Y. TOMITA, T. ANDO,AND K. UENO
404
Infrared Spectra of Nitrilotriacetate Chelates in Aqueous Solution
by Yuko Tomita, Takesbi Ando, and Keihei Ueno Contridutwn No. 66 from the Departmat of Organic Synthesis, Faculty of Engineering, Kyushu University, Fukuoka, Japan (Received June 8,1964)
The nitrilotriacetic acid chelates of magnesium and copper in aqueous solution have been studied at various pH values by infrared spectroscopy using deuterium oxide m the solvent. On the basis of the antisymmetric stretching bands of the ionized, un-ionized, and coordinated carboxyl groups, the schemes of the chelate formation reactions are proposed. In the case of the magnesium chelate, the stability constant has been estimated from the effect of pH upon the spectra.
Introduction Recently, the applications of infrared spectroscopy to the study of metal chelates in aqueous solution have provided new information.l-a In our previous paper,4 the chelating behavior of nitrilotriacetic acid (NTA) was investigated by infrared spectroscopy in the solid state as well as in solution, and the antisymmetric carbonyl stretching frequencies of NTA in aqueous solution were utilized to determine the state of the carboxyl group. In this investigation, the infrared spectra of NTAmetal chelates in aqueous solution were taken as a function of the pH to supplement the information obtained by potentiometric measurements. The metal chelates chosen in this study are magnesium- and copper(I1)-NTA, because they are suiliciently soluble in deuterium oxide for the infrared spectral measurement. On the bases of the spectral changes with pH and of our previous knowledge of the metalNTA chelates, the coordinating structures of NTA with these metal ions in solution and a new method of determining their chelate stability constant are proposed.
Experimental Infrared Spectra. The infrared spectra in aqueous
metal chloride in deuterium oxide (99.5%) a t concentrations approximately 5-100/, by weight. The i n o r g a b salts used in preparfng the complexes were rettgent grade in all casea The free acid of NTA (donated by Dojindo and Co.) was analytical grade and was used without further purification. The pH values of the solutions were controlled by the careful addition of concentrated NaOD solution from a microburet.
Table I: Antisymmetric Carboxyl Absorption Bands (cm.-1) Mg-NTA complex
PH
3.2 4.2 5.5-10.0 11.6
COOH
1730
N H +CHxCOO-
1625 1625 1625
COO-Mg
Predominant species
I
I1 1610 1610
11, v
v
Cu-NTA complex
PH
COOH
NH+CHnCOO-
1.6 1 .&11 .o
1730
1625
COO-Cu
Predominant species
I 1615
v
-
solution were observed using a cell with calc3um fluoride The measurements were made with a Nippon Koken Model D8 301 double-beam spectroDhotbmeter equipped - - - with sodium chloride optics. - Preparation of Metal Chelates and pH M k r e m e n i s . Of were by dissolving equimolar amounts of NTA free acid and
The Journal of Physical c h t k t r y
(1) (a) K. Nakamoto, Y. Morimoto, and A. E. Martell, J. Am. C h a .
SOC.,84, 2081 (1962); (b) ibid., 85, 309 (1963).
J. E. Tackett, ibid., 85, 314,2390 (1963). (3) D. Chapman, D.R. Llond, and R. H. Prince, J . Chem. S O ~3645 ., (1963). (4) Y . Tomita and K. Ueno, BuU. C h m . Soc. Japan, 36, 1069 (1963). (2) D. T. Sawyer and
405
INFRARED SPECTRA OF KITRILOTRIACETATE CHELATES
The pH values listed were obtained from the solutions of the duplicate run using ordinary water and were measured with a Hitachi-Horiba ?Model P pH meter. Therefore, the pD values of the deuterium oxide solution used for spectral measurements may be slightly different from the conventional pD values, which are obtained from the equation, pD = meter reading 0.40, but they are within the tolerances required for meaningful interpretation of the results.
z
+
0
F a a:
0
Results and Discussion
v)
When NTA is mixed with a metal ion in aqueous solution, the extent of their interaction is governed mainly by the solution pH and the kind of metal ion. The possible ionic species existing in the solution are schematically shown in Figure 1, and the antisymmetric carboxyl stretching frequencies of NTA-metal chelates are listed in Table I. It has been known that the antisymmetric stretching band of the carboxyl group of NTA in aqueous solution occurs in different frequency regions according to the state of the carboxyl group, and the band due to >N-CH,COOH, >NH+-CH,COO-, or >N-CH2COO- appears a t 1730-1700, 1630-1620, or 15851575 cm. -l, respectively. la The similar investigation on the metal chelates of NTA has revealed that the coordinated carboxyl group gives rise to a band a t 1615-1605 ~ m . - l . ~ This information, together with the knowledge of the potentiometric t i t r a t i ~ n allows ,~ one to follow the scheme of reaction which occurs during the chelate formation. Magnesium-NTA Chelate. The carboxylate absorption band of magnesium-NTA chelates is shown as a function of solution pH in Figure 2. At pH 3.2 tmo peaks a t 1730 and 1625 cm.-l are observed. As the pH is increased, the weak band a t 1730 cm.-l dis-
'JI
1
I
m
a
I
' bI
L
1700
I
I
1600
FREQUENCY, C M. - I Figure 2. Infrared spectra of Mg-NTA complex in DtO
- - - -,pH 3.2; ---,pH 4.2; - - - - -,pH 5.5; - - -, pH 10.0; ---- pH 11.6.
solutions: pH 6.8;
-,
*
*
appears, and a single strong band a t 1625 cm.-1 is observed. With the further increase of pH, a new band appears a t 1610 cm.-l which grows in intensity a t the expense of the 1625-cm.-' peak. Finally, a t the most basic pH value, only the peak a t 1610 cm.-l remains. According to Nakamoto and others, two bands a t 1720 cm.-l (vw) and 1623 cm.-l (m) which appeared in a solution of pH 2-3, were assigned to the COOH and NHf-CH2COO- groups, respectively, of the monovalent NTA anion (H2A-).Ia As the spectrum of magnesium-NTA a t pH 3.2 is quite similar to that observed by Nakamoto for the free ligand in the low pH region, it is apparent that there occurs little interaction between the ligand and magnesium ion at this pH value (I in Figure 1). With the increase of pH, NTA is known to dissociate to give a divalent anion (HA2-), which shows only one band a t 1625 cm.-l. Since the change of spectrum of magnesium-NTA chelate from pH 3.2 to 4.2 is identical with that of free ligand in the same pH range, no coordination of NTA with magnesium ion is indicated in this pH region, and the predominant species existing
YII
Figure 1. The solution equilibria of NTA-metal complex.
(5) G . Sohwarzenbaoh, E. Kampitsoh, and R. Stener, Helv. China. Acta, 28, 828 (1945).
Volume 69, Number 2
February 1966
Y. TOMITA, T. ANDO,AND K. UENO
406
in the solution are the divalent ligand anion and free metal ion (I1 in Figure 1). This conclusion is also supported by the titration data of magnesium-NTA chelate since no appreciable pH drop is observed in this pH region. When the solution pH becomes higher than 4.2, a new band at 1610 cm.-l is observed. This band, which is not observed in the free ligand over the entire pH range investigated but is found in the NTA-metal chelates, has been assigned to a coordinated carboxyl group.4 Since the two bands a t 1610 and 1625 cm.-l are observed in the solution of pH 4.2 or higher, the ionic species existing in this pH range may be an equilibrium mixture of the divalent ligand anion, free metal ion, and the tetracoordinated chelate6 (I1and V, respectively, in Figure l), and the last component (tetracoordinated magnesium-NTA chelate) exists as a single ionic species only at the relatively high pH region, where a single band at 1610 cm.-l is observed. The increase in absorption a t 1610 cm.-l and the decrease in absorption at 1625 cm.-l with increasing pH indicate that the intensities of these peaks may be proportional to the concentration of ionic species I1 and V, respectively. Although the peak at 1625 cm.-1 could be associated with a complex containing an acidic hydrogen, such as IV, the pH titration study appears to rule out any appreciable amount of such a complex above pH 4.' As is known from the titration data, the chelate formation between NTA and magnesium ion occurs in the pH region higher than 4.8.5 Thus, the main equilibria involved in the pH region of 4.2 or higher are HA2-
+ M2+ = MA- + H+
A3-
+ M2+ = MA-
(1) (2)
However, eq. 2 will be valid only in the relatively high pH region, as the third dissociation constant of NTA is 10-9.73. Therefore, the main equilibrium, which is effective at pH 4-8, will be represented by eq. 1,and the equilibrium constant can be given as
where KMAand Ka are the stability constants of magnesium-NTA chelate and the third acid dissociation constant of NTA, respectively. As previously noted, it is reasonable to assume that the intensity of the 1610- and 1625-cm.-l peaks are proportional to the concentrations of ionic species MA- and HAZ-, respectively. Then, the relative concentrations of MA- and HA2- can be derived according to the usual treatment in the spectrophotoThe J O U T Wof~ Physical Chemistry
metric determination of two components mixtures. As the spectra a t pH 3.2 and 11.6 are believed to represent the divalent ligand anion (HA2-) and the tetracoordinated magnesium-NTA chelate, respectively, molecular extinction coefficients for both components can be obtained from their spectra. Thus, the concentrations of both components in a solution of given pH value may be calculated from the peak intensities of 1610- and 1625-cm.-l bands in the spectra of intermediate pH value. Plots of such concentrations in terms of the absorption intensity against pH are shown in Figure 3.
I
p' v'
3
4
I
5
7
6
8
9
1 0 1 1
PD
Figure 3.
Absorbancy of carboxylate bands of
Mg-NTA complex as function of pH:
-, 1625-cm.-l
band;
- - - -, 1610-cm.-l
band.
At the pH values where the intensity of the peak is one-half of the original intensity, the concentrations of HA2- and MA- become equ'al, as the spectral measurements were made on the equimolar mixture of metal ion and the ligand. Then, eq. 3 can be rewritten as log K ~ =ApK3
- pH - log (M2+)
(4)
The pH value which satisfies this condition is determined from Figure 3 as 5.65 (1610-cm.-l peak) or 5.9 (1625-cm.-' peak). Although the agreement of values obtained from the two different peaks is not completely satisfactory, the mean pH value 5.75 is used for the calculation of the stability constant described as follows. To determine the stability constant, KYA, the value of pK1 must be corrected for D 2 0 solutions. (6) Although there is no positive evidence for the presence of a N-Mg bond in the Mg-NTA chelate, a recent n.m.r. study proved the existence of nitrogen-metal bonds in the EDTA chelates of alkaline earth metals [R. J. K d a , D. T. Sawyer, S. I. Chan, and C. M. Finly, J. Am. C h m . SOC.,85, 2930 (1963)l. Therefore, it is likely that the Mg-NTA chelate exists as V rather than as Va in Figure 1. (7) G. Schwarzenbach and E. Freitag, Helv. Chim. Acta, 34, 1492 (1951).
INFRARED SPECTRA OF NITRILOTRIACETATE CHELATES
the 1625-cm.-l peak can be observed only as a very weak shoulder. With further increase of pH, the new becomes a well-defined, single band at 1615 band which does not change up to pH 11.0. The two bands at 1730 and 1625 cm.-l, which are observed at pH 3.2 in the spectrum of the magnesiumNTA chelate, can be assigned to COOH and NH+CH&OO-, respectively; both are characteristic of the free ligand in the low pH region. Thus, it is clear that NTA does not interact with copper(I1) ion at pH 1.6. However, in the solution of pH 1.8 or higher, copper(11)-NTA shows only one peak at 1615 em.-', which is characteristic of the coordinated carboxyl group. As no other band is observed in this pH region, it is fair to conclude that the predominant species existing in the solution is a tetradentate normal chelate (V in Figure 1). Thus, in the case of copper, the chelate formation reaction represented by the equation
I
1
2 0
F
a a
0 M
m
a
.1 L
I
I700
I
I
I
HA-
1600
FREQUENCY, CM;' Figure 4. Infrared spectra of Cu-NTA complex in DzO solutions:
407
- - - -, pH 1.6; -, pH 1.8; - - - - -,pH 2.3.
This can be done using the equations proposed by Li, Tang, and Mathur,s to give a calculated value for pKa of 10.13. The free metal ion concentration (M2+) in this condition is equal to 0 . 5 C ~where , CM is the total concentration of metal ion and log (M2+)is calculated as -0.86. Then, one could calculate the stability constant K M A from eq. 4. Although the calculated value log K M A= 5.24 does not agree very well with the literature values (log = 5.41): it is indicative of the validity of determining chelate stability constants from infrared absorption measurements. Copper-N T A Chelate. Infrared spectra of copper(11)-NTA chelate are shown in Figure 4 as a function of solution pH. At pH 1.6, two peaks a t 1730 and 1625 em.-' are observed. However, when the pH is increased to 1.8, a new band appears a t 1615 cm.-l, the 1730-cm.-1 peak has completely disappeared, and
+ M2+ = MA- + 2H+
(5)
occurs instantaneously in a fairly low pH region. Since no intermediate stage of chelate formation is observed for the copper(I1)-NTA chelate, the stability constant cannot be derived in this case. Finally, it is interesting to note that the copper(I1)-NTA chelate shows only one carboxyl band at 1615 em.-'. Coordination to the metal of the nitrogen atom and three carboxylate groups cannot result in a planar NCuo configuration. The absence of a free COO0 0 band in the 1585-1575-~m.-~region of the spectrum suggests that the nitrogen atom and three carboxylate groups occupy four octahedral coordination sites around the copper atom in aqueous solution.
Acknowledgment. The authors are grateful to Dr. K. Nakamoto of the Illinois Institute of Technology, who gave valuable suggestions at the early stage of this investigation. They also thank the Ministry of Education, Japanese Government, for the financial support for this work. (8) N. C. Li, (1961).
P. Tang, and R. Mathur, J. Phys. Chem., 65, 1074
Volume 69,Number 8 February 1966
