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Chelation of cobalt, nickel and copper with glutamic acid
ARCHIVES
OF
BIOCHEMISTRY
AND
86, 82-88 (1959)
BIOPHYSICS
Chelation of Cobalt, Nickel and Copper with Glutamic Acid Sister M. Helen Therese Nyberg’ From the Chemistry
Department,
Malloy Catholic College for Women, Rockville Long Island, New York
Centre,
Michael Cefola From
the Department
of Chemistry,
Fordham
University,
New
York,
New
York
and
David Sabine From the United
States Vitamin Yonkers,
Corporation,
Received
March
and Pharmaceutical New York 12, 1959
The present investigation is intended to supply further evidence for proposed theories regarding some general trends in the stability of metal chelates. Cobalt, nickel, and copper were examined for relative stabilities with glutamic acid. Amino acids, in general, are said to combine with the more basic metal ions as the alkaline earths, and to have an appreciable affinity for the transition metals (1). Lumb and Martell (2) found stability constants for the chelates of the alkaline-earth metals with glutamic acid. It will be interesting to compare them with those of the transition-element complexes. EXPERIMENTAL Apparatus,
Materials
and Method
The first and second ionization constants of glutamic acid were obtained by potentiometric titration at 30°C. For the determination of the metal chelate constant, a 1: 1 ratio of metal to acid was used. Solutions were made up to such concentration as to maintain ionic strength of 0.1 throughout the titration. Potassium chloride was the electrolyte used. Experimental solutions were made to contain a final acid and metal concentration of 0.01 M. 1 This thesis was presented by Sister M. Helen Therese Nyberg, 0. P., in partial fulfillment of the requirements for the degree of Master of Science at Fordham University, 1957. 82
CHELATION
WITH
GLUTAMIC
ACID
83
The nL-glutamic acid was purchased from Mann Research Laboratories, New York, and was used without further purification. The standardization of the acid was accomplished potentiometrically, using a differential curve to determine the end point. The metal-chloride solutions were prepared from Fisher Certified reagents, and were standardized with silver nitrate, using dichlorofluorescein as indicator. The carbon dioxide-free KOH solution was prepared according to the method of Schwarzenbach and Biedermann (3). It was standardized for basic strength using potassium acid phthalate with phenolphthalein as indicator, and was also standardized for chloride ion. The titration was carried out in nitrogen atmosphere using a 250.ml. flask with five necks to accommodate the tube delivering nitrogen, a mercury-seal stirrer, a Beckman # 1190-80 glass eleci:-.: e, a Beckman K calomel electrode, and a buret. Small increments of base (0.1 M) were added until a total of 20 ml. was reached. pH readings were made on a Beckman Model G pH meter.
Calculations Figure 1 shows curves obtained for the titration of glutamic acid in the absence of and in the presence of the metal ions. The dissociation constants of the acid (k, and Icz), as well as the first stability constants (KI) of the chelates, were calculated according to the algebraic method described by Chaberek and Martell (1). The Bjerrum
PH
1
1
1.0
1.5
a ( Moles
bose /mole
acid 1
FIG. 1. Titration curves for glutamic acid and for the cobalt (II), nickel (II), and copper (II) chelates of glutamic acid at 30°C. The mark ( on the curves indicates the pH at which the hydroxide precipitation begins. Points on the curve beyond these pH values are not very reliable.
84
NYBERG, CEFOLA, AND SABINE
method of plotting fi (number of ligands bound per metal ion) vs. pA (negative log of the A= concentration) was employed also in the case of the chelate constants. At rt = 0.5, pA = log K1 (4). Activities of H+ ions were obtained directly from pH measurements. The OH- ion activity was calculated from a knowledge of the fact that the ionization constant of water at 1 atm. pressure and 30°C. is 1.471 X lo-l4 (5). In obtaining actual concentrations from activities, it was assumed that y&i= yfH+ = yfoH(6). A large-scale graph for molality vs. mean activity coefficient (yf) of the chloride ion at 30°C. was made from data given by Harned and Cook (7), and extrapolated to zero concentration. From this graph, activity coefficients were read for the chloride ion and, hence, for the hydrogen and hydroxide ions also. Hydroxide formation occurred in all the metal chelate titrations. Chaberek et al. (8) made a study of this phenomenon and gave a summary of their results. Because of this hydrolysis interference at higher pH values, K1 was calculated algebraically from points along the upper buffer region below the pH where the hydroxide begins to form. Values checked very well with the results obtained by the Bjerrum method. DISCUSSION OF RESULTS
Results are represented in graphic form in Fig. 1. The curve for the free acid indicates a two-step ionization. The reaction of the metal ion with the acid is accompanied by the displacement of protons from the acid with the consequent lowering of pH. Therefore, the curves which represent the chelates have an upper buffer region lower than that of the free acid, the degree of lowering being proportional to the stability of the chelate formed. In the case of the cupric and nickelous ions, even the first buffer region is lowered. For these metals we have: M++ + H2A ti MA + 2H+. Relative positions of the metal curves indicate relative stabilities to be Cu > Ni > Co. In the case of copper and nickel, we also have an inflection point at a = 2, which indicates a rather strong 1: 1 complex formation. Table I gives values obtained for the above constants. These values are in reasonable agreement with those already reported in the literature (15), although the temperature and ionic strength at which these measurements were made are somewhat different. It is interesting to note that the pi& of glutamic acid is approximately an average of the different values listed by Lumb and Martell (2), if those which were corrected to zero ionic strength are omitted. In the case of pkz ,
CHELATION
WITH
GLUTAMIC
TABLE Chelate Stability Glutamic
acid:
I
Constants
at SO”C.
pkt = 4.23; p/c%= 9.46. p = 0.1 log
Metal
ion Algebraic
co++
TABLE
Metal
KI
method
Bjerrum
4.48 5.60 7.78
hTiTii-* cu++
Chelate Stability
85
ACID
Constants
method
4.49 5.62 7.74
II
with Glutamic
Acid
at 25°C.
logK,
ion
Mg++ Ca++ sr++ Ba++ Cd++
1.9 1.43 1.37 1.28 3.9
the value obtained in the present investigation is considerably lower than the average, but agrees very well with that of Schmidt, Kirk, and Appleman (9), namely 9.47. Although complex formation and chelation have been found to be possible with all the metals of the periodic table (6), there are vast differences in the stabilities of the products formed. Several very plausible theories have been advanced to enable one to predict, with some degree of certainty, relative chelating possibilities. We shall discuss briefly some generalizations which have bearing on the present investigation. Included is a table (Table II) of the log K1 values obtained by Lumb and Martell (2) in their study of glutamic acid chelates with the alkaline-earth metals and with cadmium. It is to be noted that the above constants for the alkaline-earth metals are much lower than those obtained with the transition elements. No doubt, this is due to the fact that the transition metals have a d-orbital available for covalent linkage, whereas the alkaline earths form bonds that are primarily ionic. According to Martell and Calvin (6), the metals which form the strongest homopolar bonds form the most stable chelates. CONCLUSION
Chelate stability depends on a large number of factors inherent in both the chelating agent and the metal ion. In investigating chelate stability
86
NYBERG,
CEFOLA,
AND
SABINE
Log
I
2
3
Electronegotivity FIG. 2. Relationship between negativity of the bivalent metal
first stability ions.
constants
of chelates
and electro-
Log
Second Ionization
FIG. 3. Relationship ionization
potentials
Potential
between first stability constants of the bivalent metal ions.
of chelates
and
the following relations have been observed. They are represented graphically in Figs. 2 and 3. 1. For a given charge type, log K, is proportional to the electronegativity of the metal ion. This holds for P-diketones (lo), and applies to the present investigation also. The plot of log K1 vs. electronegativity has been made to include both the alkaline earths and the transition metals to show the wide applicability of this theory.
87
CHELATION WITH GLUTAMIC ACID
2. Calvin and Melchior (11) and Van Uitert (10) have shown that log K1 can be expressed as a function of the second ionization potential for bivalent ions of the first transition series. Data given herein substantiates this proposition. We notice that log K, increases with I, for the alkaline earths also, but at a much slower rate. Magnesium falls on both straight lines. 3. Mellor and Maley (la), as well as Irving and Williams (13), showed that in the first transition series there was an almost linear relationship between successive stability constants of the bivalent ions and the atomic numbers. Martell and Calvin (6) concluded that the value of the stability constant must increase in a regular manner with some property related to atomic number, perhaps to the regular increase in the number of 3d electrons. This leads to a consideration of effective atomic number as a factor in stability, since the stability of coordination compounds is sometimes related to the proximity of the electronic structure of the central metal to that of the next rare gas in the period (14). Stability decreases with increasing deficiency of electrons. This is the case with the glutamic acid complexes of cobalt, nickel, and copper. However, Bailar’s book points out that this principle can be applied only to a limited number of complexes, and should not be used as a sole criterion of the stability of any complex. Therefore, this proposition warrants further investigation. This research was supported Contract AT (30.1)906.
by a grant from the Atomic
Energy
Commission
under
SUMMARY
Stabilities of chelates of glutamic acid with bivalent cobalt, nickel, and copper have been determined by potentiometric titration. Stability was seen to increase with electronegativity, second ionization potential, and atomic number of the metal. ACKNOWLEDGEMENT REFERENCES 1. CHABEREK, S., JR., AND MARTELL, A. E., J. Am. Chem. Sot. 74, 5052 (1952). 2. LUMB, R. F., AND MARTELL, A. E., J. Phys. Chem. 67,690 (1953). 3. SCHWARZENBACH, G., AND BIEDERMANN, W., Helv. Chim. Acta 31, 339 (1948). 4. BJERRUM, J., “Metal Ammine Formation in Aqueous Solution.” P. Haase and Son, Copenhagen, 1941. 5. HARNED, H. S., AND OWEN, B. B., “The Physical Chemistry of Electrolytic Solutions,” 2nd ed. Reinhold Publ. Corp., New York, 1950. 6. MARTELL, A. E., AND CALVIN, M., “Chemistry of the Metal Chelate Compounds.” Prentice-Hall, Inc., New York, 1952. 7. HARNED, H. S., AND COOK, M. A., J. Am. Chem. Sot. 59, 1291 (1937). 8. CHABEREK, S., JR., COURTNEY, R. C., AND MARTELL, A. E., J. Am. C&m. SOC. 74, 5057 (1952).
88
NYBERG,
CEFOLA,
AND
SABINE
9. SCHMIDT, C., KIRK, P., AND APPLEMAN, W., J. Biol. Chem. 88,285 (1930). 10. VAN UITERT, L. G., Ph. D. Thesis, The Pennsylvania State University, 1952; FERNELIUS, W. G., AND VAN UITERT, L. G., Acta Chem. &and. 8, 1726 (1954). 11. CALVIN, M., AND MELCHIOR, N.C., J. Am. Chem. Sot. 70, 3270 (1948). 12. MELLOR, D. P., AND MALEY, L. E., Nature 161, 436 (1948). 13. IRVING, H. M., AND WILLIAMS, R. J. P., Nature 162, 746 (1948). 14. BAILAR, J. C., JR., ed., “The Chemistry of the Coordination Compounds.” Reinhold Publ. Corp., New York, 1956. 15. BJERRUM, J., SCHWARZENBACH,G., AND SILLEN, L.G.“StabilityConstants.Part 1: Organic Ligands.” Special Publ. No. 6. The Chemical Society, London, 1957.
OF
BIOCHEMISTRY
AND
86, 82-88 (1959)
BIOPHYSICS
Chelation of Cobalt, Nickel and Copper with Glutamic Acid Sister M. Helen Therese Nyberg’ From the Chemistry
Department,
Malloy Catholic College for Women, Rockville Long Island, New York
Centre,
Michael Cefola From
the Department
of Chemistry,
Fordham
University,
New
York,
New
York
and
David Sabine From the United
States Vitamin Yonkers,
Corporation,
Received
March
and Pharmaceutical New York 12, 1959
The present investigation is intended to supply further evidence for proposed theories regarding some general trends in the stability of metal chelates. Cobalt, nickel, and copper were examined for relative stabilities with glutamic acid. Amino acids, in general, are said to combine with the more basic metal ions as the alkaline earths, and to have an appreciable affinity for the transition metals (1). Lumb and Martell (2) found stability constants for the chelates of the alkaline-earth metals with glutamic acid. It will be interesting to compare them with those of the transition-element complexes. EXPERIMENTAL Apparatus,
Materials
and Method
The first and second ionization constants of glutamic acid were obtained by potentiometric titration at 30°C. For the determination of the metal chelate constant, a 1: 1 ratio of metal to acid was used. Solutions were made up to such concentration as to maintain ionic strength of 0.1 throughout the titration. Potassium chloride was the electrolyte used. Experimental solutions were made to contain a final acid and metal concentration of 0.01 M. 1 This thesis was presented by Sister M. Helen Therese Nyberg, 0. P., in partial fulfillment of the requirements for the degree of Master of Science at Fordham University, 1957. 82
CHELATION
WITH
GLUTAMIC
ACID
83
The nL-glutamic acid was purchased from Mann Research Laboratories, New York, and was used without further purification. The standardization of the acid was accomplished potentiometrically, using a differential curve to determine the end point. The metal-chloride solutions were prepared from Fisher Certified reagents, and were standardized with silver nitrate, using dichlorofluorescein as indicator. The carbon dioxide-free KOH solution was prepared according to the method of Schwarzenbach and Biedermann (3). It was standardized for basic strength using potassium acid phthalate with phenolphthalein as indicator, and was also standardized for chloride ion. The titration was carried out in nitrogen atmosphere using a 250.ml. flask with five necks to accommodate the tube delivering nitrogen, a mercury-seal stirrer, a Beckman # 1190-80 glass eleci:-.: e, a Beckman K calomel electrode, and a buret. Small increments of base (0.1 M) were added until a total of 20 ml. was reached. pH readings were made on a Beckman Model G pH meter.
Calculations Figure 1 shows curves obtained for the titration of glutamic acid in the absence of and in the presence of the metal ions. The dissociation constants of the acid (k, and Icz), as well as the first stability constants (KI) of the chelates, were calculated according to the algebraic method described by Chaberek and Martell (1). The Bjerrum
PH
1
1
1.0
1.5
a ( Moles
bose /mole
acid 1
FIG. 1. Titration curves for glutamic acid and for the cobalt (II), nickel (II), and copper (II) chelates of glutamic acid at 30°C. The mark ( on the curves indicates the pH at which the hydroxide precipitation begins. Points on the curve beyond these pH values are not very reliable.
84
NYBERG, CEFOLA, AND SABINE
method of plotting fi (number of ligands bound per metal ion) vs. pA (negative log of the A= concentration) was employed also in the case of the chelate constants. At rt = 0.5, pA = log K1 (4). Activities of H+ ions were obtained directly from pH measurements. The OH- ion activity was calculated from a knowledge of the fact that the ionization constant of water at 1 atm. pressure and 30°C. is 1.471 X lo-l4 (5). In obtaining actual concentrations from activities, it was assumed that y&i= yfH+ = yfoH(6). A large-scale graph for molality vs. mean activity coefficient (yf) of the chloride ion at 30°C. was made from data given by Harned and Cook (7), and extrapolated to zero concentration. From this graph, activity coefficients were read for the chloride ion and, hence, for the hydrogen and hydroxide ions also. Hydroxide formation occurred in all the metal chelate titrations. Chaberek et al. (8) made a study of this phenomenon and gave a summary of their results. Because of this hydrolysis interference at higher pH values, K1 was calculated algebraically from points along the upper buffer region below the pH where the hydroxide begins to form. Values checked very well with the results obtained by the Bjerrum method. DISCUSSION OF RESULTS
Results are represented in graphic form in Fig. 1. The curve for the free acid indicates a two-step ionization. The reaction of the metal ion with the acid is accompanied by the displacement of protons from the acid with the consequent lowering of pH. Therefore, the curves which represent the chelates have an upper buffer region lower than that of the free acid, the degree of lowering being proportional to the stability of the chelate formed. In the case of the cupric and nickelous ions, even the first buffer region is lowered. For these metals we have: M++ + H2A ti MA + 2H+. Relative positions of the metal curves indicate relative stabilities to be Cu > Ni > Co. In the case of copper and nickel, we also have an inflection point at a = 2, which indicates a rather strong 1: 1 complex formation. Table I gives values obtained for the above constants. These values are in reasonable agreement with those already reported in the literature (15), although the temperature and ionic strength at which these measurements were made are somewhat different. It is interesting to note that the pi& of glutamic acid is approximately an average of the different values listed by Lumb and Martell (2), if those which were corrected to zero ionic strength are omitted. In the case of pkz ,
CHELATION
WITH
GLUTAMIC
TABLE Chelate Stability Glutamic
acid:
I
Constants
at SO”C.
pkt = 4.23; p/c%= 9.46. p = 0.1 log
Metal
ion Algebraic
co++
TABLE
Metal
KI
method
Bjerrum
4.48 5.60 7.78
hTiTii-* cu++
Chelate Stability
85
ACID
Constants
method
4.49 5.62 7.74
II
with Glutamic
Acid
at 25°C.
logK,
ion
Mg++ Ca++ sr++ Ba++ Cd++
1.9 1.43 1.37 1.28 3.9
the value obtained in the present investigation is considerably lower than the average, but agrees very well with that of Schmidt, Kirk, and Appleman (9), namely 9.47. Although complex formation and chelation have been found to be possible with all the metals of the periodic table (6), there are vast differences in the stabilities of the products formed. Several very plausible theories have been advanced to enable one to predict, with some degree of certainty, relative chelating possibilities. We shall discuss briefly some generalizations which have bearing on the present investigation. Included is a table (Table II) of the log K1 values obtained by Lumb and Martell (2) in their study of glutamic acid chelates with the alkaline-earth metals and with cadmium. It is to be noted that the above constants for the alkaline-earth metals are much lower than those obtained with the transition elements. No doubt, this is due to the fact that the transition metals have a d-orbital available for covalent linkage, whereas the alkaline earths form bonds that are primarily ionic. According to Martell and Calvin (6), the metals which form the strongest homopolar bonds form the most stable chelates. CONCLUSION
Chelate stability depends on a large number of factors inherent in both the chelating agent and the metal ion. In investigating chelate stability
86
NYBERG,
CEFOLA,
AND
SABINE
Log
I
2
3
Electronegotivity FIG. 2. Relationship between negativity of the bivalent metal
first stability ions.
constants
of chelates
and electro-
Log
Second Ionization
FIG. 3. Relationship ionization
potentials
Potential
between first stability constants of the bivalent metal ions.
of chelates
and
the following relations have been observed. They are represented graphically in Figs. 2 and 3. 1. For a given charge type, log K, is proportional to the electronegativity of the metal ion. This holds for P-diketones (lo), and applies to the present investigation also. The plot of log K1 vs. electronegativity has been made to include both the alkaline earths and the transition metals to show the wide applicability of this theory.
87
CHELATION WITH GLUTAMIC ACID
2. Calvin and Melchior (11) and Van Uitert (10) have shown that log K1 can be expressed as a function of the second ionization potential for bivalent ions of the first transition series. Data given herein substantiates this proposition. We notice that log K, increases with I, for the alkaline earths also, but at a much slower rate. Magnesium falls on both straight lines. 3. Mellor and Maley (la), as well as Irving and Williams (13), showed that in the first transition series there was an almost linear relationship between successive stability constants of the bivalent ions and the atomic numbers. Martell and Calvin (6) concluded that the value of the stability constant must increase in a regular manner with some property related to atomic number, perhaps to the regular increase in the number of 3d electrons. This leads to a consideration of effective atomic number as a factor in stability, since the stability of coordination compounds is sometimes related to the proximity of the electronic structure of the central metal to that of the next rare gas in the period (14). Stability decreases with increasing deficiency of electrons. This is the case with the glutamic acid complexes of cobalt, nickel, and copper. However, Bailar’s book points out that this principle can be applied only to a limited number of complexes, and should not be used as a sole criterion of the stability of any complex. Therefore, this proposition warrants further investigation. This research was supported Contract AT (30.1)906.
by a grant from the Atomic
Energy
Commission
under
SUMMARY
Stabilities of chelates of glutamic acid with bivalent cobalt, nickel, and copper have been determined by potentiometric titration. Stability was seen to increase with electronegativity, second ionization potential, and atomic number of the metal. ACKNOWLEDGEMENT REFERENCES 1. CHABEREK, S., JR., AND MARTELL, A. E., J. Am. Chem. Sot. 74, 5052 (1952). 2. LUMB, R. F., AND MARTELL, A. E., J. Phys. Chem. 67,690 (1953). 3. SCHWARZENBACH, G., AND BIEDERMANN, W., Helv. Chim. Acta 31, 339 (1948). 4. BJERRUM, J., “Metal Ammine Formation in Aqueous Solution.” P. Haase and Son, Copenhagen, 1941. 5. HARNED, H. S., AND OWEN, B. B., “The Physical Chemistry of Electrolytic Solutions,” 2nd ed. Reinhold Publ. Corp., New York, 1950. 6. MARTELL, A. E., AND CALVIN, M., “Chemistry of the Metal Chelate Compounds.” Prentice-Hall, Inc., New York, 1952. 7. HARNED, H. S., AND COOK, M. A., J. Am. Chem. Sot. 59, 1291 (1937). 8. CHABEREK, S., JR., COURTNEY, R. C., AND MARTELL, A. E., J. Am. C&m. SOC. 74, 5057 (1952).
88
NYBERG,
CEFOLA,
AND
SABINE
9. SCHMIDT, C., KIRK, P., AND APPLEMAN, W., J. Biol. Chem. 88,285 (1930). 10. VAN UITERT, L. G., Ph. D. Thesis, The Pennsylvania State University, 1952; FERNELIUS, W. G., AND VAN UITERT, L. G., Acta Chem. &and. 8, 1726 (1954). 11. CALVIN, M., AND MELCHIOR, N.C., J. Am. Chem. Sot. 70, 3270 (1948). 12. MELLOR, D. P., AND MALEY, L. E., Nature 161, 436 (1948). 13. IRVING, H. M., AND WILLIAMS, R. J. P., Nature 162, 746 (1948). 14. BAILAR, J. C., JR., ed., “The Chemistry of the Coordination Compounds.” Reinhold Publ. Corp., New York, 1956. 15. BJERRUM, J., SCHWARZENBACH,G., AND SILLEN, L.G.“StabilityConstants.Part 1: Organic Ligands.” Special Publ. No. 6. The Chemical Society, London, 1957.