1972
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
are shown in Figure 6, (a) and (b), which were obtained by column B before and after hydrogenation reaction followed by thermal decomposition of about 200 pg of P E at 650 "C. In the hydrogenation mode, the carrier gas was changed to hydrogen and a catalyst tube (id., 3 mm X 15 cm) packed with 10% Pt coated on Diasolid L (80-100mesh) was inserted between the pyrolyzer and the inlet tube. Both the catalyst and the inlet tube were maintained a t 200 'C. Further work on the high-resolution pyrograms of various polyolefines is currently in progress.
(2) K . Grob and G. Grob, J . Chromatogr. Sci., 7, 584 (1969). (3) K. Grob and G. Grob, J . Chromatogr. Sci., 7, 587 (1969). (4) J. P. Schrnid, P . P. Schrnid, and W. Simon, Chromatographia, 9 , 597 ( 1 976). (5) H. L. C. Meuzelaar, H. G. Ficke, and H. C den Harink, J . Chromatogr. Sci., 13. 12 (1975). (6) J. W. de Leeuw, W. L. Maters, D. v. d. Meent, and J. J. Boon, Anal. Chem., 49, 1881 (1977). (7) S. Tsuge and T. Takeuchi. Anal. Chem., 49, 348 (1977). (8) R. G. McKeag and F. W. Hougen, J . Chromatogr., 136, 308 (1977). (9) A. Mitchell and M. Needlernan, Anal. Chem., 50, 668 (1978).
LITERATURE CITED
RECEIVED for review July 25, 1978. Accepted September 12,
(1) G Schornburg, R Dielman, H Husmann, and F Weeke, J Chromatogr , 122, 55 (1976)
1978.
Stability Constants of Calcium and Lanthanide Ions with Murexide K. S. Balaji, S. Dinesh Kumar, and P. Gupta-Bhaya" Department of Chemistry and The Biosystems Laboratories, Indian Institute of Technology, Kanpur 2080 16, U.P., India
pH titration less suitable as compared to spectrophotometric methods. However, for these determinations, one requires accurate values of stability constants at well defined pH for metal-murexide equilibria. It is well known that the positions of equilibria are pH-dependent (3). In this paper, we report values of stability constants for murexide complexes of Ca2+,Gd3+,Eu3+,Tb3+,and La3+ a t well defined pH, maintained by buffer solutions (pH < 6, acetate; pH > 6, phosphate). A method has been developed to take into account the binding of metal ions to buffer ions (acetate and phosphate). This binding reduces the concentration of free metal ion. The values so determined differ significantly from older values reported in literature, where metal ion-buffer ion binding was ignored.
The stability constants of the metallochromic indicator murexide (Ammonium Purpurate) with Caz+, Eu3+, Gd3+, La3+, and Tb3+ have been determined at several well defined pHs maintained by buffer solutions. The binding of buffer ions to metal ions has been taken into account in the analysis of the data. I n the case of Caz+, two different methods have been used to determine the stability constants and the results agree very well. The stability constant values are significantly different from the values published in literature, because in the earlier determinations the binding of buffer ions to metal ions was neglected. As expected, the discrepancy between the older values and ours is larger under conditions where the binding of buffer ions to metal ions should be stronger. For the trivalent lanthanide ions, our values differ from the older values by several orders of magnitude.
EXPERIMENTAL All reagents used were of AnalaR grade. La, Gd, and Tb were purchased as their trichlorides (99.99% pure) from Indian Rare Earths Corporation, and were used as such. Eu203(99.970pure) was purchased from Sigma Chemical Company and was converted to its perchlorate by repeated treatment with perchloric acid. Water used was doubly distilled and deionized. Spectrophotometric determinations were made with a precalibrated Beckman DU spectrophotometer. Murexide solutions were prepared in the respective buffers. Their concentrations were determined spectrophotometrically at 506 nm ( t = 1.26 X IO4 mol-' cm-' L (3),this value was determined in our laboratory). The purity of the murexide used was greater than 99.9% according to microanalytical data obtained in our laboratory. Metal solutions were standardized volumetrically against standard EDTA ( 4 ) . In all measurements an appropriate quantity of NaC104or KC1 was added to maintain the total ionic strength at the desired values. All measurements were made at the temperature specified. The titrations were carried out in the following way. Method A. Metal solution was added in small steps using a microliter pipet to a buffered solution of murexide. After each addition, the absorbance at 470 nm (480 nm for Eu3+)was noted with murexide solution in the reference cell. The choice of wavelength is dictated by the maximum in the difference spectrum. Method B. Murexide solution was added in small steps, using a microliter pipet, t o a buffered solution of metal ion. After each addition. the absorbance at 470 nm (480 nm for Eu3+)was noted
The equilibrium between metal ions and the metallochromic indicator murexide (Ammonium Purpurate) has been used to determine the stability constants of metal-ligand equilibria by monitoring spectrophotometrically the displacement of metal-murexide equilibrium due to metal-ligand binding. Some of the most important metal-ligand equilibria involve biological macromolecules as ligands. Calcium ion is one of the most important metal ions in biochemistry (1). Trivalent lanthanide ions are useful spectroscopic probes for calcium binding sites ( 2 ) . This makes the stability constants for binding of calcium and trivalent lanthanide ions to biological macromolecules important. A detailed analysis of titration data on binding of metal ions t o biomolecules provides us with equilibrium constants of individual sites and free energy of coupling between mutually interacting binding sites in a macromolecule. T o carry out such an analysis, one needs the concentrations of free metal ions as a function of the total concentration of metal ions. The titration has to be carried out under well defined conditions of temperature, ionic strength, and, in particular, pH because macromolecular structure and binding are strongly pH sensitive. This requirement makes alternative methods like 0003-2700/78/0350-1972$01.00/0
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1978 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
Table I.
1973
Values ( m o l - ' c m - ' L ) for Metal-Murexide Systems system PH
Ca-murexide 9.10 x 9.22 x 9.23 x 9.30 x 9.40 x 9.40 x 9.47 x 9.55 x 9.60 x 9.62 x 9.65 x 9.71 x
4.00 4.50 4.6jb 5.00 5.40 5.50 6.00 6.50 6.8OC 7.00 7.50 8.00
103
103 103 103 103 103 103
Eu-murexide
Gd-murexide
Tb-murexide
La-murexide
1.07 X 10' 1.10 x l o 4
1.30 X 10' 1.32 X l o 4
1.27 x 104 1.28 x 10'
9.94 x 103 9.96 x i o 3
1 . 1 5 X 10' 1.17 x 104 1.19 x 104
1.35 X 10" 1.39 X l o 4 1 . 4 0 X 10'
1.30 X 10" 1.33 X 10' 1.36 x 104
LOO x i o 4 1.03 X 10' 1.05 x 104
...
103
...
103 103 103
...
...
... ...
103
At 470 nm (480 nm for E u ) and
y
...
...
...
...
... ...
...
I . .
..I
... ... ... ...
...
...
... ...
...
...
...
...
= 0.10. Identical values were obtained at 25 "C and 30 "C.
with water in the reference cell. The path length was 1 cm. The buffer solutions used were CH,COONa-CH,COOH at pH 6. The concentrations used are mentioned in the Tables. Determination of A€. Le is the difference in the molar absorptivity of murexide and the metal murexide complex at the particular wavelength. Murexide solutions of known concentration were saturated to convert all the murexide to the complex. For example, 10.' M murexide solution was saturated with M metal ion. Saturation was confirmed by the fact that increasing the concentration of metal ion did not alter the absorbance. Measuring the absorbance and knowing the concentration of Metal Murexide present, molar absorptivity of the complex, and hence the value of A€, was calculated.
-
-
RESULTS AND DISCUSSION T h e data were analyzed in the following way. Method A. Let us denote the molar absorptivity of metal murexide as C M M ~the , molar absorptivity of murexide as tMu, the differential absorbance observed (with murexide as blank) as AA, the concentration of total murexide (mol L-') as CTMu, the concentration of free murexide (mol L-') as CMu,and the concentration of metal murexide complex (mol L-I) as CMhfu. Then it follows ( 5 ) :
+
(using CThfU= C M u CMMusince murexide forms 1:1 complexes only ( 6 ) ) ,where A C = ~ M -MtMu. ~ Hence, if one knows the value of A( a t a particular wavelength, one can calculate the concentration of the metal murexide complex in solution, by measuring the differential absorbance AA a t that wavelength. Le measurements were carried out for all metals that are considered in this paper. and are presented in Table I. Let us suppose that buffer binds to metal and forms two species, MB and MB2 with equilibrium constants K 1 and &. The metal ions investigated do not form higher species under conditions employed (7-9). However, the method can be easily generalized to include their effect. When CTM= total concentration of metal, C1 = concentration of MB species in solution, C2 = concentration of MB2 species in solution, and CMM"= concentration of MMu species in solution; then the concentration of free metal C M = CTv (CMbiu + C1 + C 2 ) . Let the equilibrium constant of the reaction MMu + M + Mu be K . Then
-
Now, we may approximate the concentration of free buffer
0.01.
y
= 0.2.
Table 11. p K d i s s n Values for Metal-Murexide Systems (Method A )
system Eu-murexide La-murexide Tb-murexide Gd-murexide Ca-murexide Ca-murexideb Ca-murexidec Ca-murexided
PKQ 5.442 i 0.009 4.577 z 0.008 5.016 r 0.009 5.031 i 0.008 2.803 t 0.006 2.780 r 0.006 2.741 i 0.005 3.352 i 0.005
pK values reported in literature ref. 4.18'
3.43e 3.95' 4.08e
13 13 13 13
. I .
... 2.68f
...
3
a In acetate buffer (0.1 M) at pH 5.40 and y = 0.10 (maintained by NaC10,) at 30 i- 0.2 "C. I n acetate buffer (0.01 M ) , pH 5.40, y = 0.01. In acetate buffer (0.01 M), pH 4.65, p = 0.01. In phosphate buffer (0.07M),pH6.80,y=0.20. eAt12"C,y=0.1(KN0,). f In 0.1 CaCll at 25 "C, pH = 4.65. e Average of three measurements, each measurement being average of six titration points.
= concentration of total buffer (denoted by CH). This is a good approximation since concentration of buffer >> concentration of metal (0.1M) (1 x 10-4 M)
Even if all the metal ions bind to buffer, our approximation is valid. Then
cl = Kl c2 = P 2
CB (CBJ2
cd (CThf - c h l h f u - c, - c2)
(CTM
ClIhfu
C1 -
On rearrangement. we obtain c1
where
= 71 (CThl - CMhlMu), and 7 . are given by
Y1 = Y2
-
K = -C M C M ~
y =
I t follows that
=
c2 = 72 (CTM - C M M u )
Kl C B___ (1 + KICB + $?(cB)') 132(c~);' (1 + KLCB
+ /32(CBI2)
(3) (4)
1974
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
Table III.a pKdiss Values of Metal-Murexide Systems as a Function of pH PH 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 8.0
Eu-murexide
Gd-murexide
system Tb-murexide
5.339 5.382 5.420 5.553
4.902 4.958 5.002 5.130
4.889 4.940 4.982 5.125
0.010 0.006 0.006 0.005
i i i t
0.009 0.007 0.005 0.006
La-murexide
Table IV. pKdiss Values of Calcium-Murexide System (Method B ) pKa conditionsb 2.732
i
0.009
pH = 4.65, (acetate buffer 0.01 M),
2.799
i
0.013
pH = 5.40 (acetate buffer 0.01 M),
2.802
*
0.014
3.347
i
0.010
pH = 5.40 (acetate buffer 0.10 M), p = 0.10 pH = 6.80 (phosphate buffer 0.07 M), p = 0.20
p
= 0.01
p =
0.01
NaC10, was used a Average of three measurements. for maintaining specified total ionic strength. Temperature 30 i 0.2 "C. From Equation 1, C M is known; ~ ~ CTMis known and 7 can be calculated. Knowing these quantities CM can be determined. T h e direct substitution of CMMU,CA4u,and CM values in Equation 2, gives the equilibrium constant K at each titration point. T h e values of K thus determined are given in Tables I1 and Table 111. Method B. Equations 2 and 5 together give KMMu
=
(1 - Y)(CTM- C M M ~ ) ( C T-MC~M M J
CMM~
Let KMhlu/(l- y) = K'. If the experimental condition is such that CTMu 6 , phosphate 0.05 M. Temperature 25 i 0.1 "C. i
0.008
t
0.009
rivation is given in the Appendix.) The K values determined by this method for the Ca2+murexide system (Table IV) agree very well with the K values determined by the Method A. For the rare earth ions, this method could not be applied because one of the assumptions, necessary for the validity of the method, Le., K' >> CTMu breaks down. In the case of lanthanides, K' -lo-' M and we use CTMu of the order of lo4 M. (Lower CTMuwould make titration data unsatisfactory.) The metal-buffer stability constants used in this work were obtained by modifying the values reported in literature (7-9) to include the effect of different conditions of temperatures and ionic strength used in our work. The effect of ionic strength was calculated by evaluating the activity coefficieiits from the equation log f = -0.509z'pL"' at 25 "C, (where f is the activity coefficient. z is the charge, p is the ionic strength). The effect of temperature was calculated by using the van't Hoff isochore. The W values were taken from the literature (7, 9, 11, 12). The W value for Tb3+-acetate equilibrium is not reported in the literature. We took the mean of the AH for Gd3+-acetate and Dy3+-acetate equilibrium. These AH values were determined a t p = 2.0. By noting the variation of lH with ionic strength for Cd2+,we found that the use of this 1H value at p = 0.1 introduces an error which is less than the standard deviation quoted. The literature values of these equilibrium constants and the enthalpy changes with the references and the corrected values used in our work are presented in Table V. It is worth noting that the standard deviations are somewhat larger under lower pH conditions. This is perhaps due to instability of murexide a t lower p H (3). Our values differ significantly from the values published in literature (3,13). This difference is due to neglect of "inert"
Table V. Values of Metal Ion-Buffer Ion Stability Constants Used in This Work literature values A H , , a kcal
system Ca-a ce tat e Ca-phosphate
log K ,
logp?
1.24c 1.69gC''
...
ref.
...
9
mol-' 0.91 3.3
1.408C'f
...
9
3.4
7
corrected values used in this work
AN^,^ kcal mol-'
ref.
...
...
7 9
...
9
log K , 1.884,' 1.8953, 2.987,' 3.027,' 3.56Qk 2.052,' 2.093; 2.360' 2.328,: 2.3451 2.183,: 2.206: 2.093,: 2.125' 2.047,' 2.0741
log Pz I . ,
... . . I
11 3.952,: 3.9921 Eu-acetate 2.31d 3.061h 3.91d 8 1.4 12 3.801,: 3.840', 3.245 Gd-acetate 8 1.868 3.76d 2.16d Tb-acetate 2.07d 3.66d 8 2.396' 3.842e ... 3.699,: 3.737: 12 3.354; 3.308' 3.786 La-acetate 2.02d 3.26d 8 2.181 a M + L + ML. CaHPO,. M + 2L ML,. p = 0.1 (NaCIO,), T = 20 "C. e Ca2++ HPO,?. p = 0, T = 25 "C. Ca2++ H,PO,- e CaH,PO,+. Mean of A H , values of Gd3'-acetate and Dy"-a,cetate (ref. 1 2 ) values. Mean of A H , T = 30 "C, p = values of Sm'+-acetate and Gd3+-acetate (ref. 1 2 ) values. T = 25 "C, p = 0.10, 3 T = 30°C, p = 0.10. 0.20.
*
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
1975
Rearranging Equation A1 and dividing by Ch141u,
buffer ions in previous studies. The binding of the "inert" ions t o metal ions is, in reality, very significant. Recently, in at least two publications (14, Is),this factor has been stressed. It should be pointed out that since murexide can exist in several protonated forms, the stability constant values determined in this paper are effective equilibrium constants. As such, they are expected to depend on pH. Our data on stability constants of Ca2+and lanthanide ions with murexide as a function of pH are summarized in Table 111. A generalization of our method can take into account the effect of binding of cations of the buffer solution to murexide. We ignored this effect because Na+ does not bind significantly to murexide.
Substituting from Equations A4 ,and A5 into A6 and rearranging, we get,
c r s This suggests a linear plot of o,'[(Crr,,/(~) - lis]cs. a,which gives (CT,f1t) = I (intercept on the cr axis). - C ' T b l l t / K ' = t (slope). Therefore K ' = - I / t . This method, without consideration of buffer binding to metals, has been used in other laboratories (10).
APPENDIX Method B. Substituting for C1 and C2 from Equations 3 and 4 in Equation 2,
LITERATURE CITED R H. Krestsinger and D. J Nelson, Coord Chem Rev , 18. 29 (1976) R . J. P. Williams, 0.Rev. Chem. Soc., 24, 331 (1970). G. Swarzenbach and H. Gysling, Helv. Chim. Acta, 32, 1314 (1949). I . M. Kolthoff and P. J. Elving, Ed., "Treatise on Analytical Chemistry", Part 11, Vol. 8. Interscience Publishers, New York and London (1963). S. Pal, M.Sc. Thesis, I.I.T. Kanpur (1977). S . P. Sanyal, Chemist-Analyst, 55, 104 (1966). G. H. Nancollas, J . Chem. Soc., 744 (1956). R . S. Kolat and J. E. Powell, Inorg. Chem.. 1, 293 (1962). A . Chughtai, R. Marshall, and G. H Nancollas, J . Phys. Chem.. 72, 208 (1968) T. M. Jovtn. unpublished notes, Max Planck Institut fur Biophysikalische Chemie, Gottingen. G. R. Choppin and J. K. Schneider, J . Inorg. Nuci. Chem., 32,3283 ( 1 970). I Grenthe. Acta Chem. Scand., 18, 283 (1964). V. G. Geiger, Ber. Bunsenges., 69. 617 (1965). H. Hauser, C. C. Hinckley, J. Krebs, B. A . Levine, M. C. Phillips. and R. J , P. Williams, Bioch/m B0phys. Acta, 468, 364 (1977). E. A-Noack and E. M. Heinen. Eur. J . Biochem., 79, 245 (1977). Interunion Commission on Biothermcdynamtcs, J. Bioi. Chem.. 251, 6879 ( 1976).
h''CMhlu. Rearrangement of Equation A1 then gives CMMU
=
CTMCTM~ K' + CTXl
RECEIVED for review November 18, 1977. Accepted August 17. 1978. The authors thank the Council of Scientific and
A plot of
N
vs.
Industrial Research, New Delhi, and the Board of Rexarch in Nuclear Sciences. Department of Atomic Energy, Bombay, for financial support. One of the authors (K.S.H.) thanks N.C.E.R.T., New Delhi. for a National Science Talent Search Fellowship held during the tenure of this work. A part oft his work was submitted in partial fulfilment of the Degree , ) f Master of Science of K.S.B. at I.I.T. Kanpur.
CT31u gives
CTM-lC
K' +
CTM
= S, as slope
Determination of Acid-Base and Solubility Behavior of Lignite Fly Ash by Selective Dissolution in Mineral Acids John B. Green' and Stanley E. M a n a h a n " Department of Chemistry, 123 Chemistry Building. University
of Missouri, Columbia, Missour/ 6520 7
As part of an evaluation of coal humic acid-fly ash mixtures as scrubber media for the removal of sulfur dioxide from stark gas ( I ) , methods described in this paper were developed for the general characterization of fly ash with special emphasis upon acid neutralizing qualities. LVith the exception of elemental analysis. such methods have not heen generally available, which may be a factor in the low utilization of fly ash. Recently there has been a great deal of interest in thc analysis of trace elements in fly ash (3,as indicated hy the
A method has been developed for the chemical analysis of fly ash which gives the total available base in the ash. This method is based upon the dissolution of fly ash in mineral acid accompanied by analysis of major metal ion constituents and sulfate. The dissolution occurs in discrete steps with increasing acidity. This behavior is indicative of specific fractions within the ash. P r e s e n t address, B a r t l e s v i l l e E n e r g y T e c h n o l o g y Center, U.S. D e p a r t m e n t of Energy, P.O. Box 1398, Bartlesville. Okla. 74003. 0003-2700/78/0350-1975$01 O O / O
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1978 American Chemical Society
