increase the matrix problems already introduced by high Ca concentrations. Quartz and other samples with very low metal concentrations are also more difficult to analyze. In these instances, flameless AAS could be used for all elemental determinations or a smaller aqueous:solid ratio could be introduced. For those instances where matrix problems are introduced by the buffers proposed here (Table 11) or due to an increased buffer capacity system, standards can be prepared in the buffer. We found that AAS analysis of the leachates should be carried out within 1-2 weeks since precipitation and discoloration occurred in some pH 5 and 6 samples after 4-6 weeks. No differences in concentration were found for any samples analyzed over a 2 week period following leaching. Sediment and soil samples are as variable as the many leaching techniques that have been used to study them. When the processes being investigated are pH related, we encourage this approach to leaching or one of the recommended modifications, realizing the importance of knowing the final pH and being sensitive to differences in sample composition. Registry No. Cu, 7440-50-8; Cd, 7440-43-9;Fe, 7439-89-6;Mn, 7439-96-5; Pb, 7439-92-1; Zn, 7440-66-6.
LITERATURE CITED Goldberg, E. D.; Arrhenius, G. 0. S. Geochim. Cosmochlm. Acta 1958, 73, 153-212. Carroll, D.; Starkey, H. Clays Clay Miner., R o c . 7th Nafl. Conf. 1960, 80-101. Gibbs, R. J. Geol. SOC.Am. Bull. 1977, 88,829-843. Bascomb, C. L. J. SoliScl. 1966, 79, 251-268. Arshad, M. A.; St. Arnaud, R. J.; Huang, P. M. Can. J. SoilSci. 1972, 52, 19-26. McKeague, J. A.; Day, J. A. Can. J. Soil Sci. 1966. 46, 13-22. Coffin, D. E. Can. J. SoilSci. 1963, 4 3 , 7-17.
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(8) Malo, 8. A. Environ. Sci. Techno/. 1977, 7 7 , 277-282. (9) Presley, B. J.; Kolodny, Y.; Nissenbaum, A,; Kaplan, I. R. Geochim. Cosmochim. Acta 1972, 3 6 , 1073-1090. (IO) Rosholt, J. N.; Emiliani, C.; Geiss, J.; Koczy, F. F.; Wangersky, P. J. J. Geol. 1961, 69, 162-185. (11) Hirst, D. M.; Nlcholis, 0.D. J. Sediment. Petrol. 1958, 28,461-468. (12) Chester, R.; Hughes, M. J. Chem. Geol. 1967, 2 ,249-262. (13) Sorensen, R. C.; Oelsligie, D. D.; Knudsen, D. Sol Sci. 1971, 7 7 7 , 352-359 - - - - - -. (14) Kltano, Y.; Sakata, M.; Matsumoto, E. Geochim. Cosmochim. Acta 1980, 44, 1279-1285. (15) Agemian, H.; Chau, A. S . Y . Ana/yst (London) 1976, 707,761-767. (16) Sinex, S. A,; Cantillo, A. Y.; Helz, G. R. Anal. Chem. 1960, 52, 2342-2346. (17) Luoma, S. N.; Bryan, G. W. J. Mar. Bioi. Assoc. U.K. 1978, 58, 793-802. (18) Luoma, S. N.; Bryan, G. W. I n "Chemical Modeling-Speciation, Sorption, Solubility, and Kinetics In Aqueous Systems"; Jenne, E. A., Ed.; American Chemical Society: Washington, DC, 1979; pp 577-61 1. (19) Gambrell, R. P.; Khalid, R. A.; Patrlck, W. H., Jr. Environ. Sci Technoi. 1980, 74, 431-436. (20) Van Valin, R.; Morse, J. W. Mar. Chem. 1982, 7 7 , 535-564. (21) Baas-Becking, L. G. M.; Kaplan, I.R.; Moore, 0. J. Geol. 1960, 68, 243-284. (22) Krug, E. C.; Frink, C. R. Science 1983, 227,520-525. (23) Owen, G. I n "Physiology of Mollusca"; Wilbur, K. M., Yonge, C. M., Eds.; Academic Press: New York, 1966; pp 53-88. (24) Barnard, E. A. I n "Comparative Animal Physiology"; Prosser, C. L., Ed.; W. 8. Saunders: Philadelphia, PA, 1973; pp 138-139. (25) Mlller, T. G.; MacKay, W. C. Wafer Res. 1980, 74, 129-133. (26) Clark, W. M.; Lubs, H. A. J. 6/01.Chem. 1916, 25, 479. (27) Trefry, J. H.; Shokes, R. F. I n "Marine Environmental Pollution, Vol. 2"; Geyer, R. A., Ed., Elsevier: Amsterdam, 1981; pp 193-208. (28) Trefry, J. H.; Presley, B. J. Geochim. Cosmochim. Acta 1982, 1715-1726. (29) Vetter, T. W.; Trefry, J. H.; Metz, S. Trans. Am. Geophys. Union 1983, 250.
RECEIVED for review October 3, 1983. Accepted January 3, 1984. This investigation was funded by the National Oceanic and Atmospheric Administration, Project P-PRIME (Pollutant-Particle Reactions in the Marine Environment), Contract No. NA82RAC00036.
Fluoride Standards in Determination of Equilibrium Constants of Metal Ion-Fluoride Complexes G . T. Hefter* School of Mathematical & Physical Sciences, Murdoch University, Murdoch, Western Australia 6150, Australia
C. B. Chan and N. H. Tioh Chemistry Department, University of Malaya, Kuala Lumpur 22-11, Malaysia
Precise measurements of the stablilty constants of the monofluoride complexes of LI', Ag', Na', and Pb2+have been made In aqueous nltrate and perchlorate medla. These results show that the expected advantage of replacing NaF by the less assoclated KF as the source of fluoride Ion in the study of metal Ion-fluoride equliibrla is largely nullified by having to substitute KNO, for NaCIO, as the ionic medium. Literature data lndlcate that apparent fluorlde stablllty constants measured In nltrate media are often lower than those In perchlorate media as a result of metal Ion-nitrate complexlng.
The study of metal ion-fluoride equilibria M"+(aq)
+ nF-(aq) + MF,("-")+(aq)
(1)
continues to attract considerable experimental attention because of the intrinsic interest of such complexes as model electrostatic systems (1,2) and their importance in geochemistry (31, dentistry (41, oceanography ( 5 ) ,etc. The study of fluoride equilibria requires an appropriate fluoride ion standard. The almost universal choice for this purpose has been sodium fluoride. However, in establishing activity standards for the fluoride ion selective electrode (ISE), Robinson et al. (6, 7) found that NaF was appreciably more associated (ion paired) than KF and thus recommended adoption of the latter as the fluoride activity standard. Potassium fluoride is now available as a standard reference material for this purpose from the U.S. National Bureau of Standards (NBS, SRM 2203). The findings of Robinson et al. (6, 7) have been independently confirmed (8,9)and have important implications for the study of fluoride equilibria. Not only is the fluoride
0003-2700/84/0356-0749$0 1.50/0 0 1964 American Chemlcal Society
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Table I. Typical Potentiometric Data for the Determination of Fluoride Stability Constantsa System Na’lF-; I = 1.00 M (KNO,); T = 25.00 “C Ej,C mV 104[F-]T,M d 104[F-],“*f M [Na’], M F,( X), f , g M “ -Eobsd> mV 40.3 0.53 4.498 4.395 0.1071 0.218 38.0 0.83 4.198 4.066 0.1667 0.194 32.8 1.44 3.599 3.400 0.2857 0.205 28.5 1.91 3.149 2.932 0.3750 0.197 24.7 2.28 2.799 2.569 0.4444 0.202 18.6 2.83 2.290 2.066 0.5455 0.199 16.0 3.03 2.099 1.882 0.5833 0.198 Averagea (12 pts): bI(NaF(aq))= 0.200 i 0.006 M ” a Abbreviated summary: full set contained 12 data points. Observed cell potential. Calculated from the Henderson equation ( I 7): Ej = -59.16 log (1- 0.191[Nat]). Total fluoride concentration corrected for background fluoride impurity ( 1 5 ) , data not shown. e Calculated from the Nernst equation: E,,, = E“’ - 59.16 log [F-I, where E,,, = Eobd Ej. E”’ = 239.40 I0.10 mV determined by calibration ( 1 5 ) , data not shown. Calculated iteratively (15). g As defined in eq 2. -~
ISE undoubtedly the most popular tool for investigating such equilibria but if NaF is appreciably associated, equilibrium constants for reactions of the type shown in eq 1 when determined in Na+-containing media will only be “apparent” rather than “true” (i.e., depending only on the ionic strength, and not upon the nature, of the medium) constants because of competition for F- between Na+ and Mm+. Thus, it appears that fluoride equilibria (eq 1)would be better determined by using K F as the fluoride ion source. There is, unfortunately, a problem in adopting KF as the fluoride standard for equilibrium studies. Equilibrium constants are nowadays almost always measured in solutions of constant ionic strength (maintained by the addition of a “swamping electrolyte”) to minimize activity coefficient variation: the so-called “constant ionic medium principle” (10, 11). Although a diversity of electrolytes have been employed for “swamping” purposes, perchlorates are usually preferred because of their limited tendency to form complexes with the (metal) ion of interest. However, if KF is used as the fluoride source, the sparing solubility of KC104 (12)precludes the use of a perchlorate as the swamping electrolyte. Although there are other potentially suitable salts such as KPF,; from the point of view of cost, ease of purification and convenience, KNOBis the obvious choice of swamping electrolyte for use in conjunction with KF. Since nitrate is known to be more strongly complexing than perchlorate (13),any advantage gained by replacing NaF by K F may be lost by having to replace NaC104 by KNOBas the swamping electrolyte. Because the vast bulk of fluoride equilibria have to date been determined in NaF/NaC104 media, it is clearly important to establish the likely errors in existing data and the possible advantages of a switch to KF/KN03 media. As NaF is only slightly associated (6-9) and as NO3- is usually weakly complexing (13),any effects will be best observed either for weak fluoride complexes or for those metal ions which form unusually strong complexes with NO3-. For this purpose a detailed study of the fluoride complexes of Li+, Na+, Ag’, and Pb2+was undertaken in NaC104, KN03, NaN03, and (for one system) KC104 media. The technique chosen for this investigation was fluoride ion selective electrode potentiometry for both its convenience and accuracy (14, 15). EXPERIMENTAL SECTION Reagents. Lithium perchlorate and nitrate were prepared by neutralizing L.R. grade lithium carbonate with the appropriate A.R. grade acids. The crude products were recrystallized twice from water. Sodium nitrate, potassium nitrate, and potassium perchlorate (A.R. grade) were also recrystallized twice from water. Sodium fluoride (A.R. grade) was further purified by the method of Lingane (16),but potassium fluoride (NBS, SRM 2203) was used as received, AU salts were thoroughly dehydrated and dried mmHg) for several days. Standard at 110 OC under vacuum
stock solutions were prepared by weighing the appropriate salts and diluting to volume in calibrated volumetric glassware. Glass-distilled water was used throughout. After the pH was adjusted to 5.5 (to minimize HF formation and OH- interference to the fluoride ISE) all solutions were stored in preleached airtight polyethylene containers. Apparatus and Procedures. Apparatus and procedures for determining equilibrium constants were standard and have been described in detail elsewhere (14,15). The potential of the fluoride ISE (Orion,Model 94-09) was measured to hO.1 mV with a digital voltmeter (Orion, Model 901A). The reference electrode was a silver wire in contact with solid silver chloride. Electrochemical cells were analogous to those employed previously (14,15). Liquid junction potentials (LJPs) were estimated from the Henderson equation which can be done reliably for the constant ionic medium type of liquid junction used (17). The conductivity and transference data required for these calculations (17) were obtained from the literature. All measurements were made at 25.00 f 0.02 “C.
Calculations. The equilibria in aqueous metal ion-fluoride solutions, assuming only “mononuclear”complexes to be formed, are represented by eq 1. Under conditions used in this work, [M”+] >> [F-] and, assuming activity coefficients are constant at constant ionic strength, the following equation may be derived (14, 15):
where p, are the overall concentration stability constants corresponding to eq 1, the subscript T denotes total or analytical concentration, and Fo(X)is introduced for convenience to represent the experimentally determinable left-hand side of eq 2. Since [M”+] >> [F-] and the metal ions in the present study have a weak affinity for the fluoride ion, only one complex is likely to form in each case and so Fo(X) = PI. Higher order complexes, if present, will be detected by a dependence of F & X ) on [F-] (eq 2). RESULTS AND DISCUSSION Results obtained for a typical potentiometric titration are shown in Table I. The constancy of the Fo(X)values indicates only one complex MF(aq) is being detected over the entire concentration range. Results were independent of metal ion concentrations indicating no “polynuclear” complexes were present (IO). Table I1 summarizes the equilibrium constants obtained in the present study. All values quoted are the average of at least three, and usually more than five, independent titrations. With the exception of PbF+(aq) all the complexes are extremely weak (& N 1 M-l). The Effect of Replacement of Na+ by K+ i n t h e Ionic Medium. In order to study the effect of replacing Na+ by K+ in the ionic medium, uncomplicated by anion changes, a
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Table 11. Summary of 0 Values Obtained for Lit-, Na+-, Ag+-, and PbZ+-FluorideComplexes in Various Media at 25 "C' ionic strength ion o,, M-I and medium, M Lit 0.75 = 0.05 1 (NaClO,) 0.58 i 0.06 1 (NaNO,) 0.93 i 0.07 1 (KNO,) N a+ 0.20 r 0.02 1 (KNO,) 0.47 i 0.01 0.1 (KC10,) n.d. 0.1 (KNO,) Ag' 0.30 i 0.03 1 (NaClO,) 0.10 i 0.02 1 (NaNO,) 0.41 i 0.04 1 (KNO,) PbZ 25i lb 1 (NaCIO,) 12.6 i 0.1 1 (NaNO,) 13.5 I0.2 1 (KNO3) a Present work; uncertainty limits are standard Taken from ref 1 4 ; deviations; n.d., not detected. obtained under identical conditions to the present work. +
series of measurements were made in NaN0, and KNO, media. The results (Table 11) show P1(MF("-')+(aq)) values are always larger in K N 0 3 media consistent with a higher degree of association of NaF. This may be seen more clearly by writing the equilibrium of interest as a competition reaction for F- (X = Na+ or K'): Mm+(aq)
+ XF(aq) + MF(m-l)+(aq) + Xf(aq)
(3)
(There is a subtle point here. It might be thought that complexation of F by X+ would result in a higher apparent value of P,(MF) because there is less free F- to be detected by the F- electrode. The reason this is not so is because the X F complexation is not allowed for in the calibration of the electrode. If this effect is taken into account it can be shown that P1(MF)apparent &(MF)true/(l+ PxF[X+I)and the apparent constant is always lower than the true value. This effect is discussed more fully elsewhere (18).) Some intrinsic variation in the observed equilibrium constants in the two media is expected on the basis of activity coefficient effects, differences in L J P corrections, etc. However, the fact that the (relative) effect of replacement of Na+ by K+ is greater the weaker the fluoride complex (compare Ag', Li+, and Pb2+) suggests the difference in association of NaF and KF is the important factor. The Effect of Replacement of C10, by NO3- in the Ionic Medium. This effect is best seen by comparing the data of Li+, Ag+, and Pb2+ in NaC10, and NaN03 media (Table 11). For each metal ion there is a significant decrease in ,f31(MF(m-1)+(aq)) when C10, is replaced by NO3- in the ionic medium, consistent with a competition reaction (Y = C104or NO3-) MY("-l)+(aq)
+ F-(aq) * MF("-l)+(aq)
f Y-(aq)
(4)
The magnitude of the effect on Pl(MF("-l)+(aq)) correlates with the known strength of the MN03(m-1)+(aq)complexes (13). The effect of replacement of C10; by NOs- was qualitatively confirmed in KC104 and KNOBmedia by measurements on the association of NaF. Thus &(NaF(aq)) was readily determinable in I = 0.1 M (KClO,) but, despite considerable experimental effort, no association of NaF could be detected in I = 0.1 M (KNOB)(suggesting &(NaF(aq)) P1(KN03(aq)) under these conditions). The Relative Merits of NaF and KF as Fluoride Standards for Equilibrium Studies. From the above discussion it is clear that the relative merits of NaF and KF as fluoride standards for the study of metal ion-fluoride equilibria (eq 1)must be considered in terms of all the possible competing equilibria. That is, combining eq 1, 3 and 4 and
751
omitting charges for simplicity MY(aq)
+ XF(aq) + MF(aq) + XY(aq)
(5)
When written in this way it is clear that the observable stability constant (equilibrium 5) may differ markedly from the required "true" value (equilibrium 1)depending on the relative magnitudes of the various competing reactions. It might be expected that, for measurements employing the constant ionic medium approach, the advantage of replacing NaF by KF (see above) as the fluoride standard may be largely nullified by having to use K N 0 3 rather than NaC104 as the swamping electrolyte. Thus, even for the very weak fluoride complexes of Li+ and Ag+, the &(MF(aq)) values obtained in the various media fall in the order NaNO, < NaC104 C KN03, with the difference between the NaC10, and KNO, values barely outside the limits of experimental precision (Table 11). Furthermore, evidence from conductivity (8)and potentiometry (9) suggests that even KF is slightly associated and thus P,,(MF,("-")+(aq)) values obtained in KF/KN03 may still not be strictly "truen values. For example, measurements in I = 1 M (&NO3) give a value of &(LiF(aq)) = 1.17 f 0.06 M-' (9) which is appreciably higher than the value obtained in KNOB(Table 11). The potential problems associated with the use of nitrate media are most clearly demonstrated by the P b 2 + / Fsystem. This system was chosen deliberately because of the relatively strong complexation of Pb2+by NO, (&(PbN03+(aq)) 15 M-l (13)).In the present context PbF+(aq) is a moderately strong complex; ,nevertheless, the replacement of C104- by NO, in the ionic medium results in a dramatic change in the apparent value of P1(PbF+(aq))which far outweighs the effect of replacement of Na+ by K+ (Table 11). In summary, the present results suggest that there is a slight advantage in using KF for fluoride equilibrium studies only under special circumstances, Le., for very weak fluoride complexes or when no swamping electrolyte is present (or when KC104can be used). However, if a swamping electrolyte (other than KC104) is present there seems to be little to be gained and sometimes a lot to be lost in switching from NaF/NaC104 to KF/KN03 media for the determination of fluoride equilibria. An important corollary to this finding is that despite the association of NaF, P,(MF,(m-n)+(aq)) values obtained in NaF/NaC104 media should be sufficiently close to "true" equilibrium constants (Le., true for the particular ionic strength used) for most purposes (3-5). Unfortunately, as very few fluoride equilibria have been measured in non-Na+-containing media (13,19), there are insufficient data to prove this assertion. However, those data which are available (8)support it. For example, log Pl(HF) = 3.31 in both I = 3 M NaC104 and KC1; log P1(FeF2+(aq))= 5.15 and 5.18 in I = 0.5 M NaC104 and KN03, respectively, etc. (19). The Use of Nitrate Media for the Study of Fluoride Complexes. Although the use of K F / K N 0 3 media may be advantageous for some systems (see above), nitrate media are not generally the optimum choice for the study of fluoride equilibria. This point is made clear by Table I11 which summarizes all the available data obtained under comparable conditions in nitrate and perchlorate media containing the same cation (19). These data show that with few exceptions equilibrium constants obtained in nitrate media are lower than the corresponding values in perchlorate media. This is consistent with the known strength of metal ion nitrate complexes (13). Such constanb are of value only as "apparent" constants for the specific medium in which they are determined: they cannot be assumed to be directly applicable to other media such as seawater (5),geological brines ( 3 ) ,etc. Certainly if nitrate media are to be used at all for the study of fluoride equilibria, KNO, (or, even better, RbNO, or CsN03) is to be
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Anal. Chem. 1984, 56,752-754
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Table 111. Comparison of Metal Ion-Fluoride Stability Constants in Perchlorate and Nitrate Mediaa log 0 (MF(m-‘)-(aq)) perchlorate nitrate H‘ 2.95 2.91‘ Mg2 1.35d 1.31d Ca2 0.58d 0.57 Sr2+ 0.13d -0.05 Ba2 -0.19d -0.40 ZnZT 0.74 ( I = 0.5) 0.49 ( I = 0.5) Cd” 0.48 ( I = 3) 0.46 ( I = 3) CU2+ 0.71 ( I = 0.5) 0.73 ( I = 0.5) Np4 4.65d (HClO,) 4.23 (“0,) Pu4+ 4.20 (HC10,) 4.00 (”0,) 1n3 3.76 ( I = 0.5) 3.72 ( I = 0.5) a Unless otherwise indicated all data have been taken from ref 1 9 and refer to I = 1 M NaCIO, and NaNO,. Average of a large number of independent determinaAverage tions. Reference 20; see also reference 19. of two independent determinations.
ion
+
+
+
preferred over the more popular NaN03. Registry No. NO
