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Kinetics and mechanism of divalent-metal ion catalyzed hydrolysis of a β-ketoimine
J. inorg,nucl. Chem., 1968, Vol. 30, pp. 1931 to 1940. PergamonPress. Printedin Great Britain
KINETICS AND ION CATALYZED
MECHANISM OF DIVALENT-METAL HYDROLYSIS OF A fl-KETOIMINE D. F. M A R T I N
Department of Chemistry, University of South Florida, Tampa, Florida
and F. F. C A N T W E L L * Department of Chemistry and Chemical Engineering, University of Illinois, Urbana, Illinois
(First received 13 March 1967; in revised form 14 December 1967)
A b s t r a c t - T h e kinetics of acid hydrolysis of a quadridentate/3-ketoimine, bisacetylacetoneethylenediimine (H~Ch), in the presence of divalent-metal ions have been measured at constant pH by a titrimetric procedure, using 50 per cent (v/v) dioxane-water solutions at 30 °. Divalent metal ions studied included: copper, nickel, beryllium, and magnesium ions. Pseudo-first-order rate constants, kobs., are shown to vary with the ratio CM/Cmch, and two limiting cases are established. The first, (k'obs.) attained as the ratio approaches unity, appears to correspond to the rate of hydrolysis of the pure metal-chelate compound; the other limiting rate (kohs.) is taken to be the rate of hydrolysis in the absence of metal ions. The ratio k'obs/kobs, is evaluated, and correlations with formation constants and the el~ctronegativity of the metal are considered.
ALL SCHIFF bases of the type R2C------NR' are subject to acid hydrolysis, though /3-ketoimines of the type RCOCH2C(R')------NR" appear to be much less susceptible to this reaction. In the presence of metal ions, marked changes in the rate of hydrolysis have been noted. Depending upon the structure of the ligand, the metal ion can increase or decrease the hydrolytic tendency. Bis-thiophenal ethylenediimine, C4H3SCH~NCH2CH2N~-------CHC4H3S, is hydrolyzed much more readily in 50 per cent aqueous ethanol in the presence of copper(II) ions [1 ], whereas the terdentate ligand o-OCrH4CH--NCH2COOR is stabilized[2] by the presence of copper(II) ions, and exists as a metal chelate compound, CuCh-H20. The destabilization has been attributed to a polarization of the carbon-nitrogen bond which results in increased susceptibility to electrophilic attack, and the stabilization has been attributed to the increased stability of the metal complex. The participation of metal ions in the hydrolysis of carbon-nitrogen bonds is of considerable interest, and it seemed desirable to have available more kinetic data and more information on the influence of representative- and transition-metal ions in order to test the proposed mechanisms of hydrolysis. A quadridentate ligand, bisacetylacetoneethylenediimine, was selected for study because the rates of hydrolysis for this and related ligands, as well as the copper(II) derivatives have been determined [3]. *Present address: Endo Laboratories, Garden City, N e w York. 1. G. L. Eichhorn, andJ. C. Bailar, Jr.,J.Am. chem. Soc. 75, 2905 (1954). 2. G. L. Eichhorn and N. D. Marchand, J. Am. chem. Soc. 78, 2688 (1956). 3. D. F. Martin and F. F. Cantwell, J. inorg, nucl. Chem. 26, 2219 (1964). 1931
1932
D . F . M A R T I N and F. F. C A N T W E L L
CH3COCHsC(CHa)NCHsCHsN(CHa)CCHsCOCH3 + 2HsO ---> 2 C H 3 C O C H s C O C H 3 + HsNCHzCHsNHs.
(1)
It would also be desirable to know what relationship exists between the chelating tendency of a ligand with a given metal ion and the ease of hydrolysis. Unfortunately, few stability or kinetic data are available. The chelating tendencies of various bidentate/3-ketoimines with some divalent metal ions have been determined[4] as have the rates of hydrolysis of certain quadridentate fl-ketoimines and their copper(II) derivatives [3]. It seemed desirable to determine the rates of hydrolysis of a standard quadridentate fl-ketoimine in the presence and absence of metal ions and to determine the effect, if any, of expected stability trends on hydrolytic tendencies. EXPERIMENTAL Materials
Analytically pure samples of the ligand were used as before[3]. Standard solutions of metal perchlorates were used, except for beryllium nitrate. Determination o f rates o f hydrolysis Specific rate constants for the (pseudo) first-order hydrolysis ofligand in the presence of metal ions were determined titrimetrically, using the previously described [3] procedure. As before, rates were determined in 50 per cent (v/v) dioxane at 30 °. The initial pH was adjusted by addition of standard perchloric acid (or tetramethylammonium hydroxide for the copper-ligand systems). Pertinent data are summarized in Tables 1 and 2. In general, the rate constants were reproducible within _+4 per cent.
RESULTS
AND DISCUSSION
In dioxane-water solution, the ligand H2Ch may participate in several equilibria HsCh + H + ~ HsChH +
K1
(2)
HaCh++ H + ~ HsChH~ +
Ks
(3)
HsCh ~ H C h - + H +
KD1
(4)
HsCh- ~ Ch s- + H +
KD2.
(5)
And in the presence of divalent metal ions, additional equilibria may be expected including complex formation: M s+ + Ch s- ~ MCh
Ks.
(6)
It may be expected that metal species MCh will undergo protonation to form M C h H ÷ and MChH] + with protonation constants KI and K~' respectively. It may be presumed that all of the above species will undergo hydrolysis, though at significantly differing rates. On the basis of a previous study [3] it appears reasonable to suggest that the process of hydrolysis of the metal-chelate compound at pH 4.2 may be represented by the following equations: 4. D. F. Martin, G. A. Janusonis and B. B. Martin, J. A m. chem. Soc. 83, 73 (1961).
Hydrolysis of a/3-ketoimine
~C-- O \ HC//
O--C / M/ ~ , %CH ~./C-- %,N/ ~N--C / /. -CH CH 8 + ~ z z
H,O"1"
K', --~=-~i---
1933
~C--O. O--C( HC// \M / "~CH ~.C_N /
/ Slow
\ / CH # C - - O \ M / O--C~-C H \C ~-_0/ ~ O _ C / / \ (¢) 4H2NCH2 CHz NH3'
/
(d) + + H3 N-CH2CH2 -NH t
O--C
/
(a)
C--O N /O--C / HC /M "~CH C / = ~1 \ O = C / NH3+\ CHz CH2 (b)
/ XC--O\ -0--C% HC// M + /CH • \C =0 /
k'i
\ H20 fast
\N_C
\ /H I CH2CHzH/O\H
\
C--O OH2 H+ slow HC// \M+/ \ C ' N / ~NH --5---/ \ / 2 k2
CH2CHz (e) 4CH3 COCHz COCH~
It may also be assumed from previous work[3] that the observed rate of hydrolysis will be less for the metal-chelate compound than for the ligand. At a given pH, the concentration of the various metal containing species will be determined by the relative concentrations of ligand and metal ion, the pH, and the co-ordinating tendency of the ligand for a given metal ion. In this study a moderate value of pH (ca. 4.2) was selected, the concentration of the ligand was held constant, and the concentration of metal ion varied. It may be expected that there will be two limiting cases, assuming the effect of metal ion is to form the metal-chelate compound which will be more resistant to hydrolysis. The first limiting case will be the rate of hydrolysis in the absence of metal ion. The second limiting case will be the rate of hydrolysis when there is a sufficient excess of metal to insure that all of the ligand is present as the metalchelate compound*. The rate of hydrolysis should then correspond to the rate of hydrolysis for the metal-chelate compound at the same pH. This is apparent, if it *This, of course, presupposes that two conditions are satisfied; the value of Kr is relatively large and the rate of complex formation is rapid. The first condition is no doubt satisfied, based on the values of metal derivatives of bidentate Schiff bases [4]. The characteristic colors of the copper(l I) and nickel(II) derivatives of the quadridentate Schiffbase are formed instantly on mixing.
1934
D . F . MARTIN and F. F. C A N T W E L L
is assumed that the observed rate is equal to kobs.CH,ch , where CHzCh is the total concentration of quadridentate ligand, and if the species of major importance are H2Ch, H2Ch +, H4Ch 2+, MCh, MChH +, MChH22+, for which the specific hydrolyt I sis constants are kl, k2, ka, k~, k s, and k 3. It has been found that in the absence of metal ion, Equation (7) is valid: kobs. =
k, + k2K1 [H +] + k3K,K~ [H+] 2 1 + K , [ H +] + KIK2[H+] *
(7)
and that when essentially all of the ligand has been converted to metal-chelate species the rate data are best represented as: I t + t t t + 2 klP + k2KI[H ] + k3K1K2[H ] Robs. ~-" kob~.- 1+ K~ [H +] + K~K~ [H+] 2
(8)
The evaluation of the various constants was described previously [3]. In order to obtain a rate law to include all metal-ligand ratios, it would be necessary to evaluate Kj, which was not possible because of hydrolysis. Nor was it possible to evaluate KD~ and Kr~; again because of hydrolysis. However it was possible to study the second limiting cases of hydrolysis of ligand in the presence ofcopper(II) or nickel(II) ions. The pertinent data, the variation of the observed specific rate constant, kobs., as a function of the ratio metal ion-chelating agent are listed in Table 1 and plotted as Figs. 1 and 2. In general it appears that the variations ofkobs, vs. CM2+/CH~chare best interpreted as a smooth curve, rather than, say, as two linear portions. The single exception is the curve for the copper(tI)-ligand system, (Fig. t) for which two linear portions might be distinguished. This is exceptional because for this metal-ligand system alone it was necessary to add initially tetramethylammonium hydroxide to increase the pH to 4.2; all other systems required addition of acid. The values of kob~. for the hydrolysis of quadridentate figand in the presence of copper(II) or nickel(II) ions approach a minimum as the metal-ligand ratio approaches unity (Fig. 1). The minimum values of kob~. are approximately those observed for the pure metal-chelate compound. For example, the values of kob~.for the hydrolysis of the pure copper(II)- and nickel(II)-chelate compounds are 1.5 x 10-5 (calc. for pH, 4.16), and 1.91 x 10-4 (pH, 4.16) sec -1, respectively. The minimum values of k'ob~, at pH 4.16 in the presence of copper(II) and nickel(II) ions are 0.5 × 10-5 and 1.85 x 10-4 s e c -1, respectively. Thus, there is fairly good agreement between the expected and observed limiting values for these two metal ions. On the basis of this information, it appears reasonable to postulate two limiting rates, and to postulate that the reduction in rate of hydrolysis is due to the formation of the metal-chelate compound, which hydrolyzes at a reduced rate. This is not to imply that the formation constants for the copper(II) and nickel(II) complexes are the same. Quite obviously, they differ substantially as indicated by the differences in initial pH values which required adjusting with base to bring the initial pH to 4.2 in the case of some copper(II) systems. The ligand-beryllium and -magnesium systems are somewhat different, and require additional comment. The data for quadridentate ligand-beryllium(II) and -magnesium(II) systems suggest that beryllium co-ordination entities are formed, and it appears that limit-
H y d r o l y s i s of a fl-ketoimine
1935
Table 1. Specific rate c o n s t a n t s for the hydrolysis of CH3COCH2C(CH3)NCH2CHzN(CH:~)CCH2COCH3(HzCh) in the presence of divalent-metal ions in 50vol. per cent d i o x a n e - w a t e r at 30°C and pH 4.16
CH2Ch Metal ion
103 M
C M2~/CH2ch
104kob...... -~
Copper(lI)
3.99 3.99 3-96 3-98 4-00 3-93 3.89
0.079 0.264 0.529 0.790 1.05 1.06 1.32
2.57 1.90 1.30 0-672 0.066 0.138 (?) 0.046
Nickel(ll)
3-98 3-98 3"98 3"97 3-98 3.97 3 -99 3.98 3"98 3.98
0.053 0.103 0.264 0'264 0.422 0.530 0.727 0.829 0-935 1-19
2-87 2-45 2-41 2-64 2.46 2-36 2-04 1-90 1.94 1.78
Beryllium
3-98 3-98 3 98 3.96 3"98 3-85 3.98 3.99 3'99 3.99 3 '99 3'99 3"85 3'86 4-00 3'97
0-050 0-050 0.131 0-131 0.251 0-311 0.414 0.498 0-504 0.640 0-644 0.828 0.946 1-20 1.30 1"31
2'40 2.80 2.12 2.14 1-98 1.85 1-81 1.55 1.60 1.52 1.69 1.35 1.34 1.13 1.12 1.19
Magnesium
3 -98 3.98 3"98 3-98 3-98 3.98 3.97 3"98 3.97
0-075 0.151 0.226 0-251 0.405 0-603 0.989 1.03 1.21
1-96 2.62 2-43 2.38 2.06 1.92 1.82 1-72 1-70
1936
D . F . M A R T I N and F. F. C A N T W E L L
i
3"0
T
u
Ni t +
0
Cu t'+
•
2-0
I&l O)
0
O ~ o
o
I'0
o
I 0
0.2
I
I
0.4
0-6
0-8
I'0
1-2
CM / C H=Ch Fig. 1. Variation of specific rate constant (104kob,., sec -1) for hydrolysis with nickel (open circles) and copper ions (closed circles).
CM2+/CHichfor
ing rates are approached when the metal ion-chelating agent ratio is 1.0. Moreover, the limiting values of kobs. are of the expected order, being 1.33 × 10-4 sec -1 and 1.80 x 10-4 sec -x, for the beryllium and magnesium systems, respectively. The beryllium and magnesium derivatives of the quadridentate ligand have not been isolated. And until recently there was good reason to doubt that berylliumnitrogen co-ordination occurred in aqueous solution. The lack of evident coordination of beryllium ion with such strong chelating agents as picolinic acid, glycine, iminodiacetic acid, and fl-alanine has been noted[5]. On the other hand, formation constants were obtained for beryllium(I I) derivatives of Schiff bases of the type RCOCH2C(NR')R" in 50 vol. per cent dioxane-water; for these ligandberyllium systems, no evidence of hydrolysis was detected[4]. Still more recently, beryllium derivatives of the Schiff bases o-OC6HaCH(--NC4Hg)- and (o-OC6H4CH--NCH2CH2N~-----CHC6H40-o) ~- have been prepared[7]. Though these results have been repeated, attempts to isolate bisacetylacetoneethylenediiminoberyllium(lI) have not been successful. Nevertheless, there is evidence that the effectiveness of beryllium in reducing the rate of hydrolysis is due to an effect of complex formation and not just to an 5. D. Sen, Unpublished results [6]. 6. R. M. Izatt, W. C. Fernelius, and B. P. Block, J. phys. Chem. 59, 80 (1955). 7. R.W. Green and P. W. Alexander, Aust. J. Chem. 18, 1297 (1965).
Hydrolysis of a fl-ketoimine
1937
I
o
M~ 2"f 0
3.0
5
Be=+ •
2.0
(.3 i,i (/)
•
0
"
O--
"o 1.0
o 0
I
I
0-2
0.4
I
I
0 "6
0-8
1 1.0
I I-2
CM /CHINCH
Fig. 2. Variation of specific rate constant (104kobs., sec -~) for hydrolysis with CM~+/Crbch for magnesium (open circles) and beryllium ions (closed circles).
effect of ionic strength. This evidence is the pH dependence of the ligand hydrolysis, which is given by the data in Table 2 and Fig. 3 ; the pH dependence is clearly different in the presence of beryllium ion (CBe2+/CH2ch' c a . 1) (Fig. 3). At a low pH, c a . 2.0, there is little difference in rate of hydrolysis when beryllium ion is present or absent; at a pH of 4 the rate has been reduced markedly. This is attributable to nearly complete absence of complex formation at the low pH and nearly complete formation of a beryllium complex at the higher pH. From the data represented in Fig. 2, it would appear that the 1 : 1 species is the most likely'present when the limiting rate is observed. However, with the absence of formation-constant data and hydrolysis rates of the pure metal-chelate compound, it is not possible to be certain of the identity of the 1 : 1 species. This is true of beryUium; for that matter, it is true of the copper and nickel systems for which the hydrolysis data are available. With the latter two systems, it is possible only say the hydrolyzing species are the same. In this connection, it is both interesting and instructive to examine the ratio kobs./kobs., which is the ratio of the limiting rate constants in the presence and absence of divalent metal ions, respectively. The values are 0.017 (CuZ+), 0.46 (BeZ+), 0.65 (Ni 2+) and 0-61 (Mg2+). It seems reasonable to grant that the value of k'obs, approximates the constant for the hydrolysis of the pure metal-chelate compound at the pH selected. It might be tempting to seek a correlation between
1938
D . F . M A R T I N and F. F. C A N T W E L L Table 2. Specific rate constants for the hydrolysis of H2Ch, effect of pH*
Metal ion
C.,ch 103 M
Cu~+ Ca,ch
pH
Copper(II)
4.00
1.01
2"39
1 "98
1"02
3'35
1"64
3"92
1"30
3"36
1"65
Beryllium
3.99 3.99 3.97 3.97
1.16 1.16 1.18 0.90
2.22 2.41 2.58 3.36
Nickel(II)
3.94
1.04
3.36
5.58
Magnesium
3.96 3.96
0.80 0.82
3-36 3-36
5.89 6.10
104kobs.,see-' 13'2
21.0 18.3 10.8 2.7
*t = 30°; Solvent, 50 vol. per cent dioxane-water.
oo e ~ \ 3-0
_
•
O0
o o v
o
0
-J 3 . 5 _ i
e
4.0
, 2-0
i
I
i
I
i
3.0
i
i
,
I
~'
4"0 pH
Fig. 3. Variation of specific rate constant for hydrolysis ( - l o g kobs.) of H2Ch as function of pH calculated (--), observed in absence of metal ions (open circles) and in the presence of beryllium ion (CRez+/Cmc h c a . 1-0; closed circles).
Hydrolysis of a/3-ketoimine
1939
these ratios and a stability order, though such a correlation does not seem valid in this case; the positions of nickel and magnesium are clearly reversed. It might be suggested that the slope of the curved portions, or ligand-excess regions, of Figs. 1 and 2 at certain values of CM2+/CH2ch could be correlated with stability inasmuch as the slope would be a measure of the amount of metal-chelate compound being formed. On the other hand, at the limiting hydrolysis rates, related to k'obs., the concentration of complex should be substantially the same, and k'ob~. should be insensitive to the formation constant, Kz. This is true if all metalchelate compounds studied have relatively large values of Ke. It seems reasonable to ascribe the differences in values of k'obs./kobs, to two sources: (1) the basicities of the ligand vs. metal-chelate compound, largely as reflected in the values K1 and K~; and (2) the ease of breaking the protonated carbon-nitrogen bonds, largely as reflected in the values k2 vs. k,~.Two synergistic effects may be expected: the act of co-ordination should cause a greater polarization of the carbon-nitrogen bond and this should be reflected in large values of K1 [8]. On the other hand the protonated carbon-nitrogen bond is less susceptible to cleavage when the nitrogen is attached to a metal atom because an additional bond must now be broken to permit proton abstraction from solvent or hydronium ion to form the kinetically determined products, an amine and metal-/3diketone complex; this would be reflected in reduced values of k'l and k~ relative to k~ and k2. Granting the importance of the polarizing action of the metal, a correlation might be expected between kobs./kobs, and XM, the electronegativity of the metal. Using electronegativity data of Allred[9], the correlation is fair and the anomalous position of nickel is again noted. Now it becomes apparent that a more detailed analysis of the value of k~bs.for the nickel(I I) compound is in order. The significant differences between the nickel(lI) and copper(ll) derivatives appear to be in the value of k~ which is greater for the nickel(I I) derivative by a factor of 25 [ 10]. The values of k.; are comparable. Thus the anomaly in kobs. seems to be due to differences in k I. This may be due to the greater predilection of the nickel(II) derivatives to form five- and six-co-ordinate entities and the consequence that a Grotthus proton-transfer chain ends nearer the co-ordinated nitrogen. H
H
H
I
I
\/
H--O--H ... O--H ... O
H ~
H
H
I
L
+
H - - - O . . . H - - - O . . . H----O--H
Obviously, it would be desirable to test this suggestion by examining other nickel(II)-Schiffbase complexes. Unfortunately, most of the pertinent compounds, 8. For the quadridentate ligand used in this study the values are: 103kl, 2.40 sec -1 and 103k.,, 8.25 sec-'; and for the copper(II) derivative; 103k,, nil; 103k2, 0.269 sec-L 9. A. L. Allred, J. inorg, nucl. Chem. 5,264, 269 (1958). 10. E.J. Olszewski, Thesis, University of Illinois (1964).
1940
D . F . M A R T I N and F. F. C A N T W E L L
prepared in connection with other studies [11], are insoluble in 50 or even 75 vol. per cent dioxane-water media. Apparently, the synthesis of more soluble compounds must be achieved before further studies are possible. Acknowledgment-This research was supported by P. H. S. Research Grants 07873 and 01204, Division of General Medical Sciences, U.S. Public Health Service. 11. E.J. Olszewski and D. F. Martin, J. inorg, nucl. Chem. 26, 1577 (1964).
KINETICS AND ION CATALYZED
MECHANISM OF DIVALENT-METAL HYDROLYSIS OF A fl-KETOIMINE D. F. M A R T I N
Department of Chemistry, University of South Florida, Tampa, Florida
and F. F. C A N T W E L L * Department of Chemistry and Chemical Engineering, University of Illinois, Urbana, Illinois
(First received 13 March 1967; in revised form 14 December 1967)
A b s t r a c t - T h e kinetics of acid hydrolysis of a quadridentate/3-ketoimine, bisacetylacetoneethylenediimine (H~Ch), in the presence of divalent-metal ions have been measured at constant pH by a titrimetric procedure, using 50 per cent (v/v) dioxane-water solutions at 30 °. Divalent metal ions studied included: copper, nickel, beryllium, and magnesium ions. Pseudo-first-order rate constants, kobs., are shown to vary with the ratio CM/Cmch, and two limiting cases are established. The first, (k'obs.) attained as the ratio approaches unity, appears to correspond to the rate of hydrolysis of the pure metal-chelate compound; the other limiting rate (kohs.) is taken to be the rate of hydrolysis in the absence of metal ions. The ratio k'obs/kobs, is evaluated, and correlations with formation constants and the el~ctronegativity of the metal are considered.
ALL SCHIFF bases of the type R2C------NR' are subject to acid hydrolysis, though /3-ketoimines of the type RCOCH2C(R')------NR" appear to be much less susceptible to this reaction. In the presence of metal ions, marked changes in the rate of hydrolysis have been noted. Depending upon the structure of the ligand, the metal ion can increase or decrease the hydrolytic tendency. Bis-thiophenal ethylenediimine, C4H3SCH~NCH2CH2N~-------CHC4H3S, is hydrolyzed much more readily in 50 per cent aqueous ethanol in the presence of copper(II) ions [1 ], whereas the terdentate ligand o-OCrH4CH--NCH2COOR is stabilized[2] by the presence of copper(II) ions, and exists as a metal chelate compound, CuCh-H20. The destabilization has been attributed to a polarization of the carbon-nitrogen bond which results in increased susceptibility to electrophilic attack, and the stabilization has been attributed to the increased stability of the metal complex. The participation of metal ions in the hydrolysis of carbon-nitrogen bonds is of considerable interest, and it seemed desirable to have available more kinetic data and more information on the influence of representative- and transition-metal ions in order to test the proposed mechanisms of hydrolysis. A quadridentate ligand, bisacetylacetoneethylenediimine, was selected for study because the rates of hydrolysis for this and related ligands, as well as the copper(II) derivatives have been determined [3]. *Present address: Endo Laboratories, Garden City, N e w York. 1. G. L. Eichhorn, andJ. C. Bailar, Jr.,J.Am. chem. Soc. 75, 2905 (1954). 2. G. L. Eichhorn and N. D. Marchand, J. Am. chem. Soc. 78, 2688 (1956). 3. D. F. Martin and F. F. Cantwell, J. inorg, nucl. Chem. 26, 2219 (1964). 1931
1932
D . F . M A R T I N and F. F. C A N T W E L L
CH3COCHsC(CHa)NCHsCHsN(CHa)CCHsCOCH3 + 2HsO ---> 2 C H 3 C O C H s C O C H 3 + HsNCHzCHsNHs.
(1)
It would also be desirable to know what relationship exists between the chelating tendency of a ligand with a given metal ion and the ease of hydrolysis. Unfortunately, few stability or kinetic data are available. The chelating tendencies of various bidentate/3-ketoimines with some divalent metal ions have been determined[4] as have the rates of hydrolysis of certain quadridentate fl-ketoimines and their copper(II) derivatives [3]. It seemed desirable to determine the rates of hydrolysis of a standard quadridentate fl-ketoimine in the presence and absence of metal ions and to determine the effect, if any, of expected stability trends on hydrolytic tendencies. EXPERIMENTAL Materials
Analytically pure samples of the ligand were used as before[3]. Standard solutions of metal perchlorates were used, except for beryllium nitrate. Determination o f rates o f hydrolysis Specific rate constants for the (pseudo) first-order hydrolysis ofligand in the presence of metal ions were determined titrimetrically, using the previously described [3] procedure. As before, rates were determined in 50 per cent (v/v) dioxane at 30 °. The initial pH was adjusted by addition of standard perchloric acid (or tetramethylammonium hydroxide for the copper-ligand systems). Pertinent data are summarized in Tables 1 and 2. In general, the rate constants were reproducible within _+4 per cent.
RESULTS
AND DISCUSSION
In dioxane-water solution, the ligand H2Ch may participate in several equilibria HsCh + H + ~ HsChH +
K1
(2)
HaCh++ H + ~ HsChH~ +
Ks
(3)
HsCh ~ H C h - + H +
KD1
(4)
HsCh- ~ Ch s- + H +
KD2.
(5)
And in the presence of divalent metal ions, additional equilibria may be expected including complex formation: M s+ + Ch s- ~ MCh
Ks.
(6)
It may be expected that metal species MCh will undergo protonation to form M C h H ÷ and MChH] + with protonation constants KI and K~' respectively. It may be presumed that all of the above species will undergo hydrolysis, though at significantly differing rates. On the basis of a previous study [3] it appears reasonable to suggest that the process of hydrolysis of the metal-chelate compound at pH 4.2 may be represented by the following equations: 4. D. F. Martin, G. A. Janusonis and B. B. Martin, J. A m. chem. Soc. 83, 73 (1961).
Hydrolysis of a/3-ketoimine
~C-- O \ HC//
O--C / M/ ~ , %CH ~./C-- %,N/ ~N--C / /. -CH CH 8 + ~ z z
H,O"1"
K', --~=-~i---
1933
~C--O. O--C( HC// \M / "~CH ~.C_N /
/ Slow
\ / CH # C - - O \ M / O--C~-C H \C ~-_0/ ~ O _ C / / \ (¢) 4H2NCH2 CHz NH3'
/
(d) + + H3 N-CH2CH2 -NH t
O--C
/
(a)
C--O N /O--C / HC /M "~CH C / = ~1 \ O = C / NH3+\ CHz CH2 (b)
/ XC--O\ -0--C% HC// M + /CH • \C =0 /
k'i
\ H20 fast
\N_C
\ /H I CH2CHzH/O\H
\
C--O OH2 H+ slow HC// \M+/ \ C ' N / ~NH --5---/ \ / 2 k2
CH2CHz (e) 4CH3 COCHz COCH~
It may also be assumed from previous work[3] that the observed rate of hydrolysis will be less for the metal-chelate compound than for the ligand. At a given pH, the concentration of the various metal containing species will be determined by the relative concentrations of ligand and metal ion, the pH, and the co-ordinating tendency of the ligand for a given metal ion. In this study a moderate value of pH (ca. 4.2) was selected, the concentration of the ligand was held constant, and the concentration of metal ion varied. It may be expected that there will be two limiting cases, assuming the effect of metal ion is to form the metal-chelate compound which will be more resistant to hydrolysis. The first limiting case will be the rate of hydrolysis in the absence of metal ion. The second limiting case will be the rate of hydrolysis when there is a sufficient excess of metal to insure that all of the ligand is present as the metalchelate compound*. The rate of hydrolysis should then correspond to the rate of hydrolysis for the metal-chelate compound at the same pH. This is apparent, if it *This, of course, presupposes that two conditions are satisfied; the value of Kr is relatively large and the rate of complex formation is rapid. The first condition is no doubt satisfied, based on the values of metal derivatives of bidentate Schiff bases [4]. The characteristic colors of the copper(l I) and nickel(II) derivatives of the quadridentate Schiffbase are formed instantly on mixing.
1934
D . F . MARTIN and F. F. C A N T W E L L
is assumed that the observed rate is equal to kobs.CH,ch , where CHzCh is the total concentration of quadridentate ligand, and if the species of major importance are H2Ch, H2Ch +, H4Ch 2+, MCh, MChH +, MChH22+, for which the specific hydrolyt I sis constants are kl, k2, ka, k~, k s, and k 3. It has been found that in the absence of metal ion, Equation (7) is valid: kobs. =
k, + k2K1 [H +] + k3K,K~ [H+] 2 1 + K , [ H +] + KIK2[H+] *
(7)
and that when essentially all of the ligand has been converted to metal-chelate species the rate data are best represented as: I t + t t t + 2 klP + k2KI[H ] + k3K1K2[H ] Robs. ~-" kob~.- 1+ K~ [H +] + K~K~ [H+] 2
(8)
The evaluation of the various constants was described previously [3]. In order to obtain a rate law to include all metal-ligand ratios, it would be necessary to evaluate Kj, which was not possible because of hydrolysis. Nor was it possible to evaluate KD~ and Kr~; again because of hydrolysis. However it was possible to study the second limiting cases of hydrolysis of ligand in the presence ofcopper(II) or nickel(II) ions. The pertinent data, the variation of the observed specific rate constant, kobs., as a function of the ratio metal ion-chelating agent are listed in Table 1 and plotted as Figs. 1 and 2. In general it appears that the variations ofkobs, vs. CM2+/CH~chare best interpreted as a smooth curve, rather than, say, as two linear portions. The single exception is the curve for the copper(tI)-ligand system, (Fig. t) for which two linear portions might be distinguished. This is exceptional because for this metal-ligand system alone it was necessary to add initially tetramethylammonium hydroxide to increase the pH to 4.2; all other systems required addition of acid. The values of kob~. for the hydrolysis of quadridentate figand in the presence of copper(II) or nickel(II) ions approach a minimum as the metal-ligand ratio approaches unity (Fig. 1). The minimum values of kob~. are approximately those observed for the pure metal-chelate compound. For example, the values of kob~.for the hydrolysis of the pure copper(II)- and nickel(II)-chelate compounds are 1.5 x 10-5 (calc. for pH, 4.16), and 1.91 x 10-4 (pH, 4.16) sec -1, respectively. The minimum values of k'ob~, at pH 4.16 in the presence of copper(II) and nickel(II) ions are 0.5 × 10-5 and 1.85 x 10-4 s e c -1, respectively. Thus, there is fairly good agreement between the expected and observed limiting values for these two metal ions. On the basis of this information, it appears reasonable to postulate two limiting rates, and to postulate that the reduction in rate of hydrolysis is due to the formation of the metal-chelate compound, which hydrolyzes at a reduced rate. This is not to imply that the formation constants for the copper(II) and nickel(II) complexes are the same. Quite obviously, they differ substantially as indicated by the differences in initial pH values which required adjusting with base to bring the initial pH to 4.2 in the case of some copper(II) systems. The ligand-beryllium and -magnesium systems are somewhat different, and require additional comment. The data for quadridentate ligand-beryllium(II) and -magnesium(II) systems suggest that beryllium co-ordination entities are formed, and it appears that limit-
H y d r o l y s i s of a fl-ketoimine
1935
Table 1. Specific rate c o n s t a n t s for the hydrolysis of CH3COCH2C(CH3)NCH2CHzN(CH:~)CCH2COCH3(HzCh) in the presence of divalent-metal ions in 50vol. per cent d i o x a n e - w a t e r at 30°C and pH 4.16
CH2Ch Metal ion
103 M
C M2~/CH2ch
104kob...... -~
Copper(lI)
3.99 3.99 3-96 3-98 4-00 3-93 3.89
0.079 0.264 0.529 0.790 1.05 1.06 1.32
2.57 1.90 1.30 0-672 0.066 0.138 (?) 0.046
Nickel(ll)
3-98 3-98 3"98 3"97 3-98 3.97 3 -99 3.98 3"98 3.98
0.053 0.103 0.264 0'264 0.422 0.530 0.727 0.829 0-935 1-19
2-87 2-45 2-41 2-64 2.46 2-36 2-04 1-90 1.94 1.78
Beryllium
3-98 3-98 3 98 3.96 3"98 3-85 3.98 3.99 3'99 3.99 3 '99 3'99 3"85 3'86 4-00 3'97
0-050 0-050 0.131 0-131 0.251 0-311 0.414 0.498 0-504 0.640 0-644 0.828 0.946 1-20 1.30 1"31
2'40 2.80 2.12 2.14 1-98 1.85 1-81 1.55 1.60 1.52 1.69 1.35 1.34 1.13 1.12 1.19
Magnesium
3 -98 3.98 3"98 3-98 3-98 3.98 3.97 3"98 3.97
0-075 0.151 0.226 0-251 0.405 0-603 0.989 1.03 1.21
1-96 2.62 2-43 2.38 2.06 1.92 1.82 1-72 1-70
1936
D . F . M A R T I N and F. F. C A N T W E L L
i
3"0
T
u
Ni t +
0
Cu t'+
•
2-0
I&l O)
0
O ~ o
o
I'0
o
I 0
0.2
I
I
0.4
0-6
0-8
I'0
1-2
CM / C H=Ch Fig. 1. Variation of specific rate constant (104kob,., sec -1) for hydrolysis with nickel (open circles) and copper ions (closed circles).
CM2+/CHichfor
ing rates are approached when the metal ion-chelating agent ratio is 1.0. Moreover, the limiting values of kobs. are of the expected order, being 1.33 × 10-4 sec -1 and 1.80 x 10-4 sec -x, for the beryllium and magnesium systems, respectively. The beryllium and magnesium derivatives of the quadridentate ligand have not been isolated. And until recently there was good reason to doubt that berylliumnitrogen co-ordination occurred in aqueous solution. The lack of evident coordination of beryllium ion with such strong chelating agents as picolinic acid, glycine, iminodiacetic acid, and fl-alanine has been noted[5]. On the other hand, formation constants were obtained for beryllium(I I) derivatives of Schiff bases of the type RCOCH2C(NR')R" in 50 vol. per cent dioxane-water; for these ligandberyllium systems, no evidence of hydrolysis was detected[4]. Still more recently, beryllium derivatives of the Schiff bases o-OC6HaCH(--NC4Hg)- and (o-OC6H4CH--NCH2CH2N~-----CHC6H40-o) ~- have been prepared[7]. Though these results have been repeated, attempts to isolate bisacetylacetoneethylenediiminoberyllium(lI) have not been successful. Nevertheless, there is evidence that the effectiveness of beryllium in reducing the rate of hydrolysis is due to an effect of complex formation and not just to an 5. D. Sen, Unpublished results [6]. 6. R. M. Izatt, W. C. Fernelius, and B. P. Block, J. phys. Chem. 59, 80 (1955). 7. R.W. Green and P. W. Alexander, Aust. J. Chem. 18, 1297 (1965).
Hydrolysis of a fl-ketoimine
1937
I
o
M~ 2"f 0
3.0
5
Be=+ •
2.0
(.3 i,i (/)
•
0
"
O--
"o 1.0
o 0
I
I
0-2
0.4
I
I
0 "6
0-8
1 1.0
I I-2
CM /CHINCH
Fig. 2. Variation of specific rate constant (104kobs., sec -~) for hydrolysis with CM~+/Crbch for magnesium (open circles) and beryllium ions (closed circles).
effect of ionic strength. This evidence is the pH dependence of the ligand hydrolysis, which is given by the data in Table 2 and Fig. 3 ; the pH dependence is clearly different in the presence of beryllium ion (CBe2+/CH2ch' c a . 1) (Fig. 3). At a low pH, c a . 2.0, there is little difference in rate of hydrolysis when beryllium ion is present or absent; at a pH of 4 the rate has been reduced markedly. This is attributable to nearly complete absence of complex formation at the low pH and nearly complete formation of a beryllium complex at the higher pH. From the data represented in Fig. 2, it would appear that the 1 : 1 species is the most likely'present when the limiting rate is observed. However, with the absence of formation-constant data and hydrolysis rates of the pure metal-chelate compound, it is not possible to be certain of the identity of the 1 : 1 species. This is true of beryUium; for that matter, it is true of the copper and nickel systems for which the hydrolysis data are available. With the latter two systems, it is possible only say the hydrolyzing species are the same. In this connection, it is both interesting and instructive to examine the ratio kobs./kobs., which is the ratio of the limiting rate constants in the presence and absence of divalent metal ions, respectively. The values are 0.017 (CuZ+), 0.46 (BeZ+), 0.65 (Ni 2+) and 0-61 (Mg2+). It seems reasonable to grant that the value of k'obs, approximates the constant for the hydrolysis of the pure metal-chelate compound at the pH selected. It might be tempting to seek a correlation between
1938
D . F . M A R T I N and F. F. C A N T W E L L Table 2. Specific rate constants for the hydrolysis of H2Ch, effect of pH*
Metal ion
C.,ch 103 M
Cu~+ Ca,ch
pH
Copper(II)
4.00
1.01
2"39
1 "98
1"02
3'35
1"64
3"92
1"30
3"36
1"65
Beryllium
3.99 3.99 3.97 3.97
1.16 1.16 1.18 0.90
2.22 2.41 2.58 3.36
Nickel(II)
3.94
1.04
3.36
5.58
Magnesium
3.96 3.96
0.80 0.82
3-36 3-36
5.89 6.10
104kobs.,see-' 13'2
21.0 18.3 10.8 2.7
*t = 30°; Solvent, 50 vol. per cent dioxane-water.
oo e ~ \ 3-0
_
•
O0
o o v
o
0
-J 3 . 5 _ i
e
4.0
, 2-0
i
I
i
I
i
3.0
i
i
,
I
~'
4"0 pH
Fig. 3. Variation of specific rate constant for hydrolysis ( - l o g kobs.) of H2Ch as function of pH calculated (--), observed in absence of metal ions (open circles) and in the presence of beryllium ion (CRez+/Cmc h c a . 1-0; closed circles).
Hydrolysis of a/3-ketoimine
1939
these ratios and a stability order, though such a correlation does not seem valid in this case; the positions of nickel and magnesium are clearly reversed. It might be suggested that the slope of the curved portions, or ligand-excess regions, of Figs. 1 and 2 at certain values of CM2+/CH2ch could be correlated with stability inasmuch as the slope would be a measure of the amount of metal-chelate compound being formed. On the other hand, at the limiting hydrolysis rates, related to k'obs., the concentration of complex should be substantially the same, and k'ob~. should be insensitive to the formation constant, Kz. This is true if all metalchelate compounds studied have relatively large values of Ke. It seems reasonable to ascribe the differences in values of k'obs./kobs, to two sources: (1) the basicities of the ligand vs. metal-chelate compound, largely as reflected in the values K1 and K~; and (2) the ease of breaking the protonated carbon-nitrogen bonds, largely as reflected in the values k2 vs. k,~.Two synergistic effects may be expected: the act of co-ordination should cause a greater polarization of the carbon-nitrogen bond and this should be reflected in large values of K1 [8]. On the other hand the protonated carbon-nitrogen bond is less susceptible to cleavage when the nitrogen is attached to a metal atom because an additional bond must now be broken to permit proton abstraction from solvent or hydronium ion to form the kinetically determined products, an amine and metal-/3diketone complex; this would be reflected in reduced values of k'l and k~ relative to k~ and k2. Granting the importance of the polarizing action of the metal, a correlation might be expected between kobs./kobs, and XM, the electronegativity of the metal. Using electronegativity data of Allred[9], the correlation is fair and the anomalous position of nickel is again noted. Now it becomes apparent that a more detailed analysis of the value of k~bs.for the nickel(I I) compound is in order. The significant differences between the nickel(lI) and copper(ll) derivatives appear to be in the value of k~ which is greater for the nickel(I I) derivative by a factor of 25 [ 10]. The values of k.; are comparable. Thus the anomaly in kobs. seems to be due to differences in k I. This may be due to the greater predilection of the nickel(II) derivatives to form five- and six-co-ordinate entities and the consequence that a Grotthus proton-transfer chain ends nearer the co-ordinated nitrogen. H
H
H
I
I
\/
H--O--H ... O--H ... O
H ~
H
H
I
L
+
H - - - O . . . H - - - O . . . H----O--H
Obviously, it would be desirable to test this suggestion by examining other nickel(II)-Schiffbase complexes. Unfortunately, most of the pertinent compounds, 8. For the quadridentate ligand used in this study the values are: 103kl, 2.40 sec -1 and 103k.,, 8.25 sec-'; and for the copper(II) derivative; 103k,, nil; 103k2, 0.269 sec-L 9. A. L. Allred, J. inorg, nucl. Chem. 5,264, 269 (1958). 10. E.J. Olszewski, Thesis, University of Illinois (1964).
1940
D . F . M A R T I N and F. F. C A N T W E L L
prepared in connection with other studies [11], are insoluble in 50 or even 75 vol. per cent dioxane-water media. Apparently, the synthesis of more soluble compounds must be achieved before further studies are possible. Acknowledgment-This research was supported by P. H. S. Research Grants 07873 and 01204, Division of General Medical Sciences, U.S. Public Health Service. 11. E.J. Olszewski and D. F. Martin, J. inorg, nucl. Chem. 26, 1577 (1964).