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Composition and stability of some metal citrate and diglycolate complexes in aqueous solution
nnaryrrcn C/1fr,tica Acta. 74 (1975) 101-106 0 Elscvicr Scicntifk Publishing Compuny,
COMPOSITION DIGLYCOLATE TERRENCE
Amsterdam
AND STABILITY OF SOME COMPLEXES IN AQUEOUS
B, FIELD.
Depurtowu
of Clwmisrry.
(Rcccivcd
3rd June 1974)
JOHN Uhwrsity
COBURN.
JANET
r>J Worerloo.
101
- Printed in The Ncthcrlnnds
L.
Waterloo.
METAL CITRATE SOLUTION MCCOURT
nnd W.
AND
A. E.
McBRYDE
Ontcwlo N.?L 3GI (Cancrtltr)
There has been some renewed interest in the metal complexes formed by citrate ion in aqueous solution, arising from the consideration of citric acid as a potential detergent builder. Diglycolic acid has also been considered for the same purpose. and corresponding data for the stability of its complexes would be useful. Of these ligands. citrate at least is known to form protonated species MHL or MH2L where HJL represents citric acid. When complex formation is measured by proton displacement in the manner first advocated by Bjey,rum’. interpretation of the experimental results is made more difficult when an unknown amount of the hydrogen-ion balance is incorporated in such protonated complexes. The authors have had access to the powerful computer program SCOGS2 which utilizes pH-titration data for systems consisting of metal, acid and ligand, and fits stability constants for proposed complex species to the data by a non-linear least-squaresmethod. The special advantage of this computer method for the systems to be studied is that it can handle protonated or deprotonated complex species, hydrolysis products, and so forth. Accordingly suitable pH titrations of solutions containing citric or diglycolic acid have been done with and without metal present, and with an added indifferent electrolyte (KNOJ) added to try to achieve constant ionic strength. The results of these titrations were then interpreted in terms of possible complex species, those complexes and their formation constants being preferred and reported which gave minimum sum of squares of deviations between calculated and experimental titres. EXPERlMENTAL
The pH titrations were carried out in a vessel thermostatted at 25.O”C (ref. 3) with standard alkali or adid added by a calibrated syringe with micrometer drive, The solutions of metal nitrates were prepared from analytical reagent salts and standardized by EDTA titrations 4. Molar alkali solution for titration was prepared bv dilution of ampoules of B.D.H. concentrated volumetric solution, standardized with acid phthalate, and stored under COz-free nitrogen. Recrystallized potassium nitrate was added to all solutions as background electrolyte at a molarity of 0.1. The citric acid (B.D.H.) and diglycolic acid (J. T. Baker) were of reagent grade, and were used without further purification when their suff?cient purity had been established by titration. The pH measurements were made either with an Orion Model 801 meter
T. B. FIELD.
102
J. COBURN.
.I. L. MCCOURT.
W. A. E. McBRYDE
equipped with Beckman electrodes or a Radiometer Model PHM52 meter with Radiometer electrodes. The pH scale WLIS established for these two assemblies with N.B.S. buffer solutions. and any observed departure from the Nernst slope for readings between pairs of buffers was compensated by means of the tcmperaturc-adjustment control on each meter. All pH-meter readings were converted to hydrogen-ion concentrations by means of nn empirically determined factor f = 10-F”/[H]. measured under the sdme conditions and with the same electrodes as applied lo particular titrations5. During titrations the apparatus was continuously purged with CO,-free nitrogen. In the absence of metal ion, the pH titrations yield acidity constunts for the ligands. and when metal ions are prcscnt the titration curve is modified according to the amount of hydrogen ion set free by complex formation. In tither case the observations yield stoichiometric equilibrium constants valid within the particular salt background used. It may be noted in connection with acidity constants that these are here given as protonation constants, and numbered in order of increasing numbers of protons bound. This practice places these constants in a formally analogous relationship with the numbering of complex stability constants. and aids in the writing of generalized statements in computer programming. It may bc mentioned that the titrations to determine metal complex stability constants were carried out with an excess of metal over the ligand present in order to secure a suflicient degree of conversion to complexes. In the inverse situation, the titration curve of citric acid is not sufliciently displaced by the presence of a smaller amount of metal to give a precise indication of the extent of complex formation. It follows that under the experimentul conditions used in this study, no evidence of complexes with higher than 1 :l stoichiometry was detected. RESULTS
AND
DISCUSSION
The results obtained from this study are summarized in Tables I-III. They are reported as averages from the data of several titrations in each case, together with a standard deviation on each reported constant. In the various titrations from which results are combined the concentratidns of metal (if present) :tnd TABLE ACIDITY
1 CONSTANTS
(Conwntr;!tion
quolicnts.
FOR
CITRIC
2S’C.
0.1
M KNO,
IO{\ /II’ -._--
----Citric acid Diglycolic acid
5.70 * 0.02 3.95_to.o1
AND
DIGLYCOLIC
ACIDS
background)
IOWA 11:’ --_._-___--~.-
log /I:’
---.
10.06i- 0.02 6.73 & 0.0
12.87 + 0.08
I
ligand were systematically varied. nnd there was no evidence that any of the /1 vuiues derived from the experimental results showed any dependence on the overall concentration of the reagents. Typically, the levels of metal or ligand concentration were in the range 1-G. 10WJ M.
METAL
CITRATE
TABLE
II
STARILITY
AND
DIGLYCOLATE
CONSTANTS
FOR
25’C. (Conccntrulion quotients. --_._..-_.-.-.---.--. -_----..--.-..
0.
METAL ---..---..--.-
103
CITRATES
I M KNO.1
IV1i-I 1, h /I I I I -._-_ . ..__._- ..- -._-_.-_ --_..- .-.--.--.8*02_eO. I5 Culciuni 7.66&O. I9 Magnesium X.08 _e 0.07 Zinc 9.04-e 0.07 Nickel __ ____ _. __-_-._-__-.-._-_-_ ._..-..
TABLE
COMPLEXES
bilckgrottnd) - ------.--.-._..___
---.-
.-_._-
_... _..._-_..-- ._.-._.
MI, hl II I 0 I -.-.--------. -- - _ _- ..- - ..-_. ..-- -...._-._-_-._,.SO&-0.06 3.38 f. 0.07 5. IO-r_ 0.05 5.40 + 0.04 _. ._..____._.__.-.._ .__.._...__ -_._ ._.... ..-_..-. -.
.-._-
_... .-_
_ __._._._-_.
___._. ._--.
.-_-_.- .._.-- .._..- ..._..-..
III
STABILITY (Concentration ___ ..__._._-_.--
CONSTANTS
FOR
METAL
DIGLYCOLATES
25°C. 0. I M KNO., bwkground) quolicnts. .._-. .-- .___._..-.._- ._..._._ _-._ ._._..__ -__..._____.__._... -. -.-- ...____- _._......_. --.--
_ .___ -._---.-.._-_.-_----
MIIL ht /{I I L
ML Iv/ P I I,I ..__-._- ._-._-.-.._.-.-__-..-.-
6.56 + 0.04 Calcium 538 _e 0.02 M:lgncsium 6.03 +- 0.05 Zinc ____ _._. _--___-.-.._ .._ ---.-
X54+0.02 2.15~0.01 3.59*0.01 _..... - ___._-- __.._._---.__
_....-..
-... _.._....... -.. __--
._._- _.__. -. _.- .__.-..________
.__-_--
._.. -.._ ._.__..-.
___.
__...-._-._ -_-.---
___ __ __.__ ._... - _... -._.-
The values of the equilibrium constants arc those which give the minimum standard dcviution in litre (s.d.t.) for the titrations from which they arc derived. They also refer to the choice of species. from various possibilities tried out. which give the lowest s.d.t. Thus. for instance. in the case of diglycolic acid the complexes ML and MHL taken together gave in every titration or combination of titrations ;I lower s.d.t. than that calculated when ML wys the only species assumed to be prcscnt. This is the evidence on which our assumption of the formation of protonatcd diglycolate species is bused. The acidity constnnts listed in Tuble I are concentration quotients ilpplicablc only in solutions of comparable ionic strc-ngth. Owing to the multiple charge borne by some of the ions gcneratcd by thcsc acids. and the effect of this on activity cocflicicnts. the stoichiomctric ionization constants. pi~rtiCUlilrly of citric acid. show ii marked dependence on the ionic strength (c:/‘. thermodynamic values”). The values reported here appear to be in sittisfuctory accord with those reported under nearly comparable conditions by Campi et al.’ and by Li e/ trl.“. A recent summary of the literature on citrate equilibria by Pearce” including his own measurements lists five. sets of stoichiometric acidity constants for ionic strength 0.1 und at specified tcmpcrntures ranging from 20 to 30°C. The present results for pK I are slightly lower. but the others are in excellent agrecmcnt with this summary. Pearce has applied a correction to the values of pKJ for complexation of citrate ion by potassium or sodium ion: this refinement has not been introduced in the present work. With respect to the citrate complexes, it is clear that protonated species do occur, certainly. MI-IL ilnd possibly MH2L. The titration data failed to product
104
T. B. FIELD.
J. COBURN.
J. L. MCCOURT.
W. A. E. McBRYDE
evidence to support the existence of the latter, but Campi et ~1.’ and Pearce9 have reported stability constants for this with certain metals. The pH-titration method for studying metal complex formation is much less capable of defining the state of equilibrium when protonated species occur. This can be seen from the relationship GI-[HI + [OH] = ~qljpqr[M]PIH]rl[L]r+CjljlOIH]JIL] in which Cu is the total net concentration of ionizable hydrogen at any point. If q=O throughout, [L] is explicitly known, given the /$’ values. If y#O, the above relationship does not define [L], and one must resort to best-fit solutions of several mass-balance equations to define [M] and [L]. Experience has shown that in such cases various sets of p. y, and I’ can be found which satisfy the data roughly as well, so that a clear choice of constituent species cannot be made. It is a further limitation on any method of determining equilibrium constants that there be a sufficient and thus precisely delined concentration of each species involved. In the case of the species MH2L, if it exists, the pK for its dissociation to MHL may be so small that to obtain a reasonably amount of this complex would require working at a pH so low as to preclude much formation of any metal complex. With the aid of Pearce’s formation constants and the computer program SPECON”, the concentration of each species formed between pH 2 and pH 6 was calculated for a number of selected values of Chl and C,. For the case where CL= I .O * IO- 3 and C, = 5.0*10-’ (corresponding to one of our titrations), the maximum value of [MH,L] was 3.8. 10 -5 M (at pH 3.5). This is not enough conversion of either ligand or metal to yield a precise definition of /3izl. Even when C,., was raised to 1.0. 10s2, only 7”/, of the ligand was converted to [MH,L] at its maximum concentration. The values of log fi rol for calcium and magnesium complexes are in good agreement with those of Campi et al.’ and Pearceg, though slightly higher than those of Li at ~1.~. The corresponding values for nickel and zinc accord with those reported by Campi and coworkers. The present values for log /jr 1 1 for these complexes tend throughout to be slightly higher than Campi’s, and *rather closer to Pearce’s for the first two metals. The discrepancy may be linked to the inability to characterize MHzL as a discrete species: if some of this was formed in the present experiments, the effect on the data would be to augment the estimate of [MHL] and hence of /I,, 1. Values of stability constants of the diglycolate complexes may be compared with those previously published by Tichane and Bennett’ ’ and by Yasuda et ~1.‘~. Apart from the indication of protonated species at low pH and the proposed stability constants for these, the results appear to be in general agreement with earlier work. It is not clear why the value of fl io1 for magnesium digiycolate is so much lower than that for the calcium complex; a similar relationship is not apparent from a comparison of the citrate complexes, nor, for that matter, other complexes of carboxylic acids, cc-hydroxy- or oc-aminocarboxylic acids (e.g. iminodiacetic acid). The efficacy of citric acid, diglycolic acid, and NTA as chelating agents to be added to detergents to sequester metal ions may be compared by calculating the concentration of free metal for a selected set of values of total metal and ligand concentrations and of PH. An illustrative example is given in Table IV based, unless
METAL TABLE
CITRATE
AND
DIGLYCOLATE
IV
PERCENTAGE
OF
(c,=c,_=o.OOl
M)
METAL
UNCOMPLEXED
M/L
Citric uciti
Dit~lycolic
Calcium Mupncsium Zinc Nickel Coppcr( II) lron(I1) Iron(lIl)
50.7 55.4 10.8 7.78 0.6gh 14.9h 0.00023”
41.4 89.0 40.3 76.8’
u Vnluc h Vnlucs c Vuluc ” Value * Value
105
COMPLEXES
crllculntcd cnlculntcd calculntcd cnlculatcd culculatcd
by constants by constants by constants by constants by constants
other other other other other
thnn than than than than
those those those those those
AT pH G
trcitl
N/tr//otriacet/c
ctcid
77.3 9G.5 I .20 0.45 0.06Sc 9.52” 0.00 I 9” in Tnblcs in Tables in Tables in Tnblcs in Tnblcs
l-111: l-111: I-III; I-111: I-111:
see ref. 12. see rcl’. 13. see ref. 14. .scc ref. 15. .scc ref. 16.
noted otherwise, on the constants of Tables I-III. The results show that while citrate cannot match the performance of NTA for sequestering transition metal ions. it may be superior to the latter for the ions of the Group IIA metals. As an agent for the reduction of free metal ion concentration, a mixture of citrate and NTA might prove more effkaceous than either substance alone. SUMMARY
Metal complexes formed in solution by citric acid or diglycolic acid with calcium. magnesium, zinc or nickel ions have been studied by pH titrations. Normal and protonated complexes of 1: 1 stoichiometry are formed, the compositions and stability constants of which are reported which are in best agreement with the experimental data. The effectiveness of these ligands for the sequestration of metal ions is compared with that of NTA.
REFERENCES
I J. Bjcrrum. hfetcd At~ttl?~c Fornrution in A~IJWNJS Sohct/orr. P. Haasc und Son, Copenhagen. I94 I. 2 1. G. Saycc. Takrrm. I5 ( 1968) 1397. 3 D. D. Pcrrin and I. G. Suycc. Chen~. Id. (Lo~rrlon). ( 1966) 661. 4 G. Schwrlrzcnbnch nnd H. Flaschkn. Ccwpkwnwtric Titrtrtions. Mcthucn, London. 5th Gcrmnn cdn. trnnslutcd by H. M. Irving. 1969. 5 W. A. E. McBrydc. Analyst (Lonclu~~). 94 ( 19G9) 337: 96 (197 I) 739. 6 R. G. Bntcs and G. D. Pinching. J. AVIW. C/~ertt. Sot.. 71 (1949) 1274. 7 E. Campi. G. Ostxoli. M. Mcironc and G. Suini. J. Ilrc>ry. NW/. C/rem., 26 (1964) 553. 8 N. C. Li. A. Lindcnbaum and J. M. White. J. iwrg. N~rc/. C/renr.. I2 (1959) 122. 9 K. N. Pcurcc. Ph.D. Thesis. Massey University. New Zcolund. 1972. IO D. D. Pcrrin. SIJOUI. Kef~tf.~Me/tt/ A, 42 (19G9) 205. I 1 R. M. Tichanc und W. E. Bcnnctt. J. Amer. Clretn. Sot.. 79 (1957) 1273. 12 M. Ynsudu. K. Yumusaki and H. Ohtnki, BIJ//. Chwn. Sue. JN~.. 33 (1960) 1067.
106 13 14 15 16
T. W. G. E.
T. B. FIELD.
J. COBtiRN.
J. L. MkOURT.
W. A. E. McBRYDE
B. Field. J. L. McCourt and W. A. E. McBrydc. Curl. J. Ckrn.. 52 ( 1974) 3 1 19. A. E. McBrydc. J. L. McCourt nnd V. Chcam. .I. Irrcir(/. Nd. Chtrr.. 35 ( 1973) 4193. Schwarzcnbach. G. Andcrcgg. W. Schncidor and H. Scnn. ffeh:. Chirn. Arm. 38 (1955) Bottnri und G. Andcrcgg. Hch. Clriru. Acfcr. 50 (1967) 2349.
1147.
COMPOSITION DIGLYCOLATE TERRENCE
Amsterdam
AND STABILITY OF SOME COMPLEXES IN AQUEOUS
B, FIELD.
Depurtowu
of Clwmisrry.
(Rcccivcd
3rd June 1974)
JOHN Uhwrsity
COBURN.
JANET
r>J Worerloo.
101
- Printed in The Ncthcrlnnds
L.
Waterloo.
METAL CITRATE SOLUTION MCCOURT
nnd W.
AND
A. E.
McBRYDE
Ontcwlo N.?L 3GI (Cancrtltr)
There has been some renewed interest in the metal complexes formed by citrate ion in aqueous solution, arising from the consideration of citric acid as a potential detergent builder. Diglycolic acid has also been considered for the same purpose. and corresponding data for the stability of its complexes would be useful. Of these ligands. citrate at least is known to form protonated species MHL or MH2L where HJL represents citric acid. When complex formation is measured by proton displacement in the manner first advocated by Bjey,rum’. interpretation of the experimental results is made more difficult when an unknown amount of the hydrogen-ion balance is incorporated in such protonated complexes. The authors have had access to the powerful computer program SCOGS2 which utilizes pH-titration data for systems consisting of metal, acid and ligand, and fits stability constants for proposed complex species to the data by a non-linear least-squaresmethod. The special advantage of this computer method for the systems to be studied is that it can handle protonated or deprotonated complex species, hydrolysis products, and so forth. Accordingly suitable pH titrations of solutions containing citric or diglycolic acid have been done with and without metal present, and with an added indifferent electrolyte (KNOJ) added to try to achieve constant ionic strength. The results of these titrations were then interpreted in terms of possible complex species, those complexes and their formation constants being preferred and reported which gave minimum sum of squares of deviations between calculated and experimental titres. EXPERlMENTAL
The pH titrations were carried out in a vessel thermostatted at 25.O”C (ref. 3) with standard alkali or adid added by a calibrated syringe with micrometer drive, The solutions of metal nitrates were prepared from analytical reagent salts and standardized by EDTA titrations 4. Molar alkali solution for titration was prepared bv dilution of ampoules of B.D.H. concentrated volumetric solution, standardized with acid phthalate, and stored under COz-free nitrogen. Recrystallized potassium nitrate was added to all solutions as background electrolyte at a molarity of 0.1. The citric acid (B.D.H.) and diglycolic acid (J. T. Baker) were of reagent grade, and were used without further purification when their suff?cient purity had been established by titration. The pH measurements were made either with an Orion Model 801 meter
T. B. FIELD.
102
J. COBURN.
.I. L. MCCOURT.
W. A. E. McBRYDE
equipped with Beckman electrodes or a Radiometer Model PHM52 meter with Radiometer electrodes. The pH scale WLIS established for these two assemblies with N.B.S. buffer solutions. and any observed departure from the Nernst slope for readings between pairs of buffers was compensated by means of the tcmperaturc-adjustment control on each meter. All pH-meter readings were converted to hydrogen-ion concentrations by means of nn empirically determined factor f = 10-F”/[H]. measured under the sdme conditions and with the same electrodes as applied lo particular titrations5. During titrations the apparatus was continuously purged with CO,-free nitrogen. In the absence of metal ion, the pH titrations yield acidity constunts for the ligands. and when metal ions are prcscnt the titration curve is modified according to the amount of hydrogen ion set free by complex formation. In tither case the observations yield stoichiometric equilibrium constants valid within the particular salt background used. It may be noted in connection with acidity constants that these are here given as protonation constants, and numbered in order of increasing numbers of protons bound. This practice places these constants in a formally analogous relationship with the numbering of complex stability constants. and aids in the writing of generalized statements in computer programming. It may bc mentioned that the titrations to determine metal complex stability constants were carried out with an excess of metal over the ligand present in order to secure a suflicient degree of conversion to complexes. In the inverse situation, the titration curve of citric acid is not sufliciently displaced by the presence of a smaller amount of metal to give a precise indication of the extent of complex formation. It follows that under the experimentul conditions used in this study, no evidence of complexes with higher than 1 :l stoichiometry was detected. RESULTS
AND
DISCUSSION
The results obtained from this study are summarized in Tables I-III. They are reported as averages from the data of several titrations in each case, together with a standard deviation on each reported constant. In the various titrations from which results are combined the concentratidns of metal (if present) :tnd TABLE ACIDITY
1 CONSTANTS
(Conwntr;!tion
quolicnts.
FOR
CITRIC
2S’C.
0.1
M KNO,
IO{\ /II’ -._--
----Citric acid Diglycolic acid
5.70 * 0.02 3.95_to.o1
AND
DIGLYCOLIC
ACIDS
background)
IOWA 11:’ --_._-___--~.-
log /I:’
---.
10.06i- 0.02 6.73 & 0.0
12.87 + 0.08
I
ligand were systematically varied. nnd there was no evidence that any of the /1 vuiues derived from the experimental results showed any dependence on the overall concentration of the reagents. Typically, the levels of metal or ligand concentration were in the range 1-G. 10WJ M.
METAL
CITRATE
TABLE
II
STARILITY
AND
DIGLYCOLATE
CONSTANTS
FOR
25’C. (Conccntrulion quotients. --_._..-_.-.-.---.--. -_----..--.-..
0.
METAL ---..---..--.-
103
CITRATES
I M KNO.1
IV1i-I 1, h /I I I I -._-_ . ..__._- ..- -._-_.-_ --_..- .-.--.--.8*02_eO. I5 Culciuni 7.66&O. I9 Magnesium X.08 _e 0.07 Zinc 9.04-e 0.07 Nickel __ ____ _. __-_-._-__-.-._-_-_ ._..-..
TABLE
COMPLEXES
bilckgrottnd) - ------.--.-._..___
---.-
.-_._-
_... _..._-_..-- ._.-._.
MI, hl II I 0 I -.-.--------. -- - _ _- ..- - ..-_. ..-- -...._-._-_-._,.SO&-0.06 3.38 f. 0.07 5. IO-r_ 0.05 5.40 + 0.04 _. ._..____._.__.-.._ .__.._...__ -_._ ._.... ..-_..-. -.
.-._-
_... .-_
_ __._._._-_.
___._. ._--.
.-_-_.- .._.-- .._..- ..._..-..
III
STABILITY (Concentration ___ ..__._._-_.--
CONSTANTS
FOR
METAL
DIGLYCOLATES
25°C. 0. I M KNO., bwkground) quolicnts. .._-. .-- .___._..-.._- ._..._._ _-._ ._._..__ -__..._____.__._... -. -.-- ...____- _._......_. --.--
_ .___ -._---.-.._-_.-_----
MIIL ht /{I I L
ML Iv/ P I I,I ..__-._- ._-._-.-.._.-.-__-..-.-
6.56 + 0.04 Calcium 538 _e 0.02 M:lgncsium 6.03 +- 0.05 Zinc ____ _._. _--___-.-.._ .._ ---.-
X54+0.02 2.15~0.01 3.59*0.01 _..... - ___._-- __.._._---.__
_....-..
-... _.._....... -.. __--
._._- _.__. -. _.- .__.-..________
.__-_--
._.. -.._ ._.__..-.
___.
__...-._-._ -_-.---
___ __ __.__ ._... - _... -._.-
The values of the equilibrium constants arc those which give the minimum standard dcviution in litre (s.d.t.) for the titrations from which they arc derived. They also refer to the choice of species. from various possibilities tried out. which give the lowest s.d.t. Thus. for instance. in the case of diglycolic acid the complexes ML and MHL taken together gave in every titration or combination of titrations ;I lower s.d.t. than that calculated when ML wys the only species assumed to be prcscnt. This is the evidence on which our assumption of the formation of protonatcd diglycolate species is bused. The acidity constnnts listed in Tuble I are concentration quotients ilpplicablc only in solutions of comparable ionic strc-ngth. Owing to the multiple charge borne by some of the ions gcneratcd by thcsc acids. and the effect of this on activity cocflicicnts. the stoichiomctric ionization constants. pi~rtiCUlilrly of citric acid. show ii marked dependence on the ionic strength (c:/‘. thermodynamic values”). The values reported here appear to be in sittisfuctory accord with those reported under nearly comparable conditions by Campi et al.’ and by Li e/ trl.“. A recent summary of the literature on citrate equilibria by Pearce” including his own measurements lists five. sets of stoichiometric acidity constants for ionic strength 0.1 und at specified tcmpcrntures ranging from 20 to 30°C. The present results for pK I are slightly lower. but the others are in excellent agrecmcnt with this summary. Pearce has applied a correction to the values of pKJ for complexation of citrate ion by potassium or sodium ion: this refinement has not been introduced in the present work. With respect to the citrate complexes, it is clear that protonated species do occur, certainly. MI-IL ilnd possibly MH2L. The titration data failed to product
104
T. B. FIELD.
J. COBURN.
J. L. MCCOURT.
W. A. E. McBRYDE
evidence to support the existence of the latter, but Campi et ~1.’ and Pearce9 have reported stability constants for this with certain metals. The pH-titration method for studying metal complex formation is much less capable of defining the state of equilibrium when protonated species occur. This can be seen from the relationship GI-[HI + [OH] = ~qljpqr[M]PIH]rl[L]r+CjljlOIH]JIL] in which Cu is the total net concentration of ionizable hydrogen at any point. If q=O throughout, [L] is explicitly known, given the /$’ values. If y#O, the above relationship does not define [L], and one must resort to best-fit solutions of several mass-balance equations to define [M] and [L]. Experience has shown that in such cases various sets of p. y, and I’ can be found which satisfy the data roughly as well, so that a clear choice of constituent species cannot be made. It is a further limitation on any method of determining equilibrium constants that there be a sufficient and thus precisely delined concentration of each species involved. In the case of the species MH2L, if it exists, the pK for its dissociation to MHL may be so small that to obtain a reasonably amount of this complex would require working at a pH so low as to preclude much formation of any metal complex. With the aid of Pearce’s formation constants and the computer program SPECON”, the concentration of each species formed between pH 2 and pH 6 was calculated for a number of selected values of Chl and C,. For the case where CL= I .O * IO- 3 and C, = 5.0*10-’ (corresponding to one of our titrations), the maximum value of [MH,L] was 3.8. 10 -5 M (at pH 3.5). This is not enough conversion of either ligand or metal to yield a precise definition of /3izl. Even when C,., was raised to 1.0. 10s2, only 7”/, of the ligand was converted to [MH,L] at its maximum concentration. The values of log fi rol for calcium and magnesium complexes are in good agreement with those of Campi et al.’ and Pearceg, though slightly higher than those of Li at ~1.~. The corresponding values for nickel and zinc accord with those reported by Campi and coworkers. The present values for log /jr 1 1 for these complexes tend throughout to be slightly higher than Campi’s, and *rather closer to Pearce’s for the first two metals. The discrepancy may be linked to the inability to characterize MHzL as a discrete species: if some of this was formed in the present experiments, the effect on the data would be to augment the estimate of [MHL] and hence of /I,, 1. Values of stability constants of the diglycolate complexes may be compared with those previously published by Tichane and Bennett’ ’ and by Yasuda et ~1.‘~. Apart from the indication of protonated species at low pH and the proposed stability constants for these, the results appear to be in general agreement with earlier work. It is not clear why the value of fl io1 for magnesium digiycolate is so much lower than that for the calcium complex; a similar relationship is not apparent from a comparison of the citrate complexes, nor, for that matter, other complexes of carboxylic acids, cc-hydroxy- or oc-aminocarboxylic acids (e.g. iminodiacetic acid). The efficacy of citric acid, diglycolic acid, and NTA as chelating agents to be added to detergents to sequester metal ions may be compared by calculating the concentration of free metal for a selected set of values of total metal and ligand concentrations and of PH. An illustrative example is given in Table IV based, unless
METAL TABLE
CITRATE
AND
DIGLYCOLATE
IV
PERCENTAGE
OF
(c,=c,_=o.OOl
M)
METAL
UNCOMPLEXED
M/L
Citric uciti
Dit~lycolic
Calcium Mupncsium Zinc Nickel Coppcr( II) lron(I1) Iron(lIl)
50.7 55.4 10.8 7.78 0.6gh 14.9h 0.00023”
41.4 89.0 40.3 76.8’
u Vnluc h Vnlucs c Vuluc ” Value * Value
105
COMPLEXES
crllculntcd cnlculntcd calculntcd cnlculatcd culculatcd
by constants by constants by constants by constants by constants
other other other other other
thnn than than than than
those those those those those
AT pH G
trcitl
N/tr//otriacet/c
ctcid
77.3 9G.5 I .20 0.45 0.06Sc 9.52” 0.00 I 9” in Tnblcs in Tables in Tables in Tnblcs in Tnblcs
l-111: l-111: I-III; I-111: I-111:
see ref. 12. see rcl’. 13. see ref. 14. .scc ref. 15. .scc ref. 16.
noted otherwise, on the constants of Tables I-III. The results show that while citrate cannot match the performance of NTA for sequestering transition metal ions. it may be superior to the latter for the ions of the Group IIA metals. As an agent for the reduction of free metal ion concentration, a mixture of citrate and NTA might prove more effkaceous than either substance alone. SUMMARY
Metal complexes formed in solution by citric acid or diglycolic acid with calcium. magnesium, zinc or nickel ions have been studied by pH titrations. Normal and protonated complexes of 1: 1 stoichiometry are formed, the compositions and stability constants of which are reported which are in best agreement with the experimental data. The effectiveness of these ligands for the sequestration of metal ions is compared with that of NTA.
REFERENCES
I J. Bjcrrum. hfetcd At~ttl?~c Fornrution in A~IJWNJS Sohct/orr. P. Haasc und Son, Copenhagen. I94 I. 2 1. G. Saycc. Takrrm. I5 ( 1968) 1397. 3 D. D. Pcrrin and I. G. Suycc. Chen~. Id. (Lo~rrlon). ( 1966) 661. 4 G. Schwrlrzcnbnch nnd H. Flaschkn. Ccwpkwnwtric Titrtrtions. Mcthucn, London. 5th Gcrmnn cdn. trnnslutcd by H. M. Irving. 1969. 5 W. A. E. McBrydc. Analyst (Lonclu~~). 94 ( 19G9) 337: 96 (197 I) 739. 6 R. G. Bntcs and G. D. Pinching. J. AVIW. C/~ertt. Sot.. 71 (1949) 1274. 7 E. Campi. G. Ostxoli. M. Mcironc and G. Suini. J. Ilrc>ry. NW/. C/rem., 26 (1964) 553. 8 N. C. Li. A. Lindcnbaum and J. M. White. J. iwrg. N~rc/. C/renr.. I2 (1959) 122. 9 K. N. Pcurcc. Ph.D. Thesis. Massey University. New Zcolund. 1972. IO D. D. Pcrrin. SIJOUI. Kef~tf.~Me/tt/ A, 42 (19G9) 205. I 1 R. M. Tichanc und W. E. Bcnnctt. J. Amer. Clretn. Sot.. 79 (1957) 1273. 12 M. Ynsudu. K. Yumusaki and H. Ohtnki, BIJ//. Chwn. Sue. JN~.. 33 (1960) 1067.
106 13 14 15 16
T. W. G. E.
T. B. FIELD.
J. COBtiRN.
J. L. MkOURT.
W. A. E. McBRYDE
B. Field. J. L. McCourt and W. A. E. McBrydc. Curl. J. Ckrn.. 52 ( 1974) 3 1 19. A. E. McBrydc. J. L. McCourt nnd V. Chcam. .I. Irrcir(/. Nd. Chtrr.. 35 ( 1973) 4193. Schwarzcnbach. G. Andcrcgg. W. Schncidor and H. Scnn. ffeh:. Chirn. Arm. 38 (1955) Bottnri und G. Andcrcgg. Hch. Clriru. Acfcr. 50 (1967) 2349.
1147.