Potentiometric Characterization of Aluminum Aminopolycarboxylate Chelonates David A. Aikens and Fred J. Bahbah’ Department of Chemistry, Rensselaer Polytechnic Institute, Troy, N . Y. 12181 Potentiometric characterization of effective chelonate stability constants is extended to strongly hydrolyzed metal ions by combining measurement of chelonate formation with the mercury pM electrode in acid solution and pH titration of the metal chelonate. Absolute chelonate stability constants, formation constants of proton and hydroxyl chelonate derivatives, and effective chelonate stability constants are reported for AICyDTA, AIDTPA, AIEDTA and AIHEDTA. The application of these data to development of analytical methods is illustrated by development of a simple, rapid method for selective analysis of iron-aluminum mixtures.
EFFECTIVE chelonate stability constants, an important aid in selection of optimum conditions for chelometric titrations, can be calculated accurately for most common metal ions. An important exception, however, is aluminum. Although absolute stability constants have been reported recently for the common aluminum chelonates ( I ) , the unusually large influence of chelonate derivative formation on effective chelonate stability and the limited information concerning the pertinent equilibria prevent comparison of the relative merits of chelons for analysis of aluminum-containing mixtures. Formation of proton chelonate derivatives makes a substantial contribution to the effective chelonate stability in acid solution and formation of hydroxyl chelonate derivatives makes a very large contribution t o the effective chelonate stability in alkaline solution. For the protonated derivatives, only approximate values of the formation constants ( I ) are available, however, and for the hydroxyl derivatives, formation constants are available only for the cyclohexanediaminetetraacetic acid (CyDTA) and ethylenediaminetetraacetic acid (EDTA) chelonates ( 2 ) . The present paper reports absolute and effective stability constants for the aluminum chelonates of CyDTA: EDTA, hydroxyethylethylenediaminetriacetic acid (HEDTA), and diethylenetriaminepentaacetic acid (DTPA) and the estimated maximum value of the absolute stability constant of the ethylene glycol bis-(P-aminoethyl ether) tetraacetic acid (EGTA) chelonate. Application of these data to determination of aluminum is illustrated by development of a simple, selective method for analysis of iron-aluminum mixtures that requires only a single sample and avoids the need of chemicai masking. The principal reason for the lack of knowledge of stability constants of aluminum chelonates is the fact that chelonate formation can be studied only between the approximate pM limits of pH 1.5 and p H 3.5. Below p H 1.5 limited soicbility of free chelons and extensive dissociation of the ai,minum chelonate prevent study o f chelonate formation. Above
pH 3.5 aluminum hydrolyzes extensively but neither the ey.act nature nor the degree of hydrolysis is clearly established (3). Hence the concentration of aquo aluminum ion becomes indeterminate above pH 3.5 and measurement of aluminum chelonate stability constants is impossible. The narrow pH range available for study of aluminum chelonare formation led Schmid and Reilley ( 4 ) , who introduceci poteritiometric measurement of metal chelonate stability constants with the mercury electrode, t o conclude that this powerfLt technique could not be used to measure aluminum cheionate stability constants. The mercury electrode cap. in fact be used for measurement of aiuminum chelonate stability constants but the potential-pH diagram differs significantly from that usually obtained. During, the present study Moel!er and Chu ( I ) demonstrated the use of the mercury electrode for measurement of aluminum chelonate stability constanx, but two important questions must be answered to permit characterization of the aluminum chelonates for analytica: purposes. First, Moeller and Chu ( I ) did not comment on the unusual nature of the potential-pH diagrams of the aluminum chelonates, a feature which is of general interest in study c f chelonate formation of strongly hydrolyzed metai ion-. Second, their primary objective was study of the thermodynamics of the simple cheionate formation reaction rathir than characterization of the analytical potential of the cheIons and they therefore placed little emphasis on the analyrically important aspect of cheioriate derivative formation. The present paper complements that of Moeller and Chu and emphasizes the analytical aspects of the aluminum cheionatr; Chelonate Equilibria. Potentiometric measurement or chelonate stability constants depends on the same basic principle for aluminum as for other metal ions-i.e., L;e., termination of the extent of exchange between the test metai ion and the mercury chelonate. Interpretation of the result:; is more complex for aluminum than for metal ions that d : ~ not hydrolyze, however, because chelonate formation reactions of aluminum differ in two important respects frow those of metal ions that d o not hydrolyze. First, over the accessible pH range the chelonates of bot!: aluminum and mercury form protonated derivatives. Hence the simple exchange reaction usually observed between the test metal, M, the test metal chelonate. MY, and the mercur!’ cheionate, HgY, generalized in Equation 1 M
+ HgY
MY
+ Hg
fli
does not accurately describe the observable exchange reactions of aluminum chelonates. The aluminum chelonates oc CyDTA, EDTA, and HEDTA obey instead the compie;: exchange equilibrium in Eayation 2
Present address, Department of Chemistry, Geneva College, Beaver Falls, Pa.
(1) T. Moeller and S. Chu, J . Inorg. Nucl. Chem., 28, 153 (1966). (2) G. Schwarzenbach and G. Biedermann. Helo. Chim. Acia, 31,
459 (1948). 646
ANALYTICAL CHEMISTRY
(3) L. G. Sillen and A. E. Martell, “Stability Constants of Metal. Ion Complexes-Special Pubiication No. 17,” The Chemicai Society, London, 1964, pp. 65-6. (4) R. W. Schmid and C . N. Reilley, 1.Am. Chem. Soc., 78, 5513 ( 1‘1156).
3
ng H T - I
'I-
200
1.5
2.5
3
PH Figure 1. Potential-pH plots for aluminum chelonates Conditions 0.1M KNOI, 25.0" C Dashed lines indicate equal concentration of chelonate and protonated derivative
'1
HgY +H+
A1
I
-H+
AIY +H+
+ .YgYH e
1I
-H+
AlYH
(2)
-t Hg
where HgYH and AlYH represent monoprotonated derivatives of the mercurj' and aluminum chelonates. At the lower pH limit of observtition, approximately p H 1.5, protonation ol both the aluminum and mercury chelonates is essentially complete and the tixtent of exchange is determined principally by the effective stability constants of the protonated :iluminum and mercury chelonates. As the p H is raised the -eiative proportion of the unprotonated metal chelonate ini:reases and a t pH :#.5 both the protonated and unprotonated :helonate are important in the exchange reaction. The exchange reaction for DTPA is similar t o Equation 2 except :hat the predominant forms of the mercury chelonate are the Jiprotonated derivative a t u t i 1.5 and the monoprotonated fierivative at pH 3.5. T'he second unus Jal feature of aluminum chelonates is the ' k t that formation of hydroxide cheionate derivatives makes ,'n important contr bution to the effective stability constants at .?igh pH. Formiation of chelonate hydroxide derivatives Liinnot he detected ,n study of chelonate formation a t p H 1.5 to pH 3.5, but the formation constants of hydroxide chelonate >:ierivatives arc readily evaluated from p H titration of the !.:ielonates with sollium hydroxide. Combination of these -esults with the results of the potentiometric studies at low p H ?her. permits prediction of sikctive stability constants from -ri i to pH 13. ZXPERIMENTAL
Reagents. Deminera!i;ea water and analytical grade - a g e n t s were dsed throughout. Stock solutions of Al+3 ..nd of Hg+2, 0 0100M in H N 0 3 and approximately 0.1M In metal ion, were preparea from the nitrates and standardized .gainst EDTA Aluminum was standardized by backtrdtion dt pti 5 with zinc and xylenol orange indicator and iercury was ;imdardized by direct potentiometric titration .t DH 2 . Ferric Titrate stock solution for titrations was htanaardized by ba:k titration with bismuth at pH 2 using 'iJaenol orange. Chelons in free acid form (Lamont Laboradries, Dallas, Tex.) were m e d a t 100" C. Standard chelon ~iiifionsof known degree of protonation were prepared .-y dissolving the appropriate quantity of chelon in sufficient rdndard NdOH to half-neutralize the chelon followed by .iilut:on to volume.
Procedure. All measurements were made in 0.1M KNOs solution at 25.0 =t0.2" C. Potential-pH data were collected as described by Holloway and Reilley (5). A calibrated Corning Model 10 pH meter was used for p€i and potential measurements. The mercury electrode potential was measured with the 300-mV scale using a 100-rnV bucking potential when necessary. The pH titration of a 0.01M solution of each chelonate was performed from p H 1.5 t o pH 13 using standard 0.5M NaOH. The titration curve was corrected for the amount of H N 0 3 in the aluminum stock solution and for the amount of NaOH in the chelon stock solution and the formation constants of the hydroxide chelonate derivatives were calculated as described by Schwarzenbach and Heller (6). RESULTS AND DISCUSSION
Potential-pH Diagrams. Potential-pH diagrams are given in Figure 1 for CyDTA, DTPA, EDTA, and HEDTA with circles representing experimental points and soIid lines calculated from the proposed chelonate equilibrium constants. Line I in each diagram represents the Hg(l1) activity in a solution lO-3M in both mercury chelonate and free chelon and line I1 represents the Hg(I1) activity generated by the exchange reaction given in Equation 2. For study of the exchange reactions, the mercury a n a aluminum chelonates and Al(II1) were present a t ]10-3M except for the exchange reaction of CyDTA in which case it was necessary to raise the AlCyDTA concentration from 10-3M to 10--2M. At the lower concentration of AlCyDTA, the potential of the mercury electrode was repressed and a white precipitate a p peared. Schmid and Reilley ( 4 ) observed a similar effect with certain EDTA chelonates and identified the precipitate as the mercury(1) salt of mercury(l1) EDTA and it seems likely that a similar explanation fits the present results. Each dashed vertical line represents the p H a t which the metal chelonate and the protonated derivative are present a t equal concentration. The potential-pH diagrams have the same general form as those reported by Holloway and Reilley ( 5 ) for unhydrolyzed metal ions, with one important difference. Instead of the p H independent potential that is characteristic of the exchange reactions of unhydrolyzed metal ions, line I1 in each of the aluminum potential-pH diagrams shows a strong pH depend( 5 ) J. H. Hoiloway and C . W. Reilley, ANAL.CUEM., 32, 240 (1960). ( 6 ) G. Schwarzenbach and G. Heller, Helc. Chim. Aero. 34, 576 (195 I). VOL 39, NO. 6, MAY i967
647
Table I. Absolute Stability Constants and Proton and Hydroxide Derivative Formation Constants of Aluminum Chelonates Chelon EDTA HEDTA DTPA CyDTA log KabpA!\' iog K: AlYH log K , AlYOH
14.4 18.7 18.9 2.4 4.3 3.4 9.3 6.6 6.3 Conditions: O.IM KNO,, 25" C. Formation constants of proton and hydroxyl chelonate derivatives refer to reaction of the metal chelonate with either H+ or OH-, respectively. 16.5 3.4 8.0
24
20 m
.-
t
16
. ) c
D
Y
c3 0
12
-J
ence. This pH dependence, caused by the complexity of the exchange reactions for aluminum indicated in Equation 2, necessitated evaluation of the pertinent equilibria in the following manner. The absolute stability constant of the mercury chelonate was established from the potential of line I above pH 5. These values agreed with those reported by Holloway and Reilley ( 5 ) and Schmid and Reilley ( 4 ) within 0.1 log K unit. Above pH 5 the slope of line I agreed veri' well with the value predicted from the chelon acidity constants. but below pH 5 the experimental values of potential showed a n increasing systematic negative deviation from the predicted values. This deviation was attributed to protonation of the mercury chelonate and the formation constant of the protonated derivative was determined by appropriate modification of the method described by Holloway and Reilley ( 5 ) for calculation of formation constants of hydroxide chelonate derivatives. Values of the absolute stability constant of the aluminum chelonate and of the formation constant ot the protonated derivative of the aluminum chelonate were then selected to best fit the observed potentials for the exchange reaction. Calculated and observed values of potential for the exchange reaction agreed within 3 mV over the observable pH range, corresponding t o an uncertainty of 0.1 k:g K unit in the effective stability constants of the aluminum c.'"': i onii tes . ; i n important difference between aluminum chelonates and :.h -;.~ior,;ttesof most metal ions which is not apparent in the po:ential-pH diagrams is the fact that the chelonate equii!bri;m is established much more slowly for aluminum than :a:. most metal ions. Holloway and Reilley ( 5 ) report that for tile majority of metal chelonate formation reactions, stable porcntials were established at the mercury electrode within 1 n!rfiute. Much longer times were required to establish equiiibriuni in formation of aluminum chelonates, however, ranging from approximately 2-4 minutes with HEDTA to 1C-15 minutes with EDTA. DTPA, and CyDTA. These observations give a valuable qualitative indication of the relative rates at which the chelons should reach equilibrium in titrations and suggest that HEDTA with its rapid equilibration should be useful in selective titrations involving aluminurr.. Chelonate Stability Constants. Values of aluminum c!xlonn!e stability constants and formation constants of protor-, and hydroxide chelonate derivatives are summarized in Table I. Formation constants of proton and hydroxide chelonate derivatives refer to the reaction of the metal chelonate with the proton or hydroxide ion. Formation constants of proton derivatives of mercury chelonates are also included iri. Table I, but values of mercury chelonate stability constants aie omitted because they agree within 0.1 log K unit with values reported by Schmid and Reilley ( 4 j and Holloway and Reilley (Sj. These values are higher than the corresponding values reported by Moeller and Chu (I) by approximately I
~
>
648 *
ANALYTICAL CHEMISTRY
e
4
0
Figure 2. Effective stability constants of aluminum chelonates as a function of pH Conditions 25" C, 0.1M KNOI
0.5 log unit. The values obtained for the stability constar?:, of the aluminum chelonates of EDTA, CyDTA, and DTPA agree well with the values obtained potentiometrically by Moeller and Chu. The value obtained for the stability cojjstant of AIHEDTA exceeds the value obtained by .Moeiier and Chu by approximately 2 log units but the reason fo: trcdiscrepancy is not clear. Both potentiometric estimates CY the stability constant of AIEDTA agree well with the valu-. of 16.1 log units obtained polarographically t>y Schwarzeribach, Gut, and Anderegg (7) and confirmed in a pH study b? Saito and Terry (8). The vaiue of the stability constant of AlCyDTA obtained polarographically by Schwarzenbaci. Gut, and Anderegg IS approximately 1 log unit less than in#: potentiometric vaiues, however. The pniarographic measurements were made under conditions similar to those in t i t potentiometric studies and the discrepancy may reflect some unsuspected kinetic complexity in the polarographic nieasurements. Values of ef'fective stability constants for the aluminum chelonates calculated from the equilibrium constants in Tablr I are plotted from pH 1 to pH 1 3 in Figure 2. These a r c based on total concentrations of chelon and chelonate but neglect hydrolysis of aluminum. The principal feature a' these plots is the fact that OVei the entire pH iange differences in effective stability constants among the aluminum chelonates anticipated on the basis of the absolute stability constants are not observed. The absolute stability constants of the aluminum chelonates span a range of 3.5 log units DU: the maximum difference in effective stability constants at ani' given pH is only 2.0 log units. Leveling of effective chelonate stabilities for a series of cheions IS expected in acid soiutio::
(7j G . Schwarzenbach, R. G ~ and I G . Anderegg, Heiti. Chim Acra, 37, 937 (1954;. (8) K. Saito and H. Terry, i.C/7tw. Soc., 1956. 4701.
because chelons are extensively protonated and the proton affinities of a series of chelons fail in essentially the same order a s the affinities o f t he chelons for a given metal ion. At high pH protonation of the cheion is unimportant and for most metal chelonates the value of the effective stability constant approaches the value of the absolute stability constant. For aluminum chelonates in alkaline solution, however, the effective stability constants d o not accurately r e k t the values of the absolute stability constants because formation of hydroxide chelonate derivatives is extensive. The extent of hydroxide chelonate derivative formation varies approximately inversely with the absolute stability constant of the chelonate. Hence the effective stability constants of aluminum chelonates tend to be levelled at high pH. As an example the absolute stability constants of the aluminum c h e b nates of HEDTA and CyDTA are, respectively, 14.4 log units and 18.9 log units, a difference of 4.5 log units. At pH 13, however, the effective stability constants of the two chelonates are, respectively, 22.8 log units and 24.2 log units, a difference of only I .4 log units. The chelonate stability constant of AIEGTA is insufficient t o permit its quaniitative determination from the potentialpH diagram, but the maximum value can be estimated easily. The potential for the AIEGTA exchange reaction did not exhibit the approximate plateau observed with the other chelons below pH 3.5, but followed within 10 mV the potential for the HgEGT.A-EGTA exchange reaction. Apparently AIEGTA is a relatively weak chelonate and Al(III) does not displace Hg(I1) significantly from EGTA below p H 3.5. The mercury electrode potential corresponding to pH 3.5 is 230 mV GS. SCE for EGTA and the potential plateau of the exchange reaction must fall at least 30 m V more negative. On this basis the atsolute stability constant of AlEGTA must be less than 13 log units and EGTA may be a useful titrant for selective determination of other metals in the presence of aluminum. Because of the inability to obtain a quantitative estimate of the chelonate stability constant, however, this aspect was not pursued. Iron-Aluminum Analysis. The problem of selective analysis of iron-aluminum mixtures provides a good example of the potential application of aluminum chelonate stability data to developmeint of analytical methods. This problem was reviewed recen1.ly by Pribil(9) who pointed out that only EDTA and CyDTA. have been applied. Because of a lack of suitable indicators .Tor aluminum only indirect methods have been successful enough to gain general acceptance. A number of such in'direct methods have k e n developed but almost all of these are relatively minor variations of the method originated by Houcla and coworkers (10) based on masking of aluminum with fluoride and back titration of excess EDTA with zinc or lead. Two samples are required, however, and aluminum must be determined by difference. The need for two samples is eliminated in a modification proposed by Pribil and Vesely (11, 12) that uses CyDTA, but sample manipulation is increaiied substantially and iron must be determined by difference. (9) R. Pribil, Talanta. 12,925 (1%5). (IO) M. Houda, J. Korbl, V. Bazant, and R. Pribil, Colfecrim Czechosloo. Chem. Communs.,24, 700 (1959). (I 1) R. Pribil and V. Vesely, Tufanta,9 , 2 3 (1962). (12) R. Pribil and V. Vesely, Chemist-Analyst, 55,68 (1966).
Table II. Analysis of Fe-AI Mixture with HEDTA Fe, mg - AI, mg Taken Found Error, Taken Found Error, 5.25 5.27 +0.4 2.70 2.68 -0.8
__
5.25 5.25 5.25 5.25
5.23 5.20 5.21 5.26
-0.2 -1.0 -0.8 +o. 2
2.70 2.70 2.70 2.70
2.71 2.71 2.74 2.71
+0.4 f0.4 +1.6 +0.4
A more direct approach suggested by Guerrin, Sheldon, and Reilley (13)is based on direct titration of iron with EDTA in the presence of aluminurn and utilizes the low stability and slow formation rate of AlEDTA to prevent interference by aluminum. The iron titration must be performed very slowly near the end point, however, to avoid a false end point. Reexamination of this problem in light of the effective stability constants of aluminum chelonates shown in Figure 1 and the relative reaction rates of the respective chelons suggest that HEDTA would be optimal in this application. Figure i shows that at pH 2 EDTA forms one of the most stabie aluminum chelonates, and HEDTA forms the least stable aluminum chelonate. On an equilibrium basis all the chelons in Figure 1 form aluminum chelonates of sufficiently low stability a t pH 2 to prevent interference by aluminum in the determination of iron. Once formed, however, aluminum chelonates dissociate slowly as indicated by the sluggish establishment of equilibrium in the measurement of aluminum chelonate stability constants. Hence aluminum can interfere in a kinetic manner near the end point in the iron titration through reaction with a local excess of chelon followed by slow dissociation of the aluminum chelonate. To avoid this type of interference the effective stability constant of the aluminum chelonate should be as low as possibie to minimize local formation of aluminum chelonate and the dissociation rate of the aluminum chelonate should be as high as possible. Of the cheions studied HEDTA best meets both requirements and was evaluated experimentally. Iron is titrated directly in nitric acid at pH 2 using SNAZOXS indicator to a maximum deep red end point. Excess HEDTA is added, the pH adjusted to 5 with acetate buffer, and excess HEDTA is back-titrated with copper. At room temperature the iron end point is sluggish but at 50" C the color change occurs rapidiy and the titration can be performed at a normal rate. Replicate results for samples containing approximately 0. mmole each of iron and aluminum are given in Table II. For five determinations the relative standard deviations for iron and aluminum are 0.6% and 0.8 respectively, and the relative errors for iron and aluminum are -0.37;, and +0.4 %, respectively.
z,
RECEIVED for review June 30, 1966. Accepted March 13, 1967. Supported in part by a research grant from Nationzl Science Foundation. F. J. B. thanks the National Science Foundation for participation in a Summer Institute in Instrumental analysis at Rensselaer Polytechnic Institute. (1 3) G. Guemn, M. Sheldon, and C. N. Reilley, Ibid.,49, 36 (1960).
VOL 39, NO. 6, M A Y 1967
649
