add this to the separatory funnel. Agitate and discard the aqueous phase. Combine the acetate fractions and remove the ethyl acetate by evaporating on an asbestos-covered hot plate. If the presence of one or more interfering ions is known or suspected, and additional extractions (see Table I) are necessary, dissolve the uranium residue left from the acetate evaporation with four 5-ml. portions of 1N nitric acid and transfer to a 60-ml. separatory funnel. Add 20 grams of aluminum nitrate and 10 ml. of ethyl acetate and extract as described above, repeating the stripping extraction. Evaporate the ethyl acetate and rrcycle again if another extraction is indicated. After the final extraction and evaporation of ethyl acetate, add 5 ml. of 70% perchloric acid to the uranium residue and fume to dryness. When cool, dissolve and transfer the residue to a suitable volunietric flask with small portions of concentrated hydrochloric acid making the solution to volume with additional hydrochloric acid. Do not heat to dissolve the residue. The absorbance of the solution is mrasured a t 246 mp against a concentrated hydrochloric acid blank. Remove the cell holder promptly after each measurement to prevent corrosion of the cell compartment. RELIABILITY
The success of this procedure depends on the isolation of uranium in a
relatively pure form. For optimum results considerable care must be exercised in the extraction process. Reliability data for the recommended procedure are given in Table 11. The data show the effects of u p to four extraction cycles in removing the strongly interfering mercury(I1) ion. Acceptable results are obtained after three extractions. The standard deviation is less than 2% at the 20-p.p.m. level. LITERATURE CITED
(1) Buck, R. P., Singhadeja, S., Rogers, L. B., ANAL.CHEM.26, 1240 (1954). (2) Cheng, K. L., Ibid., 30, 1027 (1958). (3) Delahay, P., “Instrumental Analysis,” pp. 207-8, Macmillan, New York, 1957. (4) Grimaldi, F. S., May, I., Fletcher,
M. H., Titcomb, J., “Collected Papers on Methods of Analysis for Uranium and Thorium,” Geological Survey Bulletin 1006, pp. 125-31, U. S. Government Printing Offire, Washington, D. C.,
1954. (5) Hillebrand, W. F., Lundell, G. E. F., Bright, H. A,, Huffman, J. I., “Applied Inorganic Analysis,” pp. 46&71, Wiley, New York, 1955. (6) Holcomb, H. P., Yoe, J. H., ANAL. CHEM.32, 612 (1960). (7) Horton, C. A., White, J. C., Ibid., 30, 1779 (1958). (8) Katz, J. J., Seaborg, G. T., “The Chemistry of the Actinide Elements,” pp. 171-203, Methuen, London, 1957. (9) Legge, D. I., ANAL.CHEM.26, 1617 (1954). (10) Maeck, W. J., Booman, G. L.,
Elliott, M. C., Rein, J. E., Zbid., 31,
1130 (1959). (11) Meloan, C. E., Holkeboer, P., Brandt, W. W., Ibid., 32, 791 (1960). (12) Merritt. C.. Hershenson. H. M..’ ’ Rogers, L.’B., Zbid., 25, 572 (1953). (13) Moeller, T., “Inorganic Chemistry,” pp. 598-901, Wiley, New York, 1952. (14) Moeller, T., Brantley, J. C., ANAL. CHEM.22, 433 (1950). (15) Motajima, K., Yoshida, H., Iyowa, K., Zbid., 32, 1083 (1960). (16) Paige, B. E., Elliott, M. C., Rein, J. E., Ibid., 29, 1029 (1957). (17) Rodden, C. J., editor, “Analytical
Chemistry of Manhattan Project,” pp. 122-35, McGraw-Hill, New York, 1950. (18) Rodden, C. J., ANAL. CHEM. 25,
1598 (1953). (19) Rodden, C. J., J. Research Natl. Bur. Standards 26. 557 (1941): , , 28. , 265 (1942). (20) Sandell, E. B., “Colorimetric,, Determination of Traces of Metals, 3rd ed., pp. 83, 900, Interscience, New York, 1959. (21) Seaborg, G. T., Nucleonics 5 (No. 5), 16 (1949). (22) Stewart, D. C., U . S. Atomic Energy Comm. Rept. ANL-4812, (February 1952). (23) Vosburgh, W. C., Cooper, G. R., J . Am. Chem. SOC.63, 437 (1941). (24) Yoe, J. H., Jones, A. L., IND.ENQ. CHEM.,ANAL.ED. 16, 111 (1944). (25) Yost, D. M., Russell, H., Garner, C. S., “The Rare-Earth Elements and ~
Their Compounds,” Wiley, New York, 1947.
RECEIVEDfor review April 6, 1961. Accepted July 10, 1961. Division of Analytical Chemistry, 139th Meeting, ACS, St. Louis, Mo., March 1961.
Increased Selectivity in Chelometric Titrations through End Point Location by Linear Extrapolation Copper as a Photometric Indicator DAVID A. AIKENS, GABY SCHMUCKLER,’ FAWZY S. SADEK? and CHARLES N. REILLEY Department o f Chemistry, University of North Carolina, Chapel Hi//, N. C. b The inherent selectivity of end point location by linear extrapolation permits selective determinations to be carried out when the difference in effective chelonate stability constants of soughtafter and interfering ions is as small as 2 log K units. The necessary masking often can b e provided by careful choice of chelon and buffer systems. Photometric titrations with a Cu(ll) indicator are presented as an example. Ca is determined in the presence of Mg in 0.5M ammonium hydroxide a t pH 10 with ethylene glycol bis(&xninoethyl ether)-N,N’tetraacetic acid. Selective determinations of Cd and of Zn in the presence of the other are carried out with diethylenetriaminepentaacetic acid under the same conditions. Analytical and 1664
ANALYTICAL CHEMISTRY
empirical methods for the determination of appropriate solution conditions are presented.
HE problem of insufficient selecT t i v i t y in chelometric titrations stems in large measure from the method of end point detection generally employed. I n particular, this refers to end points determined potentiometrically or by means of visual indicators. For equal concentrations of soughtafter ion and interference, the use of a visual end point requires that the effective chelonate stability constant of the sought-after ion exceed that of the interfering ion by 5 to 6 log K units, while the difference necessary for a suecessful potentiometric end point is about
4 log K units. The degree of masking required to provide the necessary difference in chelonate stability is often difficult t o provide. I n contrast, end point location by linear extrapolation of points far from end point, such as the amperometric, conductometric, or photometric method, is applicable when the difference in effective chelonate stabilities is only 2 log K units. This difference corresponds to formation of 1% of the less stable chelonate a t the point of 50% formation of the more stable chelonate, so that extrapolation to the end point 1 Present address, Department of Chemistry, Technion, Haifa, Israel. * On leave from National Research Center, Cairo, U. A. R.
I
I
1
2
0
EGC
3
I
C f CHELON
‘,rL;j\TS
Figure 1. Effect of sequence of metal chelation reactions on utility of copper as a photometric indicator One equivalent each of MI, Mz,Cu
Path 1-3-5 1-4-5 1-2-4-5 1-2-5
Sequence of Chelation CU, MI, Mz CU f MI, MI MI, CU, hr.2 Mz, CU M2
is possible with negligiblp loss of accuracy. An effective stability constant difference of 2 log K units often can be obtained simply by careful selection of the chelon and buffer system. The inherent selectivity of linear estrapolation techniques of end point location greatly broadens the scope of chelometric methods. Determinations which are difficult with conventional end point detection methods often can be carried out easily when the end point is located by linear extrapolation. It is the purpose of this paper to demonstrate the applicability of linear extrapolation techniques as exemplified by photometric end point detection and to present empirical and analytical methods for the determination of appropriate titration conditions. Although photometric end point detection is discussed in detail, the same approach is applicable to any linear estrapolation tcchnique. A11 give the same increase in selectivity and the choice of a particular linear estrapolation technique would depend on the experimental circumstances. COPPER AS A PHOTOMETRIC INDICATOR
The choice of photometric end point detection raises the question of a snitable indicator and several possibilities arise. The use of extremely selective metallochromic indicators is a possibility, although not a likely one when the interfering ion is similar to that being determined. The ultraviolet absorbance of a titrate ion, the free chelon, or a metal chelonate may be monitored as suggested by Malmstadt and Gohrbandt (S),but simplicity suggests the addition of a metal ion which undergoes a change in visible absorbance on chelation. Copper was chosen, as its chelo-
+
nates are stable and the chelation reaction is readily reversible. The marked tendency of copper to form complexes with other ligands offers wide control of the effective chelonate stability, so that the absorbance change accompanying the chelation of copper can be made to occur a t the proper point in the titration. ilpplications of copper as a photometric indicator in chelometric titrations can be divided into two broad categories. I n the first, the absolute stability constant of the copper chelonate is less than that of the titrate ion and the absorbance change occurs only after the titrate ion is completely rheIated. Examples are the use of copper in titrations of ferric ion (11) and of bismuth (12) by Underwood and in the titration of thorium by Malmstadt and Gohrbandt (3). I n the second category, which is not so well known, the absolute stability constant of the copper chelonate is greater than that of the titrate ion, but the effective stability of the copper chclate is reduced to a value below that of the titrate ion by adjustment of solution conditions. Although the absolute stability constant of Cu-EDTA is 18.8 log units and that of Ca-EDTA is only 10.6 log units (9),Ramiah and Vishnu (6) employed copper as a photometric indicator in EDTA titrations of calcium in 1.5M ammonium hydroxide. I n this medium, the effective stability constant of Ca-EDTA is 10.6 log units, while that of Cu-EDTA is reduced to 5.8 log units. These examples effectively illustrate the versatility of copper as a photometric indicator in chelometric titrations, but the most promising application, a n indicator in selective titrations, has not been explored. Copper, when
added to a binary mixture of metal ions, may be utilized as a photometric indicator in two distinct ways, depending 011 the sequence in which the metal chelonntes are formed. The significance of this sequence is rradily apparent from examination of Figure 1, in which are presented several poqsible sequences for the cheletion of copper and AI, and Mz, the two ions of the original misture. hfl is assumed to form a more stable chelonate than hIL. The sequence of chelonate formation reactions in turn depends on the effective chelonate stabilities as determined by solution conditions. The first mode of indicator action involves the chelation of copper after the chelation of MI and before the titration of Mz, only M I being determined. This corresponds to line 1-2-4-5 of Figure 1 and is illustrated by the determination of calcium in the presence of magnesium with ethylene glycol bis(8aminoethyl ether)-N,N’-tetraacetic acid (EGTA). The second mode of indicator action permits the selective determination of both &I1and MZin a single titration. &Il is titrated first, followed by the simultaneous chelation of RI. and copper. The titration follows path 1-2-5 of Figure 1. The amount of AI1 is determined from the first break in absorbance and the amount of hIz from the amount of titrant consumed between the first and second absorbance breaks minus the known amount of copper indicator added. This mode of indicator action is illustrated by the simultaneous determination of cadmium and zinc with dicthylenetriaminepentaacetic acid (DTPA). where the chelation of cadmium occurs first and is followed by the simultaneous chelation of copper and zinc. EXPERIMENTAL
All chemicals were specified as analytical reagent grade, except diethylenetriaminepentaacetic acid (DTPA) and ethylene glycol bis(paminoethy1 ether)-N,N’-tetraacetic acid (EGTA), which were used as obtained from Lamont Laboratories, 5002 Mockingbird Lane, Dallas, Tex. Demineralized water was used in the preparation of all sohtions. Chelon stock solutions were made up approximately 0.lM and standardized against calcium carbonate or copper metal. Metal ion solutions, 0.1M, were standardized against standard EDTA. Appropriate dilutions were used in the titration experiments. Stock ammonia buffer (pH 10) was 3.OM in ammonium hydroxide and ea. 1d.I in ammonium nitrate. Apparatus. All absorbance measurements were obtained with a Cary Model 14 recording spectrophotometer. T h e cell holder was replaced with a 400-ml. borosilicate glass beaker which served simultaneously Reagents.
VOL 33, NO. 12, NOVEMBER 1961
1665
as titration vessel and absorption cell.
1 magnetic stirrer provided agitation. A wood cell conipartnient cover provided entry for the stirrer leads and a 10-m!.microburet. The buret belowthe stopcock and all surfaces inside the cell compartment were flat black to minimize stray light. When the initial absorbance of the sample solution \+as greater than about 0.7, attenuation of the relerence with a solution having an absorbance slightly loner than that of the sample resulted in more precise absorbance measurements. Titration Procedure. Synthetic mixtures were prepared. the proper amount of concentrated buffer was added, the volume adjusted t o about 300 ml., and t h e usual technique of photometric titration employed. The adjustment of titration conditions is described in detail below. Calcium was determined in t h e presence of magnesium with EGTA iu solutions containing 0.5H ammonium hydroxide a t p H IO. Titrations nere performed a t a wave length of 550 mp using 0.4 mmole of copper indicator. Selective determinations of zinc and cadmium Were carried out with DTPA in solutions of 0.5M ammoniuni hydroxide a t p H 10 a t a wave length of io0 mr. The copper indicator was equal to 0.1 mmole or to the amount of zinc, whichever was greater. Metal Chelonate Displacement Reactions. Aliquots of 0.1 mmole each of copper, calcium, and E G T A or of 0.1 mmole each of copper, cadmium, and D T P A were taken and diluted t o 300 ml., and t h e p H was adjusted t o 10 with sodium hydroxide. T h e absorbance was monitored using t h e titration apparatus described previously, as aliquots of p H 10 ammonia buffer were added. The DTPA chelonate system was studied a t io0 nip and the EGTA chelonate system a t 550 mp. Chelonate Derivative Studies. T h e formation constants of ammonia derivatives of copper chelonates in 1.OM ammonium nitrate were determined by measurement of t h e absorbance of t h e copper chelonates a s t h e free ammonia concentration was varied. T h e free ammonia concentration was determined using a p H meter which had been calibrated previously in 1.ON ammonium nitrate solutions containing known amounts of ammonium hydroxide. Solutions were made up by taking 0.25 m o l e of Cu(II), 0.1 mole of ammonium nitrate, and 2.5 mmoles of E D T A or 12.5 mmoles of DTPA or EGTA. The p H of chelon solutions were adjusted to the approximate p H of the final solution prior to addition to the other constituents. The free ammonia concentration was adjusted approximately by addition of ammonium hydroxide, the solution diluted to 100 ml., and the free ammonia concentration determined by p H measurement. The absorbance was measured differentially in 2-cm. cells at 720 mfi for the EDTA and EGTA chelonates and a t 650 m p for the DTPA chelonate. The most transparent solution ill each series was placed in the reference beam. Measurements were carried out in a n air-condi1666
ANALYTICAL CHEMISTRY
tioiied rooni niaintained a t 25.0" 0.5" C.
I
FACTORS AFFECTING CHELONATE STABILITY
Knowledge of the stabilities of metal chelonates and the factors upon which they depend is essential to the rational application of chelometric methods. The necessity is accentuated mith the problem of selective titrations. Proper choice of chelon, pH, and buffer system composition is a prerequisite to obtaining optimum results. The numerous factors affecting chelonate stability may appear bewildering at first, but they can be treated in a simple, systematic manner. Detailed analysis and methods of calculation have been presented ( 7 ) . As a knowledge of these factors is essential in the determination of suitable titration conditions, a brief summary is presented here. -4s a first approximation, the stability of a metal chelonate is given by its absolute stability constant, which is simply the equilibrium constant for the reaction M f Y s M Y (1) where M is the aquated metal ion, Y the unprotonated form of the chelon, and MY the metal chelonate. The simultaneous existence of M and Y as major representatives of their respective species is unlikely and the absolute stability constant is a rather poor criterion of the extent of metal chelonate formation. Much more informative is a n effective stability eonstant which can be derived from the absolute stability constant by consideration of the other competitive equilibria which involve M, Y, and MY. Chelons are proton acceptors and the hydronium ion competes with metal ions for the chelon, reducing the effective stability of the chelonate. Metal ions forming very stable chelonates, such as mercury(I1) ion and iron(II1) ion, can be titrated in acid solution, while those forming weak chelonates such as calcium must be titrated in alkaline solution. The effective stability of a chelonate increases with p H until the fraction of the free chelon in the unprotonated form is essentially unity. Further increase in p H does not increase the stability of the chelonate and may lead to undesirable side effects. The concentration of free metal ion generally decreases with increasing p H because of stepwise hydrolysis, which eventually culminates in the precipitation of insoluble hydroxides. It is often possible to mask certain ions by deliberate precipitation, permitting selective titration of others. The composite effect of pH on the concentration of free chelon and free metal ion generally results in maximum chelonate stability at a n intermediate pH, but
of much IOU (T magnitude than might be anticipated fyom thP abpolute stability constant. This variation in effective stability oftm permits selective titiations, using appropriate p H conditions. A metal forming a very stable chelonate (such as thorium) can thus be titrated in the presence of one forming a much neaher chelonate (such as copper) by the use of acidic titration conditions. Conversrly, the precipitation of insoluble hydroxides a t high p H allows the selective determination of, for example, calcium in the presence of magnesium. Other complexing or precipitating agents may react with the metal, decreasing its concentration and the extent of chelonate formation. Such agents are often introduced intentionally to mask certain ions and improve selectivity. Complexing agents are also useful in alkaline solution to prevent precipitation of titrate ions as insoluble hydrolides, which, although theoretically titratable, react very sluggishly. The constituents of buffer systems are complesing agents and their masking ability can be used to advantage. In such cases the effective concentration of the complexing agent is p H dependent and must be considered in a fashion analogous to that employed for the chelon. Metal chelonate derivatives may be formed under certain conditions, increasing the extent of metal chelonate formation. I n acidic solution certain chelonates form derivatives of the type hIYH and in alkaline solution derivatives of the type MYOH may be formed. Other species such as ammonia may react with chelonates, yielding derivatives of the type MYNHa. Estimation of the extent of metal chelonate formation and hence of suitable titration conditions is possible if the pertinent equilibrium data are available. Two !imitations should not be overlooked, hoivever. First, the problem of slow attainment of equilibrium is often a serious one, especially in the titration of mixtures. One example, the reaction of a chelon with a precipitate, has already been mentioned. Less well documented are the relatively slow dissociation and displacement reactions of chelonates. Consider the chelonietric titration of a metal in the presence of a second which also forms a chelonate although of much lower stability. Both metals react initially with each drop of titrant and if the exchange reaction is so slow t h a t equilibrium is not attained, the end point will be late, indistinct, or both. The roughly inverse relationship between dissociation rate of a chelonate and its effective stability constant indicates the use of conditions giving the lowest chelonate stability consistent with adequate end point detection.
"' %
-
0.5
1.0
0.4 W
0 z 0.3a
0
cn 0.2 m
a
\
;-
m rr
0.5
1
,
\\
-1.01 -I . S I
I
I
Figure 2. Formation of ammonia derivative of CU-EDTA as a function of ammonia concentration
Second, the results will be no more accurate than the data from which they are calculated. I n some cases the data represent conditions which diverge widely from those under mhich the titration is to be performed. Often the data have appreciable uncertainty, and for some equilibria, notably those involving chelonate derivatives, few or no data are available. The lack of information has necessarily resulted in the omission of such reactions in the calculation of effective chelonate stability constants, and in turn in the tacit assumption that such equilibria are unimportant. However, the formation of chelonate derivatives can increase the effective stability constant of a metal chelonate by as much as 6 log K units and enhancement by 2 or more log K units often occurs under usual titration conditions (6). The sparseness of such data has prompted a brief investigation of ammonia derivatives of copper chelonates. As only a few systems have received attention, the study of chelonate derivative formation should be a very fruitful endeavor. AMMONIA DERIVATIVES OF COPPER CH EL ONATES
Ammonia often shows a marked tendency to participate in the reaction MY
+ NHI
4
MYNHs
(2)
where MY is a metal chelonate and MYSH3 is the ammonia derivative of the metal chelonate. The present study is confined to the ammonia derivatives of the EDTA, EGTA, and DTPA chelonates of copper. The extent of derivative formation was followed spectrophotometrically, as the absorption spectrum of the copper chelonate changes markedly upon formation of the ammonia derivative. If only the monoammine derivative is formed, the equilibrium constant
I
I
I
I
I
I
Figure 3. Graphical determination of stability constant of ammonia derivative of CU-EDTA
for the reaction may be evaluated readily from a plot of absorbance us. pNH3, where pNHa represents the negative logarithm of the ammonia concentration. These data for CuEDTA a t 720 mp are presented in Figure 2. The equilibrium constant for Reaction 2 is
or expressed logarithmically
The term (MY"3) ( M Y ) is given directly by A / B from Figure 3, so that log K = log i l / B pNHB. Hence, the value of log K is equal to the pNH3 a t which log d / B equals 0. The linearity of the plot of log A / B us. pNH3 in Figure 3 indicates the validity of the assumption that only the monoammine derivative is formed. The value of log K for the ammonia derivative of Cu-EDTA is 2.1. Similarly, values of log K were estimated to be about -0.1 for CuDTPA and Cu-EGTA. I n these cases, the instability of the ammonia derivative prevents ready achievement of complete reaction and hence the absorbance does not reach a steady value even a t an ammonia concentration of 1OAV. Because the limiting absorbance of the derivative must be estimated, some uncertainty in log K was introduced, For analytical purposes, the uncertainty produces no difficulty because, as the magnitude of the constant demonstrates, the extent of derivative formation is negligible under the usual analytical conditions.
+
ADJUSTMENT OF TITRATION CONDITIONS
Determination of Ca in Presence of Mg. The conventional chelometric determination of calcium in the pres-
ence of magnesium with E D T A involves titration of calcium a t p H 12.5, where magnesium is precipitated as the hydroxide and does not interfere. When the ratio of magnesium to calcium is greater than 1, calcium tends t o be coprecipitated, giving low results and an indistinct end point. The addition of citrate gives some improvement, but is not completely effective (10). Zinc in strongly alkaline medium has been employed as an amperometric indicator with a large reduction in coprecipitation because occluded calcium has an opportunity to equilibrate with the titrant during the deaeration period required before each measurement ( 2 ) . The selectivity of EGTA for calcium in the presence of magnesium has been reported by Sadek, Schmid, and Reilley (8). This titrant was chosen for the present study, since its selectivity for calcium obviates the precipitation of magnesium. The absolute stability constants of the calcium and magnesium chelonates are 11.O and 5.2 log units, respectively, while that of the copper chelonate is 17.8 log units ( 1 ) . The low stability of the magnesium chelonate precludes the determination of magnesium with this reagent, but this is a relatively minor disadvantage, as the sum of calcium and magnesium is determined readily by conventional methods. The relatively low stability of the calcium chelonate points to the use of alkaline titration conditions. A p H of 10 was chosen, as the chelonate has attained essentially maximum stability and this p H is nithin the effective buffer range of ammonia, which effectively complexes copper. I n order for copper to act as a photometric indicator for calcium, the effective stability of its chelonate must be reduced to a value less than that of the calcium chelonate. The formation of the copper chelonate will then occur
VOL 33, NO. 12, NOVEMBER 1961
1667
0.1 0,o
0
2
I
3
4
P "3
Figure 4. Dependence of effective stabilities of EGTA chelonates of copper, calcium, and magnesium on ammonia cdncentration
0.1
0.2
0.3
0.4
0.5
AMMONIA MOLARITY Figure 5. Extent of displacement of copper from EGTA by calcium as a function of ammonia concentration
Selective Determination of Cd and
Zn. Present chelometric procedures only after the calcium is completely chelated, the reaction sequence following path 1-2-4-5 in Figure 1. The appropriate ammonia concentration is readily determined from the plot of effective chelonate stabilities of CaEGTA, Mg-EGTA, and Cu-EGTA as a function of ammonia concentration given in Figure 4. An ammonia concentration of 0.3M (pSH3 0.5) reduces the effective stability of Cu-EGTA to approximately 3.6 log K units below that of Ca-EGTA. The effects of chelonate derivative formation and metal ion hydrolysis have been neglected. To check the calculated titration conditions, an empirical study was made of the effect of ammonia concentration on the relative stabilities of Ca-EGTA and Cu-EGTA. The competition of the two metals for one equivalent 6f chelon was measured by the absorbance a t 550 mp, where Cu-EGTA absorbs slightly and copper-ammonia species absorb strongly. Initially, in the absence of ammonia, the copper is completely chelated by the EGTA, leaving the calcium free. A titration of equimolar amounts of copper, calcium, and magnesium with EGTA under these conditions would follow path 1-3-5 in Figure 1, the preferential chelation of copper rendering it useless as an indicator for calcium. As ammonia is added, the copper is displaced from EGTA by calcium with the formation of copper-ammonia species
Ca+Z
+ CuY-2 + z NH,
+
-c
Cay-* Cu(NHJ.+*
(3)
as indicated by the increase in absorbance shown in Figure 5. Complete displacement of the copper, indicated by constant absorbance, occurs a t a n ammonia concentration of 0 . 5 M . Under these conditions the addition of EGTA 1668
ANALYTICAL CHEMISTRY
to a n equimolar mixture of calcium, magnesium, and copper would result in a titration corresponding to path 1-2-4-5 in Figure 1, the complete chelation of calcium corresponding to the first break in absorbance a t point 2. Ammonia concentrations much greater than 0.5M are undesirable. Continued increase in the ammonia concentration eventually reduces the stability of Cu-EGTA to such an extent that point 4 in the titration curve is depressed and the sharpness of the absorbance break a t point 2 is decreased. One point is worthy of comment in passing. The photometric nonlinearity of the displacement of copper from EGTA by calcium in the presence of ammonia in Figure 5 raises the question as t o whether similar results are to be expected in the titration of copperammonia species with EGTA, which must, of course, be linear for accurate end point location. As a general rule linearity in the chelometric titration can be expected whenever the titration is performed in the presence of an essentially constant excess of complexing agent. The nonlinearity of the displacement reaction is caused by the constantly changing distribution of the various copper-ammine species produced by the continuous stepwise addition of ammonia to copper combined with variations in the absorbancies of the individual copper-ammine species. In the chelometric titration, the reaction is a one-step addition of chelon to the copper. Further, if the complexing agent is present in essentially constant excess, the relative proportions and, thus, the average absorptivity of the copper reactant species will remain constant throughout the titration. Formally, the titration is equivalent to the conversion of a single reactant to a single product, with the result that the titration plot is linear.
for the selective determination of either cadmium or zinc in the presence of one another involve either a preliminary separation or the masking of one of the components followed by a second titration in which the sum is determined. Only the latter technique will be considered further. Sweetser and Bricker (10) determined cadmium in the presence of zinc by masking the latter with 0.7M sodium hydroxide and adding a small amount of cyanide to prevent precipitation of the cadmium. The end point was detected readily by the ultraviolet absorbance of free EDTA, but a considerable blank correction was reported. Pribil (4, 13) found that diethyldithiocarbaniate can be used to mask cadmium after the determination of the sum of cadmium and zinc. The liberated EDTA is titrated visually with a standard magnesium solution and Eriochrome Black T indicator; however, the bulky cadmium precipitate tends to obscure the end point color change. DTPA was selected as a titrant because of the unusually high stability of Cd-DTPA and Zn-DTPA. The absolute stability constants of CuDTPB, Cd-DTPA, and Zn-DTPA are 21.5, 19.0, and 18.8 log units, respectively (1). The use of copper as an indicator for the selective determination of cadmium or of zinc in the presence of the other is possible only if the effective stabilities of Zn-DTPA and Cd-DTPA differ by a t least 2 log K units, and that of Cu-DTPA is very nearly equal to that of the less stable. The appropriate order of chelonate stability is readily attained by the use of ammonia buffer a t p H 10. A plot of the effective stability constants of the three chelonates as a function of free ammonia concentration is presented in Figure 6. A free ammonia concentration of 0.1M gives the desired order
2
I
0,201
W 0 2
a
m
LT
a
m
a
0
2
I
3
4
TITRANT, ML.
,"P Figure 6. Dependence of effective stabilities o f DTPA chelonates o f copper, cadmium, and zinc on ammonia concentration
of chelonate stability, the effective stability of Zn-DTPX being within 0.4 log K unit of that of Cu-DTPA, while the effective stability of Cd-DTPA exceeds these by about 3 log K units. The importance of solution conditions in deteImining the shape of the titration curve is apparent from Figuie 1. I n the abscnce of ammonia, addition of D T P A to an equimolar solution of copper, cadmium, and zinc results in the preferelitid chelation of copper, folloi?ed by the simultaneous chelation of cadmium and zinc. The chelation sequence follons path 1-3-5 and copper is ineffective as an indicator. I n the presence of 0.1V ammonia, however, the sequence of chelation reactions follows path 1-2-5, segment 1-2 corresponding to the chelation of cadmium, and segment 2-5 corresponding to the siniultaneous chelation of copper and zinc. The titrations of equimolar amounts of copper and zinc or of copper and cadmium under these conditions are given by paths 1-4-5 and 1-2-4-5, respectively. Photometric study of the displacement of copper from DTPX by cadmium indicated that an ammonia concentration of about 0.3M was required to attain complete reaction. Under these conditions the copper chelonate required somewhat more than a minute to reach equilibrium after each titrant addition. This was true in chelometric titrations of binary mixtures of copper and zinc and of copper and cadmium. Increasing the ammonia concentration to 0.5-11 or decreasing the p H to 8 increased the rate of the exchange reaction, stable absorbance readings being obtained within 15 seconds of adding each increment of titrant. All titrations were performed in 0.5.44 ammonia at p H 10. The amount of copper indicator added
Figure 7. Deviation from ideal behavior in photometric titration of equimolar mixture of copper and zinc
was 0.1 mmole or an amount equal to the zinc, whichever was greater. The use of larger amounts of copper for samples containing zinc 1%as dictated by the cotitration of zinc and copper, a factor which decreases the absorbance change per unit yolume of titrant. The loner limit of 0.1 mmole m s set by the minimum absorbance change which could be detected n ith sufficient accuracy. The greater stability of the copper chelonate and the chelation of all the copper permit the use of a much smaller amount of copper than was employed in the calcium titrations. RESULTS AND DISCUSSION
Determination of Ca in Presence of Mg. Results for t h e determination of calcium, alone and in the presence of a tnentyfold molar excess of magnesium, are tabulated in Table I. A t t h e 8-mg. calcium level t h e accuracy is about +0.5yo and t h e uncertainty in results increases steadily as t h e calcium level is decreased, reaching +4.07, a t t h e l-mg. calcium level. S o significant difference in accuracy is apparent between titrations performed in the absence of magnesium and in the presence of a twentyfold excess of magnesium, nor was there a n y noticeable trend as the ratio of magnesium to calcium was varied from 0 to 20 a t the 4-mg. calciuin level. Selective Determinations of Cd and Zn. Results for t h e titrations of cadmium alone and in t h e presence of twentyfold molar excess zinc, and for zinc alone and in t h e presence of twentyfold molar excess cadmium, are presented in Table I1 and several titrations of equimolar mixtures of cadmium and zinc are given in Table 111. Considering first t h e titration of t h e individual metals in t h e absence
of the other, the results for cadmium are noticeably more reproducible and more accurate than those for zinc. The cadmium is determined directly,
Table 1. Determination of Calcium in the Presence and Absence of Magnesium
Mole
Ratio hl g ./Ca 0: 1
Calcium, M g . Found
Taken 8.52
10: 1 20: 1
0:1
4.26
5:1 10: 1
20: 1 0: 1
2.13
20: 1
0 :1 20: 1
1.065
8.57 8.55 8.52 8.48 8.49 8.52 8.52 8.49 8.52 4.26 4.22 4.24 4.23 4.26 4.22 4.23 4.28 4.27 4.26 4.28 4.27 2.14 2.13 2.14 2.09 1.99 2.09 2.15 2.07 2.11 1.015 1.044 1.092 1,025 1.112 1.043
VOL. 33, NO. 12, NOVEMBER 1961
70
Error $0.6 $0.4
10.0 -0.4 -0.4
10.0 10.0 -0.4
1 00 10.0 -1.0 -0.5 -0.7 10.0 -1.0 -0.7 $0.4 +0.2 1k0.0 $0.5 -0.2 $0.5 $0.5 $0.5 -2.0 -6.0 -2.0 +1.0 -3.0 -1.0 -5.0 -2 1 +2.7 -4.0 $4 7 -2.2
1669
whereas the ainc is determined by the difference between the total amount of chelon equivalent to the sum of copper and zinc and the independently determined amount of copper. Under the most favorable conditions, the zinc was equal to the copper, while in some cases the mole ratio of copper to zinc was as high as 10, and small errors in locating the end point mere unavoidably magnified when converted to relative error in the amount of zinc. The uncertainty of determination of cadmium or of zinc in the presence of a twentyfold excess of the other is mark-
Selective Titration of Cadmium-Zinc Mixtures
Table 11.
Mole Cd/Zn
Zn, .M g 9% Taken Found Error Determination of Zinc
0: 1
7.14
7.09 7.22 7.19 7.17 7.28 7.22 7.07 7.16 3.58 3.52 3.62 3.54 3.60 3.44 3.50 3.65 1.719 1,824 1,758 1.765 1.814 1.812 1.753 1.851 1.699 1.691 1,824 1.841 0.740 0.623 0.7!40 0 . 765 0.680 0.6G8
20: 1
0: 1
3.57
20: 1
0: 1
1.785
20: 1
0 :1
0.714
-0.8
+ I .3 +0.8 +0.4 $2.2 $1.3 -1.0 +0.3 +0.3 -1.4 $1.4 -0.9 $0.9 -1.3 -2.0 $2.2 -3.7 $2.2 -1.5 -1.1 +3.3 $1.5 -1.8 +3.7 -4.8 -5.1 +2.2 +3.3 +3.7 -12.8 $3 i Si.2 -3.7 -6.4 $3.7 -6.4 +4.6 +7.4 $4.6 -5.5
0,740
20: 1
edly greater than that of determination of the individual metals in the absence of the other. The average deviation in the determination of 7 mg. (0.1 mmole) of zinc increases from 0.8% in the absence of cadmium to 1.2% in the presence of a 20M excess of cadmium, while that for the determination of 11 mg. (0.1 mmole) of cadmium increases from 0.5 to 1.1% under parallel conditions. The increased uncertainty in the titrations of zinc is due to the uncertainty in the end point of the prior cadmium titration. There is some difficulty in locating the end point caused
0.668 0.747 0.767 0.747 0.675
Table 111. Cd, Mg.
Mole
Ratio Cd, Mg. % Zn/Cd Taken Found Error Determination of Cadmium 0: 1
11.51
20: 1
0: 1
11.57 11.59 11.48 11.45 11.30 11.43 11.46 11.74 11.62
5.75
20: 1
0 :1
2.88
20: 1
0: 1
1.144
20: 1
5.72 5.76 5.77 5.76 6.61 5.69 5.64 5.70 5.86 2.93 2.80 2.92 2.88 2.91 2.87 2.80 2.85 3.01 1.200 1.167 1,144 1,144 1,090 1.168 1.177 1.111 1,077 1.121 1.189 1.088
$0.5 +0.7 -0.3 -0.5 -1.8 -0.7 -0.4 $2.1 +1.0
-0.G $0.2 +0.4 f0.2 -2.5 -1.0 -1.8 -0.8 +2.1 $1.8 -2.8 +1.6
10.0 +1.1 -0.3 -2.7 -1.1 $4.5 $5.0 $2.0 2zo.o
10.0 -4.0 $2.0 +2.9 -2.9 -5.9 -2.0 $3.9 -4.9
Analysis of Equimolar Cadmium-Zinc Mixtures Zn, M g .
Taken
Found
70Error
Taken
Found
I.144
1,189 1.155 1,189 1 212
$3.9 +1 .o 13.9 +5.9
0.714
0.747 0.727 0.662 0.681
+4.6 $1.8 -7.3 -4.6
5.72
5.62 5.66 5.77 5.70
-1.7 -1.1 $0.9 -0.4
3.57
3.66 3.64 3.50 3.51
+2.5 +2 .o -2.0 -1.7
1670
ANALYTICAL CHEMISTRY
yo Error
by a barely perceptible continuous increase in absorbance during the titration of the large excess of cadmium, which n as carried nearly to completion by the addition of aliquots of concentrated titrant using a pipet. The greater uncertainty in the determination of cadmium is due to two factors. First is the smaller change in absorbance per volume of added titrant, 1%hich decreases the sharpness of the end point break. The second is somewhat less obvious and is caused by the fact that the titrations of zinc and copper are not exactly coincident. The slightly greater stability (0.4 log K ) of the zinc chelonate over that of copper evident in Figure 6 results in the preferential titration of zinc when equal concentrations of copper and zinc are present. This behavior is apparent in the titration of a binary mixture of equal amounts of copper and zinc (Figure 7 ) . The first additions of titrant produce smaller increases in absorbance than would be expected from linear interpolation between the initial and final absorbance readings. Initially, the major fraction of each titrant increment is converted into the nonabsorbing zinc chelonate. K i t h each additional increment a smaller fraction is converted to the zinc chelonate and a larger fraction to the copper chelonate. The resulting photometric titration is a continuous curve concave upward, which can be approximated M ithin the limits of experimental accuracy by two linear segments. This behavior is of no consequence in the titration of zinc alone, because the final break in absorbance can be readily determined and the initial point in the titration is fixed a t zero volume of titrant. Difficulty arises only in cases where the initial point in the zinc titrntion (the end point in the cadmium titration) must be located experimentally. Because of the large excess of zinc and copper relative to cadmium, only the first segment of the copperzinc titration appears and, further, i t is relatively linear. However, this small deviation from linearity is sufficient (when combined with the decreased slope of this segment) to cause a significant error in locating the cadmium end point. One “solution” to this problem is to add the correct amount of copper to cause cotitration of copper and zinc, but, of course, this would be of little use in the analysis of unknown samples An interesting result of the photometric behavior of copper in the presence of zinc is that the accuracy in the analysis of equimolar mixtures of cadmium and zinc is only slightly better than that for the determination of either metal in the presence of a twentyfold excess of the other. This is readily seen by comparison of the appropriate data in Tables I1 and 111. The initial
segment of the copper-zinc titration, which dcterniines the point of demarcation between the cadmium and zinc titrations, is rclatively short in the former cnse. The difficulty in locating the cadmium end point prevents the attainment of the desired accuracy, although the titration conditions appear f n i oi xble a t first glance. The prediction from stability constant data of conditions for strict cotitration of two metals is a t present impossible. The maximum difference in effective stability constants which permits a linear photometric titration is about 0.1 log K unit, while the various equilibria which must be considered in establishing the effective stability constant of a metal chelonate are often uncertain by 0.5 log K unit. Although in some cases it may he possible to adjust conditions empirically to obtain linear titration plots, the limited advantage to be gained over approximate
attainment of cotitratioii conditions docs not appear to warrant the effort involved. The most promising application to cadmium-zinc determinations seems to be the sclecti\-e determinxtion of one or the other when present as a minor constituent of a binary mixture in conjunction with a separate detcrmination of the sum. Under such conditions the decreased accuracy would not be a serious problem. As separation or masking techniques are least efficient under such conditions, this technique should effectively complement conventional chelometric methods. LITERATURE CITED
(1) Hollonay, J. H., Reilley, C. N., ANAL. CHEM.32.249 11960). (2) Laitinei, H . ~.4., Sympson, R. F., Ibid., 26, 556 (1954).
(3) Mdmstadt, A. V., Gohrbandt, E. C., Ibid., 26, 442 (1954).
(4) Pribil,
K.,“Komplexometrie,” Chemapol, Prague, 1954. 1. 5,) Ramaiah, N. A. Vishnu, 9nal. Chim. Acta 16, 569 (1957). (6) Heilley, C. N., Porterfield, W. LV., ANAL.CHEM,28,443 (19561, ( 7 ) Reilley, C. S . , Schmid, R. W., Sadek, F. S.,J . Chem. Educ. 36, 555 (1959). (8) Sadek, F. S., Schmid, R. W., Reillcy, C. PIT., Talanta 2, 38 (1959). (9) Schwarzenbach, G., Gut, R., Anderegg, G., Helv. Chim. Acta 37,937 (1954). (10) Sweetser, P. B., Bricker, C. E., A i i . 4 ~ . CHEY.26,195 (1954). (11) Underwood, A. L., I b i d . , 25, 1910 (1953). (12) Ibid , 26, 1322 (1954). (13) Welcher, F. T., “Analytical Uses of Ethglenediaminetetraacetic Acid,” p. 163, Van Nostrand, Princeton, N. J., 1958.
RECEIVED for review February 13, 1061. Accepted August 25, 1961. Research supported by the United States Air Force through the .4ir Force Office of Scientific Research, Air Research and Development Command, Contract No. AF 49(638)-333.
Improvements in the Fluorometric Determination of Submicrogram Quantities of Beryllium CLAUDE W. SILL, CONRAD P. WILLIS, and J. KENNETH FLYGARE, Jr. Health and Safefy Division, U. S. Atomic Energy Commission, ldaho Falls, ldaho
b A recently published procedure for the fluorometric determination o f beryllium using morin has been improved significantly, and its application has been extended. Stabilization o f alkaline solutions of morin toward air has been accomplished without use of stannite or other reducing agents. Use of diethylenetriaminepentaacetic acid in place o f (ethylenedinitril0)tetraacetic acid prevents formation o f fluorescent complexes of morin with scandium, yttrium, and lanthanum, and increases the selectivity greatly. A new combination of primary and secondary filters produces a threefold increase in the ratio of net beryllium fluorescence to blank fluorescence while requiring an instrumental sensitivity only one fourth that obtained with the previous combination. Since the exciting wave lengths are entirely in the visible region of the spectrum, errors produced b y colorless ions that absorb in the ultraviolet are eliminated. One o f the most important discoveries was the extensive adsorption o f beryllium from alkaline solution on the glass walls of the container. The fluorescent species contains beryllium and morin in a mole ratio of 1 to 1. Detailed procedures are presented for the determination o f beryllium in metallic thorium, zirconium, uranium, copper alloys, and
aluminum, in rare earth oxides and phosphates, and in silicates such as beryl and clay that are not decomposed completely b y either pyrosulfate fusion or hydrofluoric acid. Beryllium can b e determined in air dusts a t concentrations well below the maximum permissible levels without separations o f any kind in approximately 30 minutes.
I
P; a
previous publication, a procedure for the fluorometric determination of heryllium using niorin (2’,4’,3,5,7pentahydrosyflavone) was described (‘7). The procedure is very sensitive and precise, but suffers from three important disadvantages. Preparation of the solution for fluorometric measurement is rather tedious when many samples are being analyzed. The procedure is less selective than is desirable as thorium, yttrium, zirconium, scandium, lanthanum, and lithium each produce bright yellon-ish green fluorescence in descending order of sensitivity under the conditions used. Also, anomalous results are obtained occasionally that, for some unknon-n rea?on, are clearly inconsistent with the usual high precision. The present procedure is more rapid and sensitive than before and all sources of anomalous results have been identified and corrected.
~ i ~ t l i y l ( ~ n c t r i u i n i n e ~ e ~ i t a a cwid etic (DTPA) is very similar to (Pthylenediniti~i1o)tetraacetic acid (EDTA) both structurally and in its reactions with metals to form strong water-soluble complexes of the same type and wit’hthe s a m ~met,als. I n general, the DTPA romplexes have formation constmts that are 10 to 1000 times larger than the corresponding complexes with EDTA\(1-3). As shown by the data of Table I, DTI’A increases the selectivity of the fluorescence procedure by forming complexes with yttrium, scandium, lanthanum, and lithium of such increased stability that much laiger quantities of the metal are required t o produce significant fluorescence in comparison to the previous method using EDTA. I n contrast, the beryllium c*omplex with niorin is completely unaffected. Unfortunately, the niorin complexps of thorium or zirconiuni are d s u affected very little by DTPX. -4 iia1ytic:il 1iim:edures for these elements ai’(’ currcntlj. bcing developed to take iitlwntnge of the exceptional sensitivity :ind stability of the morin complexes. Most procedures for the fiuorometric clctcrinination of heryllium, using niorin, einploy a solution of sodium stannite to stabilize the fluorescence by preventing the xir osidation of morin. If EDTA is added, the fluorescence is stable for over VOL. 33,
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