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Complex formation by N-substituted acetoacetamides—III
J. inorg,nucl.Chem., 1969,Vol. 31, pp, 2477to 2483. PergamonPress. Printedin Great Britain
COMPLEX
FORMATION BY ACETOACETAMIDES
N-SUBSTITUTED - III*
S T A B I L I T I E S OF D I V A L E N T M E T A L C H E L A T E S OF A C E T O A C E T A N I LI DES H. J. H A R R I E S , S. S A V A G E and G. W R I G H T Department of Chemistry, Derby and District College of Technology, Derby, England and N. L O G A N Department of Chemistry, University of Nottingham, Nottingham, England
(Received 23 October 1968) Abstract-Acetoacetanilide and its derivatives have been examined as chelating agents for divalent metal ions in 50%(v/v) and 75%(v/v) aqueous dioxan. Acid dissociation constants were determined potentiometricaUy and values related to electronic effects of substituents. Stability constants of the Be(I1), Co(II), Ni(II) and Cu(II) chelates were related to acid dissociation constants and to the metal. EXPERIMENTAL
Materials. Chelating agents (BDH, Ltd.) were used without purification except (V) which was recrystallised from ethanol. Chelating agents CHa. C O . CH2. C O . N H . Ar
No I II 1II IV V
Ar
M.P. (°C)
Lit. (°C)
Ref.
phenyl 2 -methoxyphenyl 2-methylphenyl 2-chlorophenyl 2,5-dichlorophenyl
83.5-84 85 105 104.5-105 94.5-95
85 85 -86 104-106 105 94-96
[3a] [3 b] [3c] [3d] [3e]
Metal salts (BDH, Ltd.) were nitrates or suiphates, AR where available. Metal content was determined titrimetricaUy using EDTA, except Be(II) which was determined by cation exchange resin. *Previous papers in this series have reported the determination of stabilities of the Be(II) chelates of acetoacetanilide and its derivatives [ 1] and of divalent metal chelates of N-(2-pyridyl)-acetoacetamide [2]. More refined technique and calculation have been used here to obtain improved values for acid dissociation constants and Be(II) chelate stabilities. This has been extended to a study of the chelates of other divalent metals. The low water solubilities of the reagents and their metal chelates necessitate the use of mixed solvents. 1. H.J. Harries,J. inorg, nucl. Chem. 25, 519 (1963). 2. H.J. Harries, J. inorg, nucl. Chem. 29, 2484 (1967). 3. (a) E. H. Rodd, Chemistry of Carbon Compounds Vol. Ilia, p. 190 (1954). b,c,d,e. Beilstein, Handbuch der Organischen Chemie Vol. 12(I), p. 386; Vol. 12(1I) pp. 319, 337; Vol. 13(1I) p. 181. 2477
2478
H; J. H A R R I E S etal.
Dioxan was p,urified !as recommended by Weissberger:and Proskauer[4]. Buffer solutions for mixed solvents .were p r e p a r e d a s follows: 50/50 mixtures of 0-2 M sodium acetate/0:2 M acetic acid and 0.2 M ianiline/0.2 M aniline hydrochloride in 50% and 75% dioxan were prepared.~Their,pH values were calculated from pKa data of Conway [5] and activity data compiled by Harned and Owen[6]. pH Measurements. These were made using either the ElL. Vibron pH:Meter(39A) or the EILVibron Electrometer(33B) with pH-Measuring Unit(C-33-B). All measurements were made at 25 -+ 0.05°C. Electrodes. EIL glass electrodes (GG33) and (GHS33) were used as hydrogen-indicator electrodes, the latter for the high pH-values involved in the determination of pKa. Saturated calomel electrodes (EIL RJ23) were prepared using 50% and 75% dioxan. pKa determinations. These were made by adding, in ten equal aliquots, one equivalent of KOH to a solution of the chelating agent and measuring pH after each addition, pKa was calculated for each measurement:
pK a = pH + log10 [HC] + [OH] ~-A. [C] - [OH]
HC and C represent the neutral chelating agent and its singly charged anion. A, the activity correction, was calculated from the same data as used in calculation of buffer pH [6]. Stability constants. All estimations were made using solutions containing 0.1 MNaCIO4(75% dioxan) or 0.1 MKC104(50% dioxan) to provide media of constant ionic strength. Solutions containing perchloric acid (where necessary), sodium or potassium perchlorate, metal(lI) salt and chelating agent were titrated with KOH. The degree of complex formation, ~, and the concentration of free chelating anion [C-] were calculated as described previously [ 1]; n--- ([MC +] + 2[MC~])/Cm where Cm is the total metal concentration. Plots of~ vs.pC (= - log[C]) were used to obtain approximate values of log K~ and log Kz as t h e p C values where ~ = ½ and 1½ respectively [7]. KI and K2, the first and second step-wise stability constants, are defined as the equilibrium constants for the two reactions M 2 ÷ + C - ~ MC +. and MC + + C - ~ MCv Accurate values of stability constants were obtained graphically by plotting n / ( h - 1)[C] vs. (2 - a ) [ C ] / ( a - 1); this is based on the equation [8] ,~ (n-- I)[C]
(2 - ~)[C] ( n - 1) "K1K~-K1.
K1 and Ks were obtained from the slope and intercept. The best straight line for the above plot (which will, for convenience, be written Y = K~KzX- K0 was obtained by the method of least squares. Since Y, X are sensitive to errors in ~ when ~ is near 0, l(for Y) and 1,2(for X), X and Y values for near 0, 1,2, were not used in the least squares calculation. This is illustrated by data in Table 2. The uncertainty in K~ and K~ for a given determination was obtained by calculating the best straight line assuming (a) errors in X only and (b) errors in Y only. This gives an upper and a lower limit to K1 and Ks; the values quoted in Table 2 are the mean of these two -+ half the spread between the upper and lower estimates. After K~ and Kz were obtained as above, [C] was calculated for h in the range 0.1 to 1.9 using the above equation and compared with the experimental formation curve (Fig. 2.). The good agreement 4. A. Weissberger and E. S. Proskauer, Technique of Organic Chemistry, Vol. 7, p. 371. Interscience, New York (1955). 5. B. E. Conway, ElectrochemicalData, P. 196. Elsevier, New York (1952). 6. H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolytic Solutions, p. 548. Reinhold, New York (1950). 7. J. Bjerrum, MetalAmmine Formation inAqueous Solution, p. 35. Haase, Copenhagen (1941). 8. F . J . C . Rossotti and H. Rossotti, The Determination of Stability Constants, p. 91. McGraw-Hill, New York (1961)~
Complex formation by N-substituted acetoacetamides
2479
over the whole range of ~ values indicates that polynuclear species and hydroxy-complexes are not significant in this system. Stability constants, determined as above, were reproducible to 0.05 log units or less in replicate experiments, while variation of starting concentration of metal and chelating agent gave results whose variation was ± 0.10 log units or less. RESULTS AND DISCUSSION
Acid dissociation constants are presented in Table 1. In each determination, 5 ml 0-1 M K O H were added, in 0.5 ml portions, to 50 ml of 0-01 M chelating agent. These values show up to 0.2 log unit variations with the more approximate values reported earlier [ 1]. In both solvents, pKa decreases monotonically from (I) to (V). The fact that (I) is at variance with the known electronic effects of the substituents may be attributed, at least in part, to the ability of (I) alone to have the phenyl group coplanar with the chelate ring of the H-bonded enol form of the acid. This gives an added stability to the neutral acid and reduces its proton acidity. This type of relation is illustrated graphically in Fig. 1, in which pK, is plotted against Hammett o--constants [9], the latter being taken from data for o-substituted benzoic acids [10]. Stability constants. The calculation procedure is illustrated in Table 2. Stability constants are presented in Table 3. The low solubility of the metal chelates caused precipitation of solids at an early stage during many of the titrations before useful values of h were obtained, Thus, it was not possible to obtain any log K values for Mn(II), Fe(I1), or Zn(ll) and log K1 only was obtained for Co(II), Ni(II) and Cu(II). In Fig. 3 are plotted relationships between log K and pK, in 50% dioxan: similar graphs are obtained for the results in 75% dioxan. The curves are linear with slope equal or near to unity. The stability of the metal chelate is thus a function of the base strength of the donor atoms and the steric effect of the ring substituent in (I1) to (V) is not significant. This behaviour parallels that shown by many structurally similar chelating agents [I 2], including several/3-diketones [13,141. Table 1. Acid dissociation constants of reagents I-V Reagent 1 11 II1 IV V
9. 10. 11. 12. 13. 14.
pKa (50% Dioxan)
pK,, (75% Dioxan)
11-43 _ 0.01 11.33 ± 0.01 11.16±0.02 10-63 ± 0.02 9.97 ___0.02
12.76 ± 0.02 12.73 ___0-01 12.41 ±0.02 11.88 ± 0.02 11-34 ± 0-02
0.00 0.12 0-29 1.28 1.65
L. P. Hammett, Physical Organic Chemistry, Chap. 7. McGraw-Hill, New York (1940). G. B. Barlin and D. D. Perrin, Q. Rev. 20, 75 (1966). H. M. N. H. Irving and H. S. Rossotti, J. chem. Soc. 1953, 3397. H. M. N. H. Irving and H. S. Rossotti, J. chem. Soc. 1954, 2904, 2910. M. Calvin and K. W. Wilson, J. Am. chem. Soc. 67, 2003 (1945). L. G. Van Uitert, W. C. Fernelius and B. E. Douglas, J. Am. chem. Soc. 75, 457, 2736, 3862 (1953).
2480
H . J . H A R R I E S et al.
'•
O
O
75% Dioxane e
12
• "'~
Q ~.
II
o.
10 I
0
I I
0'5
/ 2
I 1.5
cr I
If
III
PV
V
L
[
I
I
I
Fig. 1. Relationship between acidity (pK.) and Hammett substituent constants (o-) for the reagents I-V. 2-0 ~X oX
X
By cok:uloflon
®
Experimentol
1.5
I~"
I-O
~5
q
%
0-5
o
s
I r
I I!t
! 9
-log [C]
Fig. 2. Formation curves for the Be(II) chelates of acetoacetanilide.
I IO
Complex formation by N-substituted acetoacetamides
2481
Table 2. Determination of stability constants for Be(ll) chelates of acetoacetanilide at 25°C in 50% Dioxan. Initial concentrations: total metal 0.001 M, total chelating agent 0.005 M, potassium perchlorate 0-100 M. Titrant: 0.100 M potassium hydroxide in 50% dioxan
T
PH
0.04 0"08 0.12 0-16 0"20 0'24 0-28 0"32 0-36 0"40 0"44 0"48 0'52 0'56 0"60 0"64 0"68 0.72 0.76 0.80 0.84 0.88 0-92 0-96 1.00
4.442 4.600 4.700 4.810 4.912 5.034 5' 152 5'271 5-404 5'537 5-689 5-848 6.023 6" 158 6.248 6,340 6.446 6'558 6.717 6.898 7-110 7.380 7.768 8.500 9"347
Anilide-BE-25 EN
C 0.5063 E-09 0-7176 E09 0.8887E-09 0-1125 E-08 0"1399 E-08 0" 1820 E-08 0-2345 E-08 0"3028 E-08 0"4036 E-08 0"5376 E-08 0"7480 E-08 0-1057 E-07 0.1549 E-07 0'2069 E-07 0'2492 E-07 0"3012 E-07 0'3760 E-07 0.4755 E-07 0"6699 E-07 0"9921 E-07 0-1577 E-06 0"2864 E-06 0"6820 E-06 0"3583 E-05 0"2452 E-04
0.1161 0.1851 0.2599 0.3355 0.4122 0-4892 0.5670 0.6453 0'7239 0"8029 0.8820 0.9614 1'0409 1.1206 1.2005 1'2804 1'3603 1.4402 1'5201 1.6000 1'6799 1-7597 1"8393 1-9163 1-9749
X
Y
-0.107923 E-08 - 0.159839 E-08 -0.208970E-08 -- 0"281998 E-08 -0-378008 E-08 - 0-538519 E-08 -0-776445 E-08 - 0' 115683 E-07 - 0" 186576 E-07 -0'326587 E-07 - 0"709007 E-07 -- 0'284580 E-06 0"362909E-06 0" 150817 E-06 0"993451E-07 0.773105 E-07 0"667554E-07 0'604732 E-07 0"618088E-07 0'661390 E-07 0"742552E-07 0'905696 E-07 0.130563 E-06 0'327132 E-06 0"629075E-06
-0.259585 E-09 - 0.316630 E-09 - 0.395342 E-09 - 0.448564 E-09 --0"501365 E-09 - 0"526254 E-09 -0"558405 E-09 - 0.600980 E-09 - 0.649841 E-09 - 0-757736 E-09 - 0.999814 E-09 -0'235757 E-10 0' 164106 E-10 0.448637 E-09 0"240213E-09 0' 151542 E-09 0-100398 E-09 0-687883E-08 0'436252 E-08 0"268755E-08 0' 156642 E-08 0.808702 E-07 0'321316 E-07 0"583569E-06 0.826068 E-05
T-volume(ml) KOH added; E N - h ; X,Y values bracketed are those used in least squares plot. Calculation of K~ and I42 by method of least squares.
Slope (= KIKz) Log K1K2 Intercept ( = - - K 0 Log K~ Log K2 (by difference)
XonY
YonX
0.8169 × 1016 15"912 -0.484 × 109 8.685 7.227
0'8248 × 10TM 15"916 +0.5893 × 10-r (= l/K2) 7.229 (= log K2) 8.687 (= log K0
Mean values: Log K1 = 8.686 ___0.001; log K2 = 7.228 ---0.001.
The values of log(Kt/Kz) for the Be(II) chelates are:
K, l°g~2
I
11
ill
IV
V
50% dioxan
1.63
1.66
1.72
1.56
1.50
75% dioxan
1.70
1.47
1.41
1.32
1.08
H . J . H A R R I E S etal.
2482
These are well within the expected range and the absence of high values implies there is little or no hindrance to the addition of the second chelate group[15] and the o-substituents in (II) to (V) do not increase this. This is to be expected for tetrahedral complexes of Be(II) where the two chelate rings are at right Table 3. Stability constants of metal chelates
Reagent I II II1 IV V
Be(II) log Kl log Ke 8.69 8.58 8.48 7.81 7.17
7.23 7.05 6.88 6.37 5.88
Co(ll) log KI 4-49 4.30 4.20 3.51 3.15
Dioxan (50%) Ni(II) Cu(II) log K~ log K~ 4.82 4.62 4.48 4.17 3.67
Dioxan (75%) Be(II) Cu(II) Log KI Log KI Log K2
7.97 7.75 7.58 7.09 6.44
10.79 10.65 10.48 9.77 9.16
9"29 9.18 9.07 8.45 8.08
9.30 9.23 9.10 8.56 8.09
_o o
I0
10-5
II
11.5
pK, I
!
~r
a
~r
~
m
I
I
if1
K for Be (1~)
G
0
K t ,'
Cu (H)
m
a
Kz"
Be(~]
A
K l ,,
Ni (17)
0
K, ,,
Co (n)
Fig. 3. Relationships between stability constants of the metal chelates (log10K) and acidity (pK~) of reagents I-V. 15. F . J . C . Rossotti, Modern Coordination Chemistry (Editedby J. Lewis and R, G. "¢¢ilkins) pp. 3439. Interscience, New York (1960).
Complex formation by N-substituted acetoacetamides
2483
angles. The highest value in 75% dioxan of 1-70 for (I) may be due to stronger solvation of the latter as compared with (II)-(V) which contain large substituents. The stabilities of the Co(II), Ni(II) and Cu(ll) chelates agree with the IrvingWilliams order for divalent ion of the 3d-series [16], in which the order of stabilities for a given ligand is Mn < Fe < Co < Ni < Cu > Zn. 16. H. Irving and R. J. P. Williams, J. chem. Soc. 1953, 3192.
COMPLEX
FORMATION BY ACETOACETAMIDES
N-SUBSTITUTED - III*
S T A B I L I T I E S OF D I V A L E N T M E T A L C H E L A T E S OF A C E T O A C E T A N I LI DES H. J. H A R R I E S , S. S A V A G E and G. W R I G H T Department of Chemistry, Derby and District College of Technology, Derby, England and N. L O G A N Department of Chemistry, University of Nottingham, Nottingham, England
(Received 23 October 1968) Abstract-Acetoacetanilide and its derivatives have been examined as chelating agents for divalent metal ions in 50%(v/v) and 75%(v/v) aqueous dioxan. Acid dissociation constants were determined potentiometricaUy and values related to electronic effects of substituents. Stability constants of the Be(I1), Co(II), Ni(II) and Cu(II) chelates were related to acid dissociation constants and to the metal. EXPERIMENTAL
Materials. Chelating agents (BDH, Ltd.) were used without purification except (V) which was recrystallised from ethanol. Chelating agents CHa. C O . CH2. C O . N H . Ar
No I II 1II IV V
Ar
M.P. (°C)
Lit. (°C)
Ref.
phenyl 2 -methoxyphenyl 2-methylphenyl 2-chlorophenyl 2,5-dichlorophenyl
83.5-84 85 105 104.5-105 94.5-95
85 85 -86 104-106 105 94-96
[3a] [3 b] [3c] [3d] [3e]
Metal salts (BDH, Ltd.) were nitrates or suiphates, AR where available. Metal content was determined titrimetricaUy using EDTA, except Be(II) which was determined by cation exchange resin. *Previous papers in this series have reported the determination of stabilities of the Be(II) chelates of acetoacetanilide and its derivatives [ 1] and of divalent metal chelates of N-(2-pyridyl)-acetoacetamide [2]. More refined technique and calculation have been used here to obtain improved values for acid dissociation constants and Be(II) chelate stabilities. This has been extended to a study of the chelates of other divalent metals. The low water solubilities of the reagents and their metal chelates necessitate the use of mixed solvents. 1. H.J. Harries,J. inorg, nucl. Chem. 25, 519 (1963). 2. H.J. Harries, J. inorg, nucl. Chem. 29, 2484 (1967). 3. (a) E. H. Rodd, Chemistry of Carbon Compounds Vol. Ilia, p. 190 (1954). b,c,d,e. Beilstein, Handbuch der Organischen Chemie Vol. 12(I), p. 386; Vol. 12(1I) pp. 319, 337; Vol. 13(1I) p. 181. 2477
2478
H; J. H A R R I E S etal.
Dioxan was p,urified !as recommended by Weissberger:and Proskauer[4]. Buffer solutions for mixed solvents .were p r e p a r e d a s follows: 50/50 mixtures of 0-2 M sodium acetate/0:2 M acetic acid and 0.2 M ianiline/0.2 M aniline hydrochloride in 50% and 75% dioxan were prepared.~Their,pH values were calculated from pKa data of Conway [5] and activity data compiled by Harned and Owen[6]. pH Measurements. These were made using either the ElL. Vibron pH:Meter(39A) or the EILVibron Electrometer(33B) with pH-Measuring Unit(C-33-B). All measurements were made at 25 -+ 0.05°C. Electrodes. EIL glass electrodes (GG33) and (GHS33) were used as hydrogen-indicator electrodes, the latter for the high pH-values involved in the determination of pKa. Saturated calomel electrodes (EIL RJ23) were prepared using 50% and 75% dioxan. pKa determinations. These were made by adding, in ten equal aliquots, one equivalent of KOH to a solution of the chelating agent and measuring pH after each addition, pKa was calculated for each measurement:
pK a = pH + log10 [HC] + [OH] ~-A. [C] - [OH]
HC and C represent the neutral chelating agent and its singly charged anion. A, the activity correction, was calculated from the same data as used in calculation of buffer pH [6]. Stability constants. All estimations were made using solutions containing 0.1 MNaCIO4(75% dioxan) or 0.1 MKC104(50% dioxan) to provide media of constant ionic strength. Solutions containing perchloric acid (where necessary), sodium or potassium perchlorate, metal(lI) salt and chelating agent were titrated with KOH. The degree of complex formation, ~, and the concentration of free chelating anion [C-] were calculated as described previously [ 1]; n--- ([MC +] + 2[MC~])/Cm where Cm is the total metal concentration. Plots of~ vs.pC (= - log[C]) were used to obtain approximate values of log K~ and log Kz as t h e p C values where ~ = ½ and 1½ respectively [7]. KI and K2, the first and second step-wise stability constants, are defined as the equilibrium constants for the two reactions M 2 ÷ + C - ~ MC +. and MC + + C - ~ MCv Accurate values of stability constants were obtained graphically by plotting n / ( h - 1)[C] vs. (2 - a ) [ C ] / ( a - 1); this is based on the equation [8] ,~ (n-- I)[C]
(2 - ~)[C] ( n - 1) "K1K~-K1.
K1 and Ks were obtained from the slope and intercept. The best straight line for the above plot (which will, for convenience, be written Y = K~KzX- K0 was obtained by the method of least squares. Since Y, X are sensitive to errors in ~ when ~ is near 0, l(for Y) and 1,2(for X), X and Y values for near 0, 1,2, were not used in the least squares calculation. This is illustrated by data in Table 2. The uncertainty in K~ and K~ for a given determination was obtained by calculating the best straight line assuming (a) errors in X only and (b) errors in Y only. This gives an upper and a lower limit to K1 and Ks; the values quoted in Table 2 are the mean of these two -+ half the spread between the upper and lower estimates. After K~ and Kz were obtained as above, [C] was calculated for h in the range 0.1 to 1.9 using the above equation and compared with the experimental formation curve (Fig. 2.). The good agreement 4. A. Weissberger and E. S. Proskauer, Technique of Organic Chemistry, Vol. 7, p. 371. Interscience, New York (1955). 5. B. E. Conway, ElectrochemicalData, P. 196. Elsevier, New York (1952). 6. H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolytic Solutions, p. 548. Reinhold, New York (1950). 7. J. Bjerrum, MetalAmmine Formation inAqueous Solution, p. 35. Haase, Copenhagen (1941). 8. F . J . C . Rossotti and H. Rossotti, The Determination of Stability Constants, p. 91. McGraw-Hill, New York (1961)~
Complex formation by N-substituted acetoacetamides
2479
over the whole range of ~ values indicates that polynuclear species and hydroxy-complexes are not significant in this system. Stability constants, determined as above, were reproducible to 0.05 log units or less in replicate experiments, while variation of starting concentration of metal and chelating agent gave results whose variation was ± 0.10 log units or less. RESULTS AND DISCUSSION
Acid dissociation constants are presented in Table 1. In each determination, 5 ml 0-1 M K O H were added, in 0.5 ml portions, to 50 ml of 0-01 M chelating agent. These values show up to 0.2 log unit variations with the more approximate values reported earlier [ 1]. In both solvents, pKa decreases monotonically from (I) to (V). The fact that (I) is at variance with the known electronic effects of the substituents may be attributed, at least in part, to the ability of (I) alone to have the phenyl group coplanar with the chelate ring of the H-bonded enol form of the acid. This gives an added stability to the neutral acid and reduces its proton acidity. This type of relation is illustrated graphically in Fig. 1, in which pK, is plotted against Hammett o--constants [9], the latter being taken from data for o-substituted benzoic acids [10]. Stability constants. The calculation procedure is illustrated in Table 2. Stability constants are presented in Table 3. The low solubility of the metal chelates caused precipitation of solids at an early stage during many of the titrations before useful values of h were obtained, Thus, it was not possible to obtain any log K values for Mn(II), Fe(I1), or Zn(ll) and log K1 only was obtained for Co(II), Ni(II) and Cu(II). In Fig. 3 are plotted relationships between log K and pK, in 50% dioxan: similar graphs are obtained for the results in 75% dioxan. The curves are linear with slope equal or near to unity. The stability of the metal chelate is thus a function of the base strength of the donor atoms and the steric effect of the ring substituent in (I1) to (V) is not significant. This behaviour parallels that shown by many structurally similar chelating agents [I 2], including several/3-diketones [13,141. Table 1. Acid dissociation constants of reagents I-V Reagent 1 11 II1 IV V
9. 10. 11. 12. 13. 14.
pKa (50% Dioxan)
pK,, (75% Dioxan)
11-43 _ 0.01 11.33 ± 0.01 11.16±0.02 10-63 ± 0.02 9.97 ___0.02
12.76 ± 0.02 12.73 ___0-01 12.41 ±0.02 11.88 ± 0.02 11-34 ± 0-02
0.00 0.12 0-29 1.28 1.65
L. P. Hammett, Physical Organic Chemistry, Chap. 7. McGraw-Hill, New York (1940). G. B. Barlin and D. D. Perrin, Q. Rev. 20, 75 (1966). H. M. N. H. Irving and H. S. Rossotti, J. chem. Soc. 1953, 3397. H. M. N. H. Irving and H. S. Rossotti, J. chem. Soc. 1954, 2904, 2910. M. Calvin and K. W. Wilson, J. Am. chem. Soc. 67, 2003 (1945). L. G. Van Uitert, W. C. Fernelius and B. E. Douglas, J. Am. chem. Soc. 75, 457, 2736, 3862 (1953).
2480
H . J . H A R R I E S et al.
'•
O
O
75% Dioxane e
12
• "'~
Q ~.
II
o.
10 I
0
I I
0'5
/ 2
I 1.5
cr I
If
III
PV
V
L
[
I
I
I
Fig. 1. Relationship between acidity (pK.) and Hammett substituent constants (o-) for the reagents I-V. 2-0 ~X oX
X
By cok:uloflon
®
Experimentol
1.5
I~"
I-O
~5
q
%
0-5
o
s
I r
I I!t
! 9
-log [C]
Fig. 2. Formation curves for the Be(II) chelates of acetoacetanilide.
I IO
Complex formation by N-substituted acetoacetamides
2481
Table 2. Determination of stability constants for Be(ll) chelates of acetoacetanilide at 25°C in 50% Dioxan. Initial concentrations: total metal 0.001 M, total chelating agent 0.005 M, potassium perchlorate 0-100 M. Titrant: 0.100 M potassium hydroxide in 50% dioxan
T
PH
0.04 0"08 0.12 0-16 0"20 0'24 0-28 0"32 0-36 0"40 0"44 0"48 0'52 0'56 0"60 0"64 0"68 0.72 0.76 0.80 0.84 0.88 0-92 0-96 1.00
4.442 4.600 4.700 4.810 4.912 5.034 5' 152 5'271 5-404 5'537 5-689 5-848 6.023 6" 158 6.248 6,340 6.446 6'558 6.717 6.898 7-110 7.380 7.768 8.500 9"347
Anilide-BE-25 EN
C 0.5063 E-09 0-7176 E09 0.8887E-09 0-1125 E-08 0"1399 E-08 0" 1820 E-08 0-2345 E-08 0"3028 E-08 0"4036 E-08 0"5376 E-08 0"7480 E-08 0-1057 E-07 0.1549 E-07 0'2069 E-07 0'2492 E-07 0"3012 E-07 0'3760 E-07 0.4755 E-07 0"6699 E-07 0"9921 E-07 0-1577 E-06 0"2864 E-06 0"6820 E-06 0"3583 E-05 0"2452 E-04
0.1161 0.1851 0.2599 0.3355 0.4122 0-4892 0.5670 0.6453 0'7239 0"8029 0.8820 0.9614 1'0409 1.1206 1.2005 1'2804 1'3603 1.4402 1'5201 1.6000 1'6799 1-7597 1"8393 1-9163 1-9749
X
Y
-0.107923 E-08 - 0.159839 E-08 -0.208970E-08 -- 0"281998 E-08 -0-378008 E-08 - 0-538519 E-08 -0-776445 E-08 - 0' 115683 E-07 - 0" 186576 E-07 -0'326587 E-07 - 0"709007 E-07 -- 0'284580 E-06 0"362909E-06 0" 150817 E-06 0"993451E-07 0.773105 E-07 0"667554E-07 0'604732 E-07 0"618088E-07 0'661390 E-07 0"742552E-07 0'905696 E-07 0.130563 E-06 0'327132 E-06 0"629075E-06
-0.259585 E-09 - 0.316630 E-09 - 0.395342 E-09 - 0.448564 E-09 --0"501365 E-09 - 0"526254 E-09 -0"558405 E-09 - 0.600980 E-09 - 0.649841 E-09 - 0-757736 E-09 - 0.999814 E-09 -0'235757 E-10 0' 164106 E-10 0.448637 E-09 0"240213E-09 0' 151542 E-09 0-100398 E-09 0-687883E-08 0'436252 E-08 0"268755E-08 0' 156642 E-08 0.808702 E-07 0'321316 E-07 0"583569E-06 0.826068 E-05
T-volume(ml) KOH added; E N - h ; X,Y values bracketed are those used in least squares plot. Calculation of K~ and I42 by method of least squares.
Slope (= KIKz) Log K1K2 Intercept ( = - - K 0 Log K~ Log K2 (by difference)
XonY
YonX
0.8169 × 1016 15"912 -0.484 × 109 8.685 7.227
0'8248 × 10TM 15"916 +0.5893 × 10-r (= l/K2) 7.229 (= log K2) 8.687 (= log K0
Mean values: Log K1 = 8.686 ___0.001; log K2 = 7.228 ---0.001.
The values of log(Kt/Kz) for the Be(II) chelates are:
K, l°g~2
I
11
ill
IV
V
50% dioxan
1.63
1.66
1.72
1.56
1.50
75% dioxan
1.70
1.47
1.41
1.32
1.08
H . J . H A R R I E S etal.
2482
These are well within the expected range and the absence of high values implies there is little or no hindrance to the addition of the second chelate group[15] and the o-substituents in (II) to (V) do not increase this. This is to be expected for tetrahedral complexes of Be(II) where the two chelate rings are at right Table 3. Stability constants of metal chelates
Reagent I II II1 IV V
Be(II) log Kl log Ke 8.69 8.58 8.48 7.81 7.17
7.23 7.05 6.88 6.37 5.88
Co(ll) log KI 4-49 4.30 4.20 3.51 3.15
Dioxan (50%) Ni(II) Cu(II) log K~ log K~ 4.82 4.62 4.48 4.17 3.67
Dioxan (75%) Be(II) Cu(II) Log KI Log KI Log K2
7.97 7.75 7.58 7.09 6.44
10.79 10.65 10.48 9.77 9.16
9"29 9.18 9.07 8.45 8.08
9.30 9.23 9.10 8.56 8.09
_o o
I0
10-5
II
11.5
pK, I
!
~r
a
~r
~
m
I
I
if1
K for Be (1~)
G
0
K t ,'
Cu (H)
m
a
Kz"
Be(~]
A
K l ,,
Ni (17)
0
K, ,,
Co (n)
Fig. 3. Relationships between stability constants of the metal chelates (log10K) and acidity (pK~) of reagents I-V. 15. F . J . C . Rossotti, Modern Coordination Chemistry (Editedby J. Lewis and R, G. "¢¢ilkins) pp. 3439. Interscience, New York (1960).
Complex formation by N-substituted acetoacetamides
2483
angles. The highest value in 75% dioxan of 1-70 for (I) may be due to stronger solvation of the latter as compared with (II)-(V) which contain large substituents. The stabilities of the Co(II), Ni(II) and Cu(ll) chelates agree with the IrvingWilliams order for divalent ion of the 3d-series [16], in which the order of stabilities for a given ligand is Mn < Fe < Co < Ni < Cu > Zn. 16. H. Irving and R. J. P. Williams, J. chem. Soc. 1953, 3192.