- Email: [email protected]
A study of copper and cadmium iminodiacetate complexes by ion-selective electrodes and application to cadmium monitoring
Analytica Chimica Acta, 152 (1983) 191-202 Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands
A STUDY OF COPPER AND CADMIUM IMINODIACETATE COMPLEXES BY ION-SELECTIVE ELECTRODES AND APPLICATION TO CADMIUM MONITORING
R. STELLA* and M. T. GANZERLI
VALENTINI
Dipartimento di Chimica Generale and Centro di Radiochimica ed Analisl per Attivazione de1 C.N.R., Universitci di Pavia, Viale Taramelll 12, 27100 Pavia (Italy) (Received 6th March 1983)
SUMMARY The chelating properties of the iminodiacetate ion (IDAl-) towards copper(I1) and cadmium(I1) were investigated through pM and pH measurements at constant ionic strength of 0.1 M and at 25°C in the pH range 4-9. The acidity constants found were in good agreement with those already reported but significant differences from reported values were obtained for the stepwise copper complex formation constants, though the overall constant was in good agreement. Evidence is given for the existence of the protonated species CuH (IDA);, which proved helpful in explaining the anomalous complexing capacity of IDA in acidic medium. Practical applications include the use of IDA for lowering the level of copper interference in monitoring cadmium ion with a cadmiumselective solid-state electrode; this provides reliable measurements which are otherwise impossible.
The reported stability constants of complexes of the iminodiacetate ion (IDA’-) with copper(I1) and cadmium(I1) ions [l-5] are sufficiently different to suggest the possibility that this ligand may be employed as a selective masking agent for copper in the presence of cadmium. Yet measurements conducted on both the Cu-IDA and Cd-IDA systems with ion-selective electrodes (i.s.e.) showed discrepancies from the values calculated from the available constants. For instance, in the acidic region the extent of copper complexation was greater than the calculated value, thus suggesting the presence of a protonated complex. Therefore it seemed worthwhile to investigate these systems by coupling i.s.e.‘s to pH measurements; this technique allows a more detailed insight into metal complex equilibria than is possible by previous methods. A preliminary investigation was also made on iminodiacetic acid dissociation constants and the values obtained were compared with those already reported [ 1,2,6,7].
0003-2670/83/$03.00
0 1983 Elsevier Science Publishers B.V.
192 EXPERIMENTAL
Reagents Sodium iminodiacetate (Fluka) was purified by double recrystallization from methanol and dried at 115°C. The corresponding acid was prepared by passing a 0.1 M disodium salt solution through two columns in series (2 cm i.d., 20 cm high) filled with Dowex 5OW-X4 resin converted to the acid form with 0.1 M perchloric acid, and repeatedly washing with deionized and twicedistilled water. Only the central fractions of the eluate were collected and the absence of sodium was checked by atomic emission spectrometry; it is very important to discard all head and tail fractions, as they may contain undesirable ligands. The acid was then diluted and potentiometrically titrated with standard sodium hydroxide or standard copper solution at pH 7 using the corresponding i.s.e. The carbonate-free 0.1 M sodium hydroxide was prepared from a 50% (w/w) solution and standardized against potassium hydrogenphthalate. A 0.1 M perchloric acid solution was prepared from the 70% acid (Merck, Suprapur) and standardized against sodium carbonate. A 10m2 M copper stock solution was prepared by dissolving a suitable amount of copper(I1) perchlorate hexahydrate (ultrapure, Alfa Products) and standardized by titration with EDTA; end-points were detected visually and potentiometrically, with murexide as indicator for the former. A 10m2M cadmium stock solution was prepared by dissolving a weighed amount of cadmium wire (ultrapure, Alfa Products) in diluted nitric acid (Suprapur, Merck) in a teflon beaker. The excess of nitric acid was evaporated on a hot plate and the residue was taken up with water. The resulting solution was standardized by titration with EDTA; both potentiometric and visual (Calcon indicator) end-points were obtained in ammoniacal medium. The water used in all experiments was deionized and then twice-distilled in a quartz apparatus. Instrumentation A digital millivoltmeter (Orion model 701) was used in all measurements. A combined glass electrode (model 91) was used for pH measurements. The glass electrode was calibrated as suggested by Rajan and Martell [2] by direct titration of acetic acid; the observed readings were compared with the actual hydrogen ion concentrations tabulated by Harned and Owen [ 81. In the pH region below 3.5, the meter was calibrated by measurements of solutions with known concentrations of perchloric acid. The concentration of the hydroxide ion was obtained by using a p(K,), value of 13.79 at 25°C in 0.1 M potassium nitrate. The copper ion-selective electrode (Orion model 94-29A) was used with a single-junction Orion 90-01 reference electrode. The electrode was carefully calibrated at ionic strength 0.1 M with Cu-EDTA buffered solutions as described in a previous paper [9] ; a tenfold change in the copper concentration yielded an electrode response change of 29.0 + 0.5 mV, as expected.
193
The cadmium ion-selective electrode (Orion 94-48A) was used similarly and calibrated following the analogous procedure for copper: the mV response plotted in a semilog scale versus the free ion concentration was found to be linear down to lo-” M in metal buffer solutions as in the procedure suggested by R8iiEka and Hansen [lo]. Procedure All experiments were done at 25°C in a constant temperature bath, the ionic strength (I) being adjusted to 0.1 M with potassium nitrate (Suprapur, Merck). The iminodiacetic acid and its disodium salt were titrated under a nitrogen atmosphere, by first bubbling a nitrogen stream for 30 min and then maintaining the stream in the cell over the solution. Solutions having the following concentrations were used: CA = 4.50 X lo-* M for the acid and CB = 8.44 X lo-* M or 9.20 X 10m3M for the disodium salt. The copper-IDA and the cadmium-IDA systems were separately investigated by simultaneously measuring the free metal ion concentration and the pH in two different sets of experiments. First, titrations of unbuffered IDA (disodium salt) solutions were done with standard copper or cadmium solutions, with Co, = 5.82 X lo-’ M, Ccd = 4.38 X lo-* M and CB = 4.22 X 10m3M or 9.20 X lo-’ M. Next, solutions of ligand and metal ion in 1:l and 2:l molar ratios were titrated with standard 0.1 M sodium hydroxide; here CA = 2.11 X lo4 M, 4.22 X 10m3M and 1.00 X 10m3M for the copper system, and CA = 3.88 X lo-’ M, 7.39 X 10W3M and 3.68 X lo-* M for the cadmium system. The results in both cases are presented in graphic form with pH and pCu or pCd plotted as a function of the titrant added. RESULTS
AND DISCUSSION
Evaluation of acidity constants Considering the iminodiacetic acid as a diprotic acid, Chabereck and Martell [l] found p& = 2.54 and pK,* = 9.12 (I = 0.1, T = 30°C); Rajan and Martell [2] reported p& = 2.50 and pK,, = 9.40 (I = 0.1, T = 25°C) and Thompson [7] found pK,, = 2.58 and pK,, = 9.33 (I = 0.1, T = 25°C). All these values agree well, even with those that may be derived from the results of Liberti and Napoli [6] : log p1,2 = 9.17 and log p2,* = 11.73 (I= 0.5, T = 25°C). Liberti and Napoli also gave a third formation constant related to a H,IDA+ form, which begins to play a significant r61e at pH < 2.5. In this paper, the evaluation of the first two constants was considered sufficient to allow an accurate description of the investigated system. The acidity constants were calculated by using the well known Bjerrum method which is based on calculation of the complex formation function ii, the average number of moles of hydrogen bound per mole of acid: fi = ([HIDA-]
+ 2 [ H,IDA] )/C, = (2C, - aCA - [H+] + [OH-] )/C,
194
where CA is the analytical acid concentration and a is the number of moles of base added per mole of acid. The plot of Iz as a function of pH is shown in Fig. 1. The pH values for which it = 0.5 and 1.5 are taken as equivalent to PK., and pKa2, respectively. The two acidity constants of iminodiacetic acid, calculated as half integral values were refined by using Bjerrum’s convergence method [ll] . The results were pKaI = 2.43 + 0.08 and pK,, = 9.31 + 0.10, in good agreement with the previously reported values. Stability constants for the copper-IDA system In order to evaluate the stability constants of the Cu-IDA system, for which the formation of two mononuclear complexes is reported, it is necessary to find a relationship between the constants and the experimentally determined or calculated variables [M] and [L] (Figs. 2 and 3). The Bjerrum approach, which makes use of the secondary concentration variable ti, the average number of moles of ligand bound per mole of metal, cannot be employed as it does not yield values lower than 0.6 or higher than 1. The relationship may be established by using the secondary concentration which is called the degree of complex formation variable Cp = [M=]/[M], [ 111. The function Q finds application in the Leden method which, notwithstanding a number of limitations, is suitable for calculations from potentiometric data. Leden defined a function f(L):
f(L) = (@- l)/[Ll = ([&I - [Ml I/‘WI [Ll PCU
T 3-
i 2.
1.
1.
I\
6. -8 9-
as
PI
,
a
,
,
a
,
Fig. 1. Formation
,
6
,
,
8
,
10
,
,
12
c
1
Pi
4
8'12
16
20
c 2:
Cu(ml)
function ii of iminodiacetic acid.
Fig. 2. Titration curves of unbuffered disodium iminoacetate with copper. CB = 4.22 X lo-” M (100 ml) and Cm = 5.82 X 10m2 M. Curves: (1) pCu; (2) pH. In this and other figures, subscripts A and B are used to indicate IDA in the acid (A) or disodium salt (B) forms when distinction is necessary; otherwise the subscript IDA is used.
1.
195
ii 1
Fig. 3. Titration curves of iminodiacetic acid and copper mixtures with NaOH. CNaoH = = 1; (2) C,IC,, = 2. 1.03 x lo-* M; Ccu = 1.00 x 10e3 M (100 ml). Curves: (1) C*/Ca Fig. 4. Formation function ii of the protonated complex monoprotic acid: (0) direct titration; (X ) back-titration.
CuH(IDA);
considered as a
which, if the number of complexes formed, N, is 2, becomes f(L) = p1 + & [L] and if f(L) is plotted versus [L] , a straight line is obtained having slope & and an intercept pi on the f(L) axis. The Leden method was applied to the data obtained from both types of titration; a cyclic procedure was adopted to refine the [L] values. Leastsquare straight line plots were obtained, and the following values were derived for the copper complexes: log b1 = 9.32 + 0.05 and log pz = 16.33 f 0.12. The absence of polynuclear species, which are highly improbable, was demonstrated by the fact that pi and pz, calculated from experiments at different concentrations, did not vary within experimental error. Calculations by the Leden method failed to yield straight lines when applied to data obtained at pH lower than 7; this indicates that in acidic media another complex species may be formed. The existence of another copper complex in the acidic region is supported by the difference between the titration with sodium hydroxide of the mixtures with C,/Cti = 2 and C,/C,, = 1 (Fig. 3). The latter merely corresponds, as already noted by Chabereck and Martell [l] , to the reaction Cu2+ + H21DA + 2 OH- + CuIDA + 2 Hz0 In the titration of the mixture containing C,/C, = 2, the first equivalence point is obtained when 3/4 of the total acid equivalents are neutralized and
196
the remaining l/4 is then neutralized, as shown by the appearance of a second equivalence point. This pattern suggests the existence of a protonated species CuH(IDA);, which does not seem to have been considered by other authors. This species is probably formed during the first titration step in accordance with the reaction Cu2+ + 2 HJDA
+ 3 OH- + CuH(IDA);
The second acid titration CuH(IDA);
+ 3 Hz0
step in Fig. 3 corresponds
+ OH- + Cu(IDA):-
therefore
to
+ Hz0
In the pH region 4-9, the Bjerrum method was then applied and the curve reported in Fig. 4 was obtained. The reverse titration was also run by adding standard perchloric acid to a preformed Cu(IDA);- complex solution (C, = 9.20 X 10-j M and CcU = 4.60 X 10s3 M) and the Bjerrum method was applied; the points thus obtained are also reported in Fig, 4. For n = 0.5, the value pK, = 6.53 + 0.10 was found. An interesting trend is shown by the curves reported in Fig. 5, which refer to measurements at constant total ligand concentration and different total copper concentrations: their spacing indicates a well defined dependence, at constant pH in the acidic region, of pCu on the square of CcU and their slopes are close to + 1 for all CcU values. Further, [ Cu( IDA)] and [ CuH(IDA);]
Fig. 5. pCu vs. pH at constant CA = 1.20 X 10m3M. Ccu (M): (1) 2.76 X lo-‘; (2) 5.62 X 10-S; (3) 8.28 x lO-s; (4) 1.12 x 1O-4; (5) 2.76 x lo+; (6) 5.62 X lo+. Fig. 6. Calculated log ( [CU(IDA)]~/[CUH(IDA);]) experimental data. pH 4.01, CA = 1.20 x 10“ M.
vs. log CQ, and pCu vs. log CQ, from
197
were calculated at constant pH, through the known constants; together with experimentally measured pCu values, these concentrations were used to plot the curves reported in Fig. 6. The trend can only be explained through the existence of the equilibrium Cuz+ + CuH(IDA);
= 2 Cu(IDA) + H+
which explains the parallel lines and the slopes of 2. The distribution of the species formed in the Cu-IDA system were calculated at different pH values; the pattern obtained is reported in Fig. 7. Stability constants for the cadmium-IDA system Two stepwise chelate formation constants have been reported in the literature for the Cd-IDA system [ 1, 51. The Leden method was again employed to calculate fll and /I2 for cadmium complexes. Examples of the results of both kinds of titration are reported in Figs. 8 and 9. The curves of Fig. 9 show that half equivalents of the acid are neutralized before any cadmium chelate is formed. Only two stepwise complex formation reactions are assumed and this appears to be valid even at pH < 7: all calculated f(L) values, even in the acidic region, fitted a straight line, from which the values log p1 = 5.48 + 0.04 and log pZ = 9.72 + 0.10 were evaluated for the cadmium complexes. No evidence was found that a species corresponding to CdH(IDA); is formed.
PCd
Fig. 7. &-IDA 7.40 x lo-’ M.
complex
distribution
Fig. 8. Titration curves of unbuffered 9.20 x lo-’ M (100 ml) and CC~ = 4.38
I
as a function
x
of pH. CB = 1.20
X
lo-’ M; Ccu =
disodium iminodiacetate with cadmium. 10ez M. Curves: (1) pCd; (2) pH.
CB=
199 TABLE 1 Acid and chelate equilibrium constants (I = 0.1;
T = 25’C)
Logarithmic constant found
Literature values
Cation
Reaction
H’
H’ + IDA’- + HIDA-
9.31 f 0.10
H+ + HIDA- + HJDA 2 H’ + IDAZ- += H,IDA H+ + Cu(IDA);- = CuH(IDA);
2.43 f 0.08 -
cuz+
Cu’+ + IDA*- + Cu(IDA) Cu(IDA) + IDA*- + Cu(IDA);Cua+ + 21DAa- --’ Cu(IDA);-
9.32 * 0.05 16.33 f 0.12
10.55 [l]; 5.65 [l]; -
Cd”
Cdl’ + IDAz- + Cd(IDA) Cd(IDA) + IDAa- --‘ Cd(IDA):Cdl’ + 21DA2- + Cd(IDA);-
5.48 f 0.04 9.72 f 0.10
5.35 [l]; 4.18 [l]; -
6.53 f 0.10
9.12 [l]; 9.40 [2];9.17 [6]; 9.33 [71 2.54 [l]; 2.50 [2]; 2.58 [7] 11.73 [6] 10.63 [3]; 10.42 [4] 6.05 [5]; 5.60 141 5.54 [5] 4.74 [5]
and [Cd2+] were measured with the appropriate i.s.e. after copper had been added to cadmium/IDA mixtures at pH 4.5. In no way can the data trend shown in Fig. 11 be ascribed to any secondary ion effect, which is usually described through the selectivity coefficient in the modified Nikolski equation [ 121. An exhaustive investigation of this problem was outside the scope of this work, and interest was restricted to practical applications. The important considerations are as follows. First, free copper ion, buffered by an excess of IDA, can be tolerated up to a concentration of 4 X lo-’ M (Fig.
Fig. 11. Response of the cadmium i.s.e. as a function of free copper ion concentration for: (1) Ccc = 6.85 X lo-’ M, Cm* = 2.50 X lo-’ M; (2) Ccd = 1.37 X lo-‘M, CIDA = 2.50 X 10m4M, and CCd = 1.37 X 10m5M, CIDA = 1.22 X 1O-3 M; (3) Cod = 2.74 X 10-6M, CmA = 2.50 X lo-’ M. Free [Cu”] measured with copper i.s.e. Fig. 12. Corresponding “Cd and cxou values at selected pHs.
200
11); the cadmium-selective electrode gives a response proportional to cadmium concentration and equilibrium is attained within reasonable times (= 10 min). Secondly, when the free copper ion concentration exceeds 4 X lo-’ M, equilibration of the electrode response takes such a long time that it becomes analytically useless, and the responses are not obviously correlated to any analytical parameter. In order to assess the limiting conditions for applicability of the cadmiumselective electrode, the ratios of total metal-to-free ion concentrations, (Ye and cCd, were calculated; the values are represented logarithmically in Fig. 12. If it is assumed that, under the conditions adopted, copper is complexed predominantly as CuIDA and CuH(IDA)i whereas cadmium is complexed only as CdIDA, then olcu = ([CL?‘] +
KJ32cu
+ [CuIDA]
+ [CuH(IDA);])/[CX?+]
= 1 + /31c,[IDA2-]
WA’-1 ‘W+l
‘Ax= ([Cd2+]
+ [CdIDA]
(1) )/[Cd’+]
= 1 + fltca [IDA’-]
(2)
The theoretical error on cadmium measurements can be related to the percentage of complexed cadmium, a%, which is a function of CX~ through the following relationship: ((Ye - 1) = a%/(100 - a%). Figure 12 can be used to identify the maximum value of (Ycu attainable by IDA addition without affecting the cadmium measurement by a factor higher than a%. The value of (Ye can be related directly to the potential decrease, AE, measured with a copper-selective electrode after addition of IDA, to reach the correct conditions for cadmium measurement. The value of AE is related to Q,, by the relationship AE = S log CX~, where S indicates the slope of the electrode response. This procedure markedly simplifies the problem of adding the correct amount of ligand, which would be difficult to calculate. However, the limiting condition that the free copper be not higher than 4 X lo-’ M implies also that an upper limit for total copper, Cc,,(max), be assessed. For a cadmium a% value of 0.5, Ccu(max) values were derived from ecu at different pH values (Table 2). Although the cadmium-selective electrode can be used over a wide pH range, hydrogen ions interfere with measurements of low concentrations of cadmium ion, thus pH levels lower than 4 were not considered. A standard addition procedure was devised for the determination of cadmium in the presence of copper and tested on artificial samples. The linearized Gran plots reported in Fig. 13 refer to samples of Cca = 1.37 X lo-’ and 1.37 X lO+ M: the lower concentration may be taken as the limit of determination for the cadmium-selective electrode, at least in the linear response region. The amount of copper added never exceeded the permissible upper limit. Results and related parameters are shown in Table 3. Samples must be prepared suitably before standard additions and potential measurements can be made with the cadmium-selective electrode. This was
201 TABLE 2 Maximum allowed copper concentrations in cadmium samples at selected pH for a fixed percentage of complexed cadmium (a%) of 0.5 Co,( max) (mol I-‘)
PH 4.0 4.5 5.0 5.5
2000 635 210 64.4
8.00 2.54 8.40 2.58
x x x X
lo+ lo4 lo+ 1O-5
TABLE 3 Determination of cadmium ion by the standard addition procedure with sample concentrations and main related parameters Initial cone. (X lo-’ M) Cd=+
cuz+
1.37 1.37 1.37 0.137 0.137 0.137
2.00 0.100 2.00 2.00 0.214 20.0
IDA added (x lo+ M)
Residual [cLP+ 1
log “cu
4.22 0.84 8.44 4.22 1.67 50.0
2.57 X lo-’ 1.60 x lo-’ 6.45 x lo-’ 2.63 X lo-’ 1.25 X lo-’ 2.27 x lo+
1.89 0.79 2.49 1.88 1.23 3.94
log(Qod-1)
-2.76 -3.32 -2.46 -2.77 -3.10 -1.72
a%
0.17 0.05 0.35 0.17 0.08 1.87
Fig. 13. Gran plots in the standard addition procedure applied to: (A) Cod = 1.37 X 10~‘M and (B) C&r = 1.37 x 10m6M in the presence of different amounts of copper ion and IDA. (A) With additions of 1.00 x 10e2 M cadmium solution: (1) Cc,, = 2.00 X lo-’ M, CmA = 4.22 X lo- M; (2) Ccu = 1.00 X lo+ M, Cr,,A = 8.44 x lo-’ M; (3) Ccu = 2.00 x lo-’ M, Cm, = 8.44 X low4 M. (B) With additions of 1.00 X low3 M cadmium solution: (1) Cou = 2.00 X lo-” M, CmA = 4.22 X lo-’ M; (2) Cc” = 2.14 X 1W6 M, Cl,,_,, = 1.67 X lo-‘M; (3) Ccu = 2.00 X lo-’ M, C~A = 5.00 x lo-’ M. V, = 100 ml.
202
done by adding the ligand and recording the corresponding potential drop at the copper-selective electrode; the addition was stopped when the reading corresponded to a free ion concentration [ Cu2’] < 4 X lo-’ M. The electrode commutator was then switched to the cadmium-selective electrode and standard additions of cadmium ion were made. As shown in Fig. 13, the slopes of the Gran plots vary depending on the residual copper ion activity of the samples but, although the abscissa units may differ, the relative positions of the lines are maintained. The last column of Table 3 shows the a% values derived from acU; they indicate that the percentage of complexed cadmium is very low in all cases, reaching a maximum of 1.87% in the last sample, which was measured with the largest actual experimental error (?12%). This work was supported by the Consiglio Nazionale delle Richerche as part of the Special Research Project “Environmental Quality Promotion”; the CNR also gave permission for publication. The authors are grateful to S. Meloni for valuable discussions concerning this work. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12
S. Chabereck, Jr. and A. E. Martell, J. Am. Chem. Sot., 74 (1962) 5052. K. S. Rajan and A. E. Martell, J. Inorg. Nucl. Chem., 26 (1964) 789. G. Anderegg, Helv. Chim. Acta, 47 (1964) 1801. R. P. Bonomo, R. Calf, F. Riggi, E. Rizzarelli, S. Sammartano and G. Siragusa, Inorg. Chem., 18 (1980) 3417. T. Ait-Hamouda and M. J. Schwing-Weill, Analusis, 9 (1981) 93. A. Liberti and A. Napoli, J. Inorg. Nucl. Chem., 33 (1971) 89. L. C. Thompson, Inorg. Chem., 1 (1962) 490. H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolyte Solutions, Reinhold, New York, 1958. R. Stella and M. T. Ganzerli Valentini, Anal. Chem., 51 (1979) 2148. J. RfbZiEkaand E. H. Hansen, Anal. Chim. Acta, 63 (1973) 115. See, e.g., F. R. Hartley, C. Burgess and R. M. Alcock, Solution Equilibria, Horwood, Chichester, 1980. G. Eisenman, in R. A. Durst (Ed.), Ion-Selective Electrodes, NBS Spec. Publ. No. 314, Washington, DC, 1969.
A STUDY OF COPPER AND CADMIUM IMINODIACETATE COMPLEXES BY ION-SELECTIVE ELECTRODES AND APPLICATION TO CADMIUM MONITORING
R. STELLA* and M. T. GANZERLI
VALENTINI
Dipartimento di Chimica Generale and Centro di Radiochimica ed Analisl per Attivazione de1 C.N.R., Universitci di Pavia, Viale Taramelll 12, 27100 Pavia (Italy) (Received 6th March 1983)
SUMMARY The chelating properties of the iminodiacetate ion (IDAl-) towards copper(I1) and cadmium(I1) were investigated through pM and pH measurements at constant ionic strength of 0.1 M and at 25°C in the pH range 4-9. The acidity constants found were in good agreement with those already reported but significant differences from reported values were obtained for the stepwise copper complex formation constants, though the overall constant was in good agreement. Evidence is given for the existence of the protonated species CuH (IDA);, which proved helpful in explaining the anomalous complexing capacity of IDA in acidic medium. Practical applications include the use of IDA for lowering the level of copper interference in monitoring cadmium ion with a cadmiumselective solid-state electrode; this provides reliable measurements which are otherwise impossible.
The reported stability constants of complexes of the iminodiacetate ion (IDA’-) with copper(I1) and cadmium(I1) ions [l-5] are sufficiently different to suggest the possibility that this ligand may be employed as a selective masking agent for copper in the presence of cadmium. Yet measurements conducted on both the Cu-IDA and Cd-IDA systems with ion-selective electrodes (i.s.e.) showed discrepancies from the values calculated from the available constants. For instance, in the acidic region the extent of copper complexation was greater than the calculated value, thus suggesting the presence of a protonated complex. Therefore it seemed worthwhile to investigate these systems by coupling i.s.e.‘s to pH measurements; this technique allows a more detailed insight into metal complex equilibria than is possible by previous methods. A preliminary investigation was also made on iminodiacetic acid dissociation constants and the values obtained were compared with those already reported [ 1,2,6,7].
0003-2670/83/$03.00
0 1983 Elsevier Science Publishers B.V.
192 EXPERIMENTAL
Reagents Sodium iminodiacetate (Fluka) was purified by double recrystallization from methanol and dried at 115°C. The corresponding acid was prepared by passing a 0.1 M disodium salt solution through two columns in series (2 cm i.d., 20 cm high) filled with Dowex 5OW-X4 resin converted to the acid form with 0.1 M perchloric acid, and repeatedly washing with deionized and twicedistilled water. Only the central fractions of the eluate were collected and the absence of sodium was checked by atomic emission spectrometry; it is very important to discard all head and tail fractions, as they may contain undesirable ligands. The acid was then diluted and potentiometrically titrated with standard sodium hydroxide or standard copper solution at pH 7 using the corresponding i.s.e. The carbonate-free 0.1 M sodium hydroxide was prepared from a 50% (w/w) solution and standardized against potassium hydrogenphthalate. A 0.1 M perchloric acid solution was prepared from the 70% acid (Merck, Suprapur) and standardized against sodium carbonate. A 10m2 M copper stock solution was prepared by dissolving a suitable amount of copper(I1) perchlorate hexahydrate (ultrapure, Alfa Products) and standardized by titration with EDTA; end-points were detected visually and potentiometrically, with murexide as indicator for the former. A 10m2M cadmium stock solution was prepared by dissolving a weighed amount of cadmium wire (ultrapure, Alfa Products) in diluted nitric acid (Suprapur, Merck) in a teflon beaker. The excess of nitric acid was evaporated on a hot plate and the residue was taken up with water. The resulting solution was standardized by titration with EDTA; both potentiometric and visual (Calcon indicator) end-points were obtained in ammoniacal medium. The water used in all experiments was deionized and then twice-distilled in a quartz apparatus. Instrumentation A digital millivoltmeter (Orion model 701) was used in all measurements. A combined glass electrode (model 91) was used for pH measurements. The glass electrode was calibrated as suggested by Rajan and Martell [2] by direct titration of acetic acid; the observed readings were compared with the actual hydrogen ion concentrations tabulated by Harned and Owen [ 81. In the pH region below 3.5, the meter was calibrated by measurements of solutions with known concentrations of perchloric acid. The concentration of the hydroxide ion was obtained by using a p(K,), value of 13.79 at 25°C in 0.1 M potassium nitrate. The copper ion-selective electrode (Orion model 94-29A) was used with a single-junction Orion 90-01 reference electrode. The electrode was carefully calibrated at ionic strength 0.1 M with Cu-EDTA buffered solutions as described in a previous paper [9] ; a tenfold change in the copper concentration yielded an electrode response change of 29.0 + 0.5 mV, as expected.
193
The cadmium ion-selective electrode (Orion 94-48A) was used similarly and calibrated following the analogous procedure for copper: the mV response plotted in a semilog scale versus the free ion concentration was found to be linear down to lo-” M in metal buffer solutions as in the procedure suggested by R8iiEka and Hansen [lo]. Procedure All experiments were done at 25°C in a constant temperature bath, the ionic strength (I) being adjusted to 0.1 M with potassium nitrate (Suprapur, Merck). The iminodiacetic acid and its disodium salt were titrated under a nitrogen atmosphere, by first bubbling a nitrogen stream for 30 min and then maintaining the stream in the cell over the solution. Solutions having the following concentrations were used: CA = 4.50 X lo-* M for the acid and CB = 8.44 X lo-* M or 9.20 X 10m3M for the disodium salt. The copper-IDA and the cadmium-IDA systems were separately investigated by simultaneously measuring the free metal ion concentration and the pH in two different sets of experiments. First, titrations of unbuffered IDA (disodium salt) solutions were done with standard copper or cadmium solutions, with Co, = 5.82 X lo-’ M, Ccd = 4.38 X lo-* M and CB = 4.22 X 10m3M or 9.20 X lo-’ M. Next, solutions of ligand and metal ion in 1:l and 2:l molar ratios were titrated with standard 0.1 M sodium hydroxide; here CA = 2.11 X lo4 M, 4.22 X 10m3M and 1.00 X 10m3M for the copper system, and CA = 3.88 X lo-’ M, 7.39 X 10W3M and 3.68 X lo-* M for the cadmium system. The results in both cases are presented in graphic form with pH and pCu or pCd plotted as a function of the titrant added. RESULTS
AND DISCUSSION
Evaluation of acidity constants Considering the iminodiacetic acid as a diprotic acid, Chabereck and Martell [l] found p& = 2.54 and pK,* = 9.12 (I = 0.1, T = 30°C); Rajan and Martell [2] reported p& = 2.50 and pK,, = 9.40 (I = 0.1, T = 25°C) and Thompson [7] found pK,, = 2.58 and pK,, = 9.33 (I = 0.1, T = 25°C). All these values agree well, even with those that may be derived from the results of Liberti and Napoli [6] : log p1,2 = 9.17 and log p2,* = 11.73 (I= 0.5, T = 25°C). Liberti and Napoli also gave a third formation constant related to a H,IDA+ form, which begins to play a significant r61e at pH < 2.5. In this paper, the evaluation of the first two constants was considered sufficient to allow an accurate description of the investigated system. The acidity constants were calculated by using the well known Bjerrum method which is based on calculation of the complex formation function ii, the average number of moles of hydrogen bound per mole of acid: fi = ([HIDA-]
+ 2 [ H,IDA] )/C, = (2C, - aCA - [H+] + [OH-] )/C,
194
where CA is the analytical acid concentration and a is the number of moles of base added per mole of acid. The plot of Iz as a function of pH is shown in Fig. 1. The pH values for which it = 0.5 and 1.5 are taken as equivalent to PK., and pKa2, respectively. The two acidity constants of iminodiacetic acid, calculated as half integral values were refined by using Bjerrum’s convergence method [ll] . The results were pKaI = 2.43 + 0.08 and pK,, = 9.31 + 0.10, in good agreement with the previously reported values. Stability constants for the copper-IDA system In order to evaluate the stability constants of the Cu-IDA system, for which the formation of two mononuclear complexes is reported, it is necessary to find a relationship between the constants and the experimentally determined or calculated variables [M] and [L] (Figs. 2 and 3). The Bjerrum approach, which makes use of the secondary concentration variable ti, the average number of moles of ligand bound per mole of metal, cannot be employed as it does not yield values lower than 0.6 or higher than 1. The relationship may be established by using the secondary concentration which is called the degree of complex formation variable Cp = [M=]/[M], [ 111. The function Q finds application in the Leden method which, notwithstanding a number of limitations, is suitable for calculations from potentiometric data. Leden defined a function f(L):
f(L) = (@- l)/[Ll = ([&I - [Ml I/‘WI [Ll PCU
T 3-
i 2.
1.
1.
I\
6. -8 9-
as
PI
,
a
,
,
a
,
Fig. 1. Formation
,
6
,
,
8
,
10
,
,
12
c
1
Pi
4
8'12
16
20
c 2:
Cu(ml)
function ii of iminodiacetic acid.
Fig. 2. Titration curves of unbuffered disodium iminoacetate with copper. CB = 4.22 X lo-” M (100 ml) and Cm = 5.82 X 10m2 M. Curves: (1) pCu; (2) pH. In this and other figures, subscripts A and B are used to indicate IDA in the acid (A) or disodium salt (B) forms when distinction is necessary; otherwise the subscript IDA is used.
1.
195
ii 1
Fig. 3. Titration curves of iminodiacetic acid and copper mixtures with NaOH. CNaoH = = 1; (2) C,IC,, = 2. 1.03 x lo-* M; Ccu = 1.00 x 10e3 M (100 ml). Curves: (1) C*/Ca Fig. 4. Formation function ii of the protonated complex monoprotic acid: (0) direct titration; (X ) back-titration.
CuH(IDA);
considered as a
which, if the number of complexes formed, N, is 2, becomes f(L) = p1 + & [L] and if f(L) is plotted versus [L] , a straight line is obtained having slope & and an intercept pi on the f(L) axis. The Leden method was applied to the data obtained from both types of titration; a cyclic procedure was adopted to refine the [L] values. Leastsquare straight line plots were obtained, and the following values were derived for the copper complexes: log b1 = 9.32 + 0.05 and log pz = 16.33 f 0.12. The absence of polynuclear species, which are highly improbable, was demonstrated by the fact that pi and pz, calculated from experiments at different concentrations, did not vary within experimental error. Calculations by the Leden method failed to yield straight lines when applied to data obtained at pH lower than 7; this indicates that in acidic media another complex species may be formed. The existence of another copper complex in the acidic region is supported by the difference between the titration with sodium hydroxide of the mixtures with C,/Cti = 2 and C,/C,, = 1 (Fig. 3). The latter merely corresponds, as already noted by Chabereck and Martell [l] , to the reaction Cu2+ + H21DA + 2 OH- + CuIDA + 2 Hz0 In the titration of the mixture containing C,/C, = 2, the first equivalence point is obtained when 3/4 of the total acid equivalents are neutralized and
196
the remaining l/4 is then neutralized, as shown by the appearance of a second equivalence point. This pattern suggests the existence of a protonated species CuH(IDA);, which does not seem to have been considered by other authors. This species is probably formed during the first titration step in accordance with the reaction Cu2+ + 2 HJDA
+ 3 OH- + CuH(IDA);
The second acid titration CuH(IDA);
+ 3 Hz0
step in Fig. 3 corresponds
+ OH- + Cu(IDA):-
therefore
to
+ Hz0
In the pH region 4-9, the Bjerrum method was then applied and the curve reported in Fig. 4 was obtained. The reverse titration was also run by adding standard perchloric acid to a preformed Cu(IDA);- complex solution (C, = 9.20 X 10-j M and CcU = 4.60 X 10s3 M) and the Bjerrum method was applied; the points thus obtained are also reported in Fig, 4. For n = 0.5, the value pK, = 6.53 + 0.10 was found. An interesting trend is shown by the curves reported in Fig. 5, which refer to measurements at constant total ligand concentration and different total copper concentrations: their spacing indicates a well defined dependence, at constant pH in the acidic region, of pCu on the square of CcU and their slopes are close to + 1 for all CcU values. Further, [ Cu( IDA)] and [ CuH(IDA);]
Fig. 5. pCu vs. pH at constant CA = 1.20 X 10m3M. Ccu (M): (1) 2.76 X lo-‘; (2) 5.62 X 10-S; (3) 8.28 x lO-s; (4) 1.12 x 1O-4; (5) 2.76 x lo+; (6) 5.62 X lo+. Fig. 6. Calculated log ( [CU(IDA)]~/[CUH(IDA);]) experimental data. pH 4.01, CA = 1.20 x 10“ M.
vs. log CQ, and pCu vs. log CQ, from
197
were calculated at constant pH, through the known constants; together with experimentally measured pCu values, these concentrations were used to plot the curves reported in Fig. 6. The trend can only be explained through the existence of the equilibrium Cuz+ + CuH(IDA);
= 2 Cu(IDA) + H+
which explains the parallel lines and the slopes of 2. The distribution of the species formed in the Cu-IDA system were calculated at different pH values; the pattern obtained is reported in Fig. 7. Stability constants for the cadmium-IDA system Two stepwise chelate formation constants have been reported in the literature for the Cd-IDA system [ 1, 51. The Leden method was again employed to calculate fll and /I2 for cadmium complexes. Examples of the results of both kinds of titration are reported in Figs. 8 and 9. The curves of Fig. 9 show that half equivalents of the acid are neutralized before any cadmium chelate is formed. Only two stepwise complex formation reactions are assumed and this appears to be valid even at pH < 7: all calculated f(L) values, even in the acidic region, fitted a straight line, from which the values log p1 = 5.48 + 0.04 and log pZ = 9.72 + 0.10 were evaluated for the cadmium complexes. No evidence was found that a species corresponding to CdH(IDA); is formed.
PCd
Fig. 7. &-IDA 7.40 x lo-’ M.
complex
distribution
Fig. 8. Titration curves of unbuffered 9.20 x lo-’ M (100 ml) and CC~ = 4.38
I
as a function
x
of pH. CB = 1.20
X
lo-’ M; Ccu =
disodium iminodiacetate with cadmium. 10ez M. Curves: (1) pCd; (2) pH.
CB=
199 TABLE 1 Acid and chelate equilibrium constants (I = 0.1;
T = 25’C)
Logarithmic constant found
Literature values
Cation
Reaction
H’
H’ + IDA’- + HIDA-
9.31 f 0.10
H+ + HIDA- + HJDA 2 H’ + IDAZ- += H,IDA H+ + Cu(IDA);- = CuH(IDA);
2.43 f 0.08 -
cuz+
Cu’+ + IDA*- + Cu(IDA) Cu(IDA) + IDA*- + Cu(IDA);Cua+ + 21DAa- --’ Cu(IDA);-
9.32 * 0.05 16.33 f 0.12
10.55 [l]; 5.65 [l]; -
Cd”
Cdl’ + IDAz- + Cd(IDA) Cd(IDA) + IDAa- --‘ Cd(IDA):Cdl’ + 21DA2- + Cd(IDA);-
5.48 f 0.04 9.72 f 0.10
5.35 [l]; 4.18 [l]; -
6.53 f 0.10
9.12 [l]; 9.40 [2];9.17 [6]; 9.33 [71 2.54 [l]; 2.50 [2]; 2.58 [7] 11.73 [6] 10.63 [3]; 10.42 [4] 6.05 [5]; 5.60 141 5.54 [5] 4.74 [5]
and [Cd2+] were measured with the appropriate i.s.e. after copper had been added to cadmium/IDA mixtures at pH 4.5. In no way can the data trend shown in Fig. 11 be ascribed to any secondary ion effect, which is usually described through the selectivity coefficient in the modified Nikolski equation [ 121. An exhaustive investigation of this problem was outside the scope of this work, and interest was restricted to practical applications. The important considerations are as follows. First, free copper ion, buffered by an excess of IDA, can be tolerated up to a concentration of 4 X lo-’ M (Fig.
Fig. 11. Response of the cadmium i.s.e. as a function of free copper ion concentration for: (1) Ccc = 6.85 X lo-’ M, Cm* = 2.50 X lo-’ M; (2) Ccd = 1.37 X lo-‘M, CIDA = 2.50 X 10m4M, and CCd = 1.37 X 10m5M, CIDA = 1.22 X 1O-3 M; (3) Cod = 2.74 X 10-6M, CmA = 2.50 X lo-’ M. Free [Cu”] measured with copper i.s.e. Fig. 12. Corresponding “Cd and cxou values at selected pHs.
200
11); the cadmium-selective electrode gives a response proportional to cadmium concentration and equilibrium is attained within reasonable times (= 10 min). Secondly, when the free copper ion concentration exceeds 4 X lo-’ M, equilibration of the electrode response takes such a long time that it becomes analytically useless, and the responses are not obviously correlated to any analytical parameter. In order to assess the limiting conditions for applicability of the cadmiumselective electrode, the ratios of total metal-to-free ion concentrations, (Ye and cCd, were calculated; the values are represented logarithmically in Fig. 12. If it is assumed that, under the conditions adopted, copper is complexed predominantly as CuIDA and CuH(IDA)i whereas cadmium is complexed only as CdIDA, then olcu = ([CL?‘] +
KJ32cu
+ [CuIDA]
+ [CuH(IDA);])/[CX?+]
= 1 + /31c,[IDA2-]
WA’-1 ‘W+l
‘Ax= ([Cd2+]
+ [CdIDA]
(1) )/[Cd’+]
= 1 + fltca [IDA’-]
(2)
The theoretical error on cadmium measurements can be related to the percentage of complexed cadmium, a%, which is a function of CX~ through the following relationship: ((Ye - 1) = a%/(100 - a%). Figure 12 can be used to identify the maximum value of (Ycu attainable by IDA addition without affecting the cadmium measurement by a factor higher than a%. The value of (Ye can be related directly to the potential decrease, AE, measured with a copper-selective electrode after addition of IDA, to reach the correct conditions for cadmium measurement. The value of AE is related to Q,, by the relationship AE = S log CX~, where S indicates the slope of the electrode response. This procedure markedly simplifies the problem of adding the correct amount of ligand, which would be difficult to calculate. However, the limiting condition that the free copper be not higher than 4 X lo-’ M implies also that an upper limit for total copper, Cc,,(max), be assessed. For a cadmium a% value of 0.5, Ccu(max) values were derived from ecu at different pH values (Table 2). Although the cadmium-selective electrode can be used over a wide pH range, hydrogen ions interfere with measurements of low concentrations of cadmium ion, thus pH levels lower than 4 were not considered. A standard addition procedure was devised for the determination of cadmium in the presence of copper and tested on artificial samples. The linearized Gran plots reported in Fig. 13 refer to samples of Cca = 1.37 X lo-’ and 1.37 X lO+ M: the lower concentration may be taken as the limit of determination for the cadmium-selective electrode, at least in the linear response region. The amount of copper added never exceeded the permissible upper limit. Results and related parameters are shown in Table 3. Samples must be prepared suitably before standard additions and potential measurements can be made with the cadmium-selective electrode. This was
201 TABLE 2 Maximum allowed copper concentrations in cadmium samples at selected pH for a fixed percentage of complexed cadmium (a%) of 0.5 Co,( max) (mol I-‘)
PH 4.0 4.5 5.0 5.5
2000 635 210 64.4
8.00 2.54 8.40 2.58
x x x X
lo+ lo4 lo+ 1O-5
TABLE 3 Determination of cadmium ion by the standard addition procedure with sample concentrations and main related parameters Initial cone. (X lo-’ M) Cd=+
cuz+
1.37 1.37 1.37 0.137 0.137 0.137
2.00 0.100 2.00 2.00 0.214 20.0
IDA added (x lo+ M)
Residual [cLP+ 1
log “cu
4.22 0.84 8.44 4.22 1.67 50.0
2.57 X lo-’ 1.60 x lo-’ 6.45 x lo-’ 2.63 X lo-’ 1.25 X lo-’ 2.27 x lo+
1.89 0.79 2.49 1.88 1.23 3.94
log(Qod-1)
-2.76 -3.32 -2.46 -2.77 -3.10 -1.72
a%
0.17 0.05 0.35 0.17 0.08 1.87
Fig. 13. Gran plots in the standard addition procedure applied to: (A) Cod = 1.37 X 10~‘M and (B) C&r = 1.37 x 10m6M in the presence of different amounts of copper ion and IDA. (A) With additions of 1.00 x 10e2 M cadmium solution: (1) Cc,, = 2.00 X lo-’ M, CmA = 4.22 X lo- M; (2) Ccu = 1.00 X lo+ M, Cr,,A = 8.44 x lo-’ M; (3) Ccu = 2.00 x lo-’ M, Cm, = 8.44 X low4 M. (B) With additions of 1.00 X low3 M cadmium solution: (1) Cou = 2.00 X lo-” M, CmA = 4.22 X lo-’ M; (2) Cc” = 2.14 X 1W6 M, Cl,,_,, = 1.67 X lo-‘M; (3) Ccu = 2.00 X lo-’ M, C~A = 5.00 x lo-’ M. V, = 100 ml.
202
done by adding the ligand and recording the corresponding potential drop at the copper-selective electrode; the addition was stopped when the reading corresponded to a free ion concentration [ Cu2’] < 4 X lo-’ M. The electrode commutator was then switched to the cadmium-selective electrode and standard additions of cadmium ion were made. As shown in Fig. 13, the slopes of the Gran plots vary depending on the residual copper ion activity of the samples but, although the abscissa units may differ, the relative positions of the lines are maintained. The last column of Table 3 shows the a% values derived from acU; they indicate that the percentage of complexed cadmium is very low in all cases, reaching a maximum of 1.87% in the last sample, which was measured with the largest actual experimental error (?12%). This work was supported by the Consiglio Nazionale delle Richerche as part of the Special Research Project “Environmental Quality Promotion”; the CNR also gave permission for publication. The authors are grateful to S. Meloni for valuable discussions concerning this work. REFERENCES 1 2 3 4 5 6 7 8 9 10 11 12
S. Chabereck, Jr. and A. E. Martell, J. Am. Chem. Sot., 74 (1962) 5052. K. S. Rajan and A. E. Martell, J. Inorg. Nucl. Chem., 26 (1964) 789. G. Anderegg, Helv. Chim. Acta, 47 (1964) 1801. R. P. Bonomo, R. Calf, F. Riggi, E. Rizzarelli, S. Sammartano and G. Siragusa, Inorg. Chem., 18 (1980) 3417. T. Ait-Hamouda and M. J. Schwing-Weill, Analusis, 9 (1981) 93. A. Liberti and A. Napoli, J. Inorg. Nucl. Chem., 33 (1971) 89. L. C. Thompson, Inorg. Chem., 1 (1962) 490. H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolyte Solutions, Reinhold, New York, 1958. R. Stella and M. T. Ganzerli Valentini, Anal. Chem., 51 (1979) 2148. J. RfbZiEkaand E. H. Hansen, Anal. Chim. Acta, 63 (1973) 115. See, e.g., F. R. Hartley, C. Burgess and R. M. Alcock, Solution Equilibria, Horwood, Chichester, 1980. G. Eisenman, in R. A. Durst (Ed.), Ion-Selective Electrodes, NBS Spec. Publ. No. 314, Washington, DC, 1969.