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Some theoretical and practical problems in the use of organic reagents in chemical analysis—VII
J. Inorg. Nucl. Chem. 1966, Vol. 28, pp. 139 to 145. Pergamon Press Ltd. Printed in Northern Ireland
SOME THEORETICAL AND PRACTICAL PROBLEMS IN THE USE OF ORGANIC REAGENTS IN CHEMICAL ANALYSIS--VII INVESTIGATIONS OF THE TRANSITION METAL COMPLEXES OF PYRIDINE-2-ALDOXIME K. BURG~, I. EGYEI) and I. RUFF Department of Inorganic and Analytical Chemistry, L. E6tvSs University,Budapest (Received 18 March 1965; in revised form 22 May 1965)
Abltraet--The composition of the mononuclear pyridin-2-aldoximecomplexes of 3dt-3d x°transition metals formed in neutral solution were determined. The successivestability constants of the manganese(ID, cobalt 0I), nickel (II)and zinc (II) complexes were measured. It was proved by proton resonance measurements that in the iron ([I) and cobalt (IT)complexespyddine-2-aldoximeis a strong field ligand, and the nickel 01) complex is octahedral. Tm~ compositions and stabilities of the dioxime complexes of 3dS-3d 1° transition metals were determined in our previous works, (l'a) and the factors influencing their stabilities explored. The analogous transition metal complexes of o-phenanthroline and ~-0~'dipyddyl were studied by IRVrN6t3) and recently by A~FJ~GG. (4) The ligands mentioned are in some respects similar. The functional group reacting with the metal ion i s :
\
/
/ C ¸ [I N
C I[ N
"\
but while the donor nitrogens of the dioximes belong to the oxime group, those of the other two ligands are heterocyclic nitrogens. All of the ligands mentioned are strong field ligands in their iron (II) complexes. The iron (II)-complexes are diamagnetic. The ligands mentioned above have free ~r-orbitals capable of accepting electrons from the central metal atom (back co-ordination). The existence of this kind of bond was proved in the case of the o-phenantroline and 0c-~'dipyddyl complexes by infra-red spectrophotometdc investigations by Busch and BAILAR,tS) and in the case of dimethylglyoxime complexes by our own infra-red studies. (2) The fundamental difference between the dioximes and the other two ligands mentioned is that the dioximes form complexes with a metal:ligand ratio 1:2, and have square planar structures; while o-phenanthroline and ~-~'dipyddyl form (1) K. BURt~ERand I. R u ~ , Talanta 10, 329 (1963). (J) K. BuRO~, I. Rul~ and F. RUFF,J. Inorg. Nucl. Chem. 27, 179 (1965). (s) H. IRVINOand D. H. M~LOR,./. Chem. Soc. 5222 (1962). (*) G. A ~ e ~ u 3 o , Helo. Chim. Acta 46, 2397, 2813 (1963). (5) D. H. Busot and J. C. B ~ Jr., J. Amer. Chem. Soc. 78, 1137 (1956).
139
140
K. BURGER, I. EGYED and I. RUFF
complexes with a metal:ligand ratio 1:3 and have octahedral symmetry. The different behaviour o f the dioximes can be explained most p r o b a b l y by the f o r m a t i o n o f strong intramolecular hydrogen bridges which stabilize the 1:2 complexes. Pyridine-2-aldoxime, a reagent used in analytical chemistry, t6.7"sJ occupies an intermediate place between the two types o f ligands mentioned. O f the two d o n o r nitrogens o f this ligand one is an oxime nitrogen, while the other is a heterocyclic one. This ligand is able to f o r m p r o t o n a t e d complexes t9'1°~ of composition M e A ( H A ) and polynuclear complexes, cm besides the regular m o n o n u c l e a r complexes. At present we are interested only in the m o n o n u c l e a r complexes o f composition analogous to that of the complexes mentioned above. The stability constants o f the copper (II) t g ' ~ and iron (II) (1~'13) pyddine-2-aldoxime complexes are given in the literature, some nickel (II) complexes with different composition have been prepared t14) and studied by infra-red spectrophotometry.t~5~ The acid dissociation constants of some p r o t o n a t e d complexes are also k n o w n P °'t2) The manganese (II), cobalt (II) and zinc (II) complexes have not yet been studied, and the stability constants o f the nickel (II) complex are unknown. Thus first o f all we had to study the mononuclear, regular (non-protonated) complexes o f manganese (II), cobalt (II), nickel (II) and zinc (II). EXPERIMENTAL Spectrophotometry. Our investigations in the ultra-violet and visible region were carried out on a Unicam SP 700 recording spectrophotometer. The spectrophotometric measurements were made in solutions of 10-~-5 × 10-~ mole/1. The molar extinction coefficients are shown in Fig. 1. The solutions of the metals and that of the ligand used for the measurements based on Job's method had a concentration of 2 × 10-' mole/l. For the determination of the successive stability constants of the cobalt II complex the degree of formation (a) of the CoAs complex was calculated from the extinction values measured at 30,000 wave numbers in solutions containing different metal and ligand concentrations at different pH values. The degree of formation of the complex (a) was plotted as a function of the logarithm of the free ligand concentration ([A-l) (Fig. 2). These (a -- log [A-]) points were fitted to a family of Y -- (log v)p normalized curves, where Y=
1 + pv + v~
v = (K1. K2)½ [A-]
and
p
=
K~J. K~-i
The successive stability constants (K1 and K0 were calculated from the value ofp of the best fitting curve and from the value of log [A-] at the point where log v = 0. The acid dissociation constants of the ligand determined from spectrophotometric measurements and used for the calculations of the free ligand concentration were: K1 = --3.4; logKa = --10,0 (where K1 was the acid dissociation constant of the heterocyclic NI-I+ group, and K~ that of the ~6~H. HARTK~a', Naturwissenschaften 45, 211 (1958). ~7~ H. H_~rgAw&, Z. Analyt. Chem. 170, 399 (1959). ~s~H. I - h R - - , Z. Analyt. Chem. 176, 185 (1960). cg~ C. H. LILTand C. FAN Lro, J. Amer. Chem. Soc. 83, 4167, 4163 (1961). clo~I. MAStmA and K. Srn~A, Science Reports (Osaka Univ.) 11, 3 (1962). c ~ B. KmsoN, Bull. Soc. Chim. France 1030 (1962). c~2~G. I. H. HArqAr,rm and D. H. I R ~ , d. Chem. Soc. 2745 (1962). ~8~ y. l S ~ L I and E. JtmoREIs, Bull. Res. Council oflsrael, 11 A, 121 (1962). ~1~ R. A. KRAtrSEand D. H. BUSCH,d. Amer. Chem. See. 82, 4830 (1960). t~s~R. A. K~usE, N. B. CoLTrxtre and D. H. Btrscrl, d. Phys. Chem. 65, 2216 (1961).
Problems in the use of organic reagents in chemical analysis---VII
141
HA 1'6-
I-2-
0"8
0"4 O0 1"6 1'2 0-8
L
(
A-
Ni
Mn
Cu
Fe
Zn
I
\
0'4 0"0 E
1"6 1'2 0"8 0"4 0"0 1'6 1"2 0"8
I
0"0,50 ] 42
[
I
34
[
i
26
i
50
42
34
26
kK
FIe. 1.--The absorption spectra of pyridine-2-aldoxime and its manganese (l]), iron (N), cobalt (II), nickel (11), copper 013 and zinc 0I) complexes. oxime group). The compositions of the solutions were: cobalt concentration 2 × 10-5-4 x 10--s mole/L, ligand concentration 10-5-2 x 1e-a; pH = 6.5--8.0; NaC10, concentration = 0"3 mole/1. For the spectrophotometric investigations of the complexes solutions of the metal and the ligand of suitable conczntration were mixed; quantitative formation of the complexes was a~ured by choicz of a suitable pH. The cobalt (11), copper (ll) and zinc (H) complexes were extracted by chloroform
142
K. BvgG~t, I. EoY]~a and I. Ru]~
I'O
0.5
o/ log ~',=8'6 log ~z= 17'2
0
I 9
i
i 8
7
- l o g EA:]
Fro. 2.--The formation degree of the cobalt (H)-pyridine-2-aldoxime complexes of CoAs composition plotted as a function of the logarithm of free ligand concentration. ( Q [ ] A mark the points calculated from the data of our measurements; the curve shown represents the best fitting normalized curve.) Cobalt concentration Q 2 x 10-~ M, [] and A 4 × 10-~ M. Pyridine-2-aldoxime concentration 10-5-2 × 10-a M. The p H of the solution O 6.5, [] 7.2, A 8.0. from the aqueous solutions.
The solid complexes were prepared by evaporating the extract in
V/IC//O.
Theptl-metricequilibriummeasurements. These measurements were carried out by BS~RRUM'SCle'17~ method. The formation curves were evaluated by the method of DxqksSSN and SILL'S,{18j based on the use of normalized curves. The points of the (t~ -- log [A-]) formation curve, {~e~calculated from experimental data were fitted to a family of Y -- 0og v)~ normalized curves, where Y=
v + 2v~p 1 + v + _pvs
v = Ka[A-]
and K,
The successivestabilityeomtants (Kt, K O were calculated from the value ofp of the curve with the best fit and from the value of log [A-] at log v = 0. t16~I. B ~ O M , Metal Ammine Formation in Aqueous Solution. Haase, Copenhagen (1941). ~aT~M. CALVZN and N. G. MY~LCmOR,Y. Amer. Chem. Soc. 70, 3270 (1948). tm D. D Y ~ N and L. O. SXLL~N,Acta Chem. $cand. 7, 663 (1953). L. (3. SILL~X%Acta Chem. Stand. 10, 186 (1956).
Problems in the use of organic reagents in chemical analysis--VII
143
The pH-metric titrations were made with an Orion KTS pH-meter with Metrohm glass electrodes. A Wilhelm bridge served as reference electrode. The ionic strength was kept at 0.3 M with sodium perchlorate. The titrations were carried out with 0.1 or 0.01 M sodium hydroxide solutions. The metal concentration was 10-4-5 x 10-4 mole/L, the pyridine-2-aldoximeconcentration was 2.5 x 10-L10 -~ mole/l. ; the acid concentration at the beginning was 5 x 10-8 mole/l. (in the case of titrations with 0.01 M sodium hydroxide 10-' mole/L). The equilibria were rne~ured in a VEB ultrathermostat set on 25.00 ± 0.05 C°. Proton resonance measurements. To decide whether the complexes were dia- or para-magnetic in aqueous solution we used a method based on proton resonance measurements.~I~ The magnetic measurements were accomplished at a complex concentration of 10-ffimole/L, the dissociation of the complex was suppressed with a ligand excess of 5 x 10-~ mole/l, at pH ~ 8. We are indebted to Dr. L. Ko~cz, Department of Atomic Physics L. EOtv6s University, for his valuable help in this measurement. The electrophoretic measurements. These measurements were carried out in a modified Buz~gh apparatus. The complex concentration in the solutions was 0.1-0.1 mole/l. Reagents, Pyridine-2-aldoxime p.a. (Fluka A.G.) Mnm+, Fes+, Co~+, Ni s+, Cu~+ and ZnI÷ solutions of 0-01 mole/1, concentration made from reagents of analytical purity; 0.1 and 0.01 M sodium hydroxide solutions; 0.01 M perchloric acid solution. Sodium perchlorate was prepared from A.R. perchloric acid and A.R. sodium carbonate (both chloride free). Chloroform A.R. (purified according to the Pharm. Hung. V.). RESULTS AND DISCUSSION Comparing the electronic spectra of the ligand with those of the complexes (Fig. 1) shows that in the spectra of the complexes--except that of the manganese (II) complex - - t h e ligand band at 34000 cm -1 shifted to 30000-32,000 cm -1, and in the spectra of the copper (II) and iron (II) complexes a new absorption band appeared at 25000 and 19000 cm -1, respectively. The existence of the manganese (II) complex was indicated by a small shift of the bands of the ligand, and by a significant decrease in the extinction coefficients calculated from the concentration of the ligand. The compositions of the complexes were determined by Job's (continuous variations) method, spectrophotometrically. The absorption of the solutions containing the central atom and the ligand in a suitably varying ratio was measured at a wavelength where under the concentration given, light was absorbed only by the complex. The cobalt (II), copper (II), zinc (II) complexes have the composition of MeA2, the iron (II), nickel (II) complexes have that of MeA3-. Analysis of the cobalt (II), copper (II) and zinc (II) complexes prepared from chloroform extracts, for C, H, N and metal confirmed the composition MeA~. To determine the charge on the complex electrophoretic investigations were carried out. It was found that the cobalt (II), copper (II) and zinc ([I) complexes have no electric charge, while the iron (II) and nickel (II) complexes carry a negative charge. In acid solution (pH = 3-4) of the copper (II) complex a species carrying a positive charge was detected. This complex seems to be the same as the complex of metal:ligand ratio 1 : l, described by KIltSON.~11> Whether the pyridine-2-aldoxime in the complex produces a strong or weak ligand field, can be decided on the base of the magnetic behaviour of the complexes. In this work the mononuclear, regular (non-protonated) complexes present in a dilute aqueous solution were studied, and their magnetic behaviour had to be determined in solutions of similar composition. According to the data in the literature, and partly to the results of our own experiments, in more concentrated aqueous solutions polynuclear complexes form, ~u> in acid solution protonated complexes/s'~°) and in relatively
144
K. BURO~R, I. EGYED a n d I. RuFF TASLE 1.--MAoNZTIC PROPERTIES O¥TRANSITIONAL METALPYRIDINE-2-ALDOXIMBCOMPI.EX.~;IN 10-= M
AQUEOUSSOLUTION
Central metal
MeA= complex
MeAs- complex
Fe II
diamagnetic
Co xx
weakly paramagnetic*
Ni Iz
paramagnetic
Cu ~x
paramagnetic
* The paxamagnetism is smaller than that of the aquocgmplex in solution of the same concentration.
f y....~-.
3.0
2.5
/
< Y/"
- - -
2.o
o 1.5
Co
Zn o
/
/',n/ /
0'51 09
•
r
8
1
7
J
6
I
5
r
4
-log [~A-'] F[(:}. 3.--Bjerrum's formation curves of the pyridine-2-aldoxime complexes of manganese (1I), cobalt (II), nickel (II) and zinc (II). (The circles mark the points calculated from the data of our measurements; the curve shown represents the best fitting normalized curve.)
concentrated alkaline solutions formation of hydroxo mixed complexes ag~ takes place. In dilute solutions magnetic susceptibilities cannot be determined by the usual methods. Thus we had to use a qualitative method based on measuring the proton resonance of water, which makes it possible to decide reliably even in very dilute (10-2 mole/L) solution whether the complex is diamagnetic or paramagnetic. The value of the ca,~R. W. GREENand M. C. K. SVAST],Australian J. of Chem. 16, 356 (1963).
Problems in the use of organic reagents in chemical analysis--VII
145
TAm.~ 2.--TtrE STA~rrY CONSTANTSOFTRANSITIONALMETALPYRIDINE-2-ALDOXIMECOMPLEXES Central metal atom MnxI
log K1
log K~
log K8
5.2 ± 0.2 3.9 ± 0.2
log t2 9.1 ± 0.2
FeZ1 8.4
7-7
5.1
CoXt
8.8 ± 0.1 8.8 ± 0-1 8.6 ± 0.2 8.6 ± 0.2
NiII
9.4 ± 0.2 7.1 ± 0.2 5.5 ± 0.2
Cut~ ZnH
8.9
5.65
5.8 -4- 0.1 5.3 ± 0.1
log 138
16.1
24.85 ± 0.15 21.2
17.6 ± 0.1 17.2 ± 0.2 16.5
Reference
Method of measurement
This work
Pot.
12 13
Spectrophot. Pot.
Pot. This work Spectrophot. 22.0 4- 0.4
This work
Pot.
14.55
11
Pot.
11.1 ± 0.1
This work
Pot.
magnetic moment cannot be determined by this method. But in the case of paramagnetic complexes one can decide, whether the paramagnetism of the aqueous solution of the complex is identical, or smaller than that of the aquocomplexes of the same concentration. The results, summarized in Table 1 show, that the pyridine-2aldoxime in its complexes formed with iron (II) and cobalt (II) is a ligand of strong field, and confirm that the nickel (II) complex is octahedral. The data obtained by spectrophotometric measurements on the composition of the complexes were confirmed by the Bjerrum's formation curves (Fig. 3) calculated from the data of pH-metric titrations. The pH-metric method showed the composition of the manganese (II) complex to be MnA~. The stability constants of the regular complexes of the 3d~3d a° transition metals with the pyridine-2-aldoxime, published in the literature and determined by us are summarized in Table 2. The data show, that (1) the stability constants do not follow the Irving-Williams c~a~ stability rule; (2) in the case of cobalt (II) complex the ratio of the successive stability constants is different from that of the statistical case, and K 1: K S = I ; according to the data of Hanania, tm in the iron (II) complex K1 < K~ < Ks.* The fundamental cause of both these deviations is that the pyridine-2-aldoxime in the complexes of iron (II) and cobalt (II) behaves as a strong-field ligand. This is the reason for the greater stability of these complexes than indicated by the IRXaNGWmLIA~tS rule. The unaccustomed ratio of the successive stability constants may be due to different causes: (1) the strong ligand field discussed above; (2) the donor •r-bond; (3) the hydrogen bridges stabilizing the MeA~ complex. In our case there is no possibility of the formation of hydrogen bridges. It was proved however that pyridine-2-aldoxime in its cobalt (II) and iron (II) complexes is a strong field ligand, and on the basis of the spectrophotometric invesitgations of ISRAELIt~°) the existence of the donor ~r-bond is likely. • The contradietious data of ~ and JtmaR~IS~18~cannot be taken reliable, for in his study he left the acid dissociation constant of the heterocyclic nitrogen of the ligand out of consideration. ito~ y . Iss~ta, Bull. Soc. Chim. Belg. 72, 123 (1963). tsa~H. IRVINOand R. J. P. WmUAMS,Nature, Lond. 162, 746 (1948). 10
SOME THEORETICAL AND PRACTICAL PROBLEMS IN THE USE OF ORGANIC REAGENTS IN CHEMICAL ANALYSIS--VII INVESTIGATIONS OF THE TRANSITION METAL COMPLEXES OF PYRIDINE-2-ALDOXIME K. BURG~, I. EGYEI) and I. RUFF Department of Inorganic and Analytical Chemistry, L. E6tvSs University,Budapest (Received 18 March 1965; in revised form 22 May 1965)
Abltraet--The composition of the mononuclear pyridin-2-aldoximecomplexes of 3dt-3d x°transition metals formed in neutral solution were determined. The successivestability constants of the manganese(ID, cobalt 0I), nickel (II)and zinc (II) complexes were measured. It was proved by proton resonance measurements that in the iron ([I) and cobalt (IT)complexespyddine-2-aldoximeis a strong field ligand, and the nickel 01) complex is octahedral. Tm~ compositions and stabilities of the dioxime complexes of 3dS-3d 1° transition metals were determined in our previous works, (l'a) and the factors influencing their stabilities explored. The analogous transition metal complexes of o-phenanthroline and ~-0~'dipyddyl were studied by IRVrN6t3) and recently by A~FJ~GG. (4) The ligands mentioned are in some respects similar. The functional group reacting with the metal ion i s :
\
/
/ C ¸ [I N
C I[ N
"\
but while the donor nitrogens of the dioximes belong to the oxime group, those of the other two ligands are heterocyclic nitrogens. All of the ligands mentioned are strong field ligands in their iron (II) complexes. The iron (II)-complexes are diamagnetic. The ligands mentioned above have free ~r-orbitals capable of accepting electrons from the central metal atom (back co-ordination). The existence of this kind of bond was proved in the case of the o-phenantroline and 0c-~'dipyddyl complexes by infra-red spectrophotometdc investigations by Busch and BAILAR,tS) and in the case of dimethylglyoxime complexes by our own infra-red studies. (2) The fundamental difference between the dioximes and the other two ligands mentioned is that the dioximes form complexes with a metal:ligand ratio 1:2, and have square planar structures; while o-phenanthroline and ~-~'dipyddyl form (1) K. BURt~ERand I. R u ~ , Talanta 10, 329 (1963). (J) K. BuRO~, I. Rul~ and F. RUFF,J. Inorg. Nucl. Chem. 27, 179 (1965). (s) H. IRVINOand D. H. M~LOR,./. Chem. Soc. 5222 (1962). (*) G. A ~ e ~ u 3 o , Helo. Chim. Acta 46, 2397, 2813 (1963). (5) D. H. Busot and J. C. B ~ Jr., J. Amer. Chem. Soc. 78, 1137 (1956).
139
140
K. BURGER, I. EGYED and I. RUFF
complexes with a metal:ligand ratio 1:3 and have octahedral symmetry. The different behaviour o f the dioximes can be explained most p r o b a b l y by the f o r m a t i o n o f strong intramolecular hydrogen bridges which stabilize the 1:2 complexes. Pyridine-2-aldoxime, a reagent used in analytical chemistry, t6.7"sJ occupies an intermediate place between the two types o f ligands mentioned. O f the two d o n o r nitrogens o f this ligand one is an oxime nitrogen, while the other is a heterocyclic one. This ligand is able to f o r m p r o t o n a t e d complexes t9'1°~ of composition M e A ( H A ) and polynuclear complexes, cm besides the regular m o n o n u c l e a r complexes. At present we are interested only in the m o n o n u c l e a r complexes o f composition analogous to that of the complexes mentioned above. The stability constants o f the copper (II) t g ' ~ and iron (II) (1~'13) pyddine-2-aldoxime complexes are given in the literature, some nickel (II) complexes with different composition have been prepared t14) and studied by infra-red spectrophotometry.t~5~ The acid dissociation constants of some p r o t o n a t e d complexes are also k n o w n P °'t2) The manganese (II), cobalt (II) and zinc (II) complexes have not yet been studied, and the stability constants o f the nickel (II) complex are unknown. Thus first o f all we had to study the mononuclear, regular (non-protonated) complexes o f manganese (II), cobalt (II), nickel (II) and zinc (II). EXPERIMENTAL Spectrophotometry. Our investigations in the ultra-violet and visible region were carried out on a Unicam SP 700 recording spectrophotometer. The spectrophotometric measurements were made in solutions of 10-~-5 × 10-~ mole/1. The molar extinction coefficients are shown in Fig. 1. The solutions of the metals and that of the ligand used for the measurements based on Job's method had a concentration of 2 × 10-' mole/l. For the determination of the successive stability constants of the cobalt II complex the degree of formation (a) of the CoAs complex was calculated from the extinction values measured at 30,000 wave numbers in solutions containing different metal and ligand concentrations at different pH values. The degree of formation of the complex (a) was plotted as a function of the logarithm of the free ligand concentration ([A-l) (Fig. 2). These (a -- log [A-]) points were fitted to a family of Y -- (log v)p normalized curves, where Y=
1 + pv + v~
v = (K1. K2)½ [A-]
and
p
=
K~J. K~-i
The successive stability constants (K1 and K0 were calculated from the value ofp of the best fitting curve and from the value of log [A-] at the point where log v = 0. The acid dissociation constants of the ligand determined from spectrophotometric measurements and used for the calculations of the free ligand concentration were: K1 = --3.4; logKa = --10,0 (where K1 was the acid dissociation constant of the heterocyclic NI-I+ group, and K~ that of the ~6~H. HARTK~a', Naturwissenschaften 45, 211 (1958). ~7~ H. H_~rgAw&, Z. Analyt. Chem. 170, 399 (1959). ~s~H. I - h R - - , Z. Analyt. Chem. 176, 185 (1960). cg~ C. H. LILTand C. FAN Lro, J. Amer. Chem. Soc. 83, 4167, 4163 (1961). clo~I. MAStmA and K. Srn~A, Science Reports (Osaka Univ.) 11, 3 (1962). c ~ B. KmsoN, Bull. Soc. Chim. France 1030 (1962). c~2~G. I. H. HArqAr,rm and D. H. I R ~ , d. Chem. Soc. 2745 (1962). ~8~ y. l S ~ L I and E. JtmoREIs, Bull. Res. Council oflsrael, 11 A, 121 (1962). ~1~ R. A. KRAtrSEand D. H. BUSCH,d. Amer. Chem. See. 82, 4830 (1960). t~s~R. A. K~usE, N. B. CoLTrxtre and D. H. Btrscrl, d. Phys. Chem. 65, 2216 (1961).
Problems in the use of organic reagents in chemical analysis---VII
141
HA 1'6-
I-2-
0"8
0"4 O0 1"6 1'2 0-8
L
(
A-
Ni
Mn
Cu
Fe
Zn
I
\
0'4 0"0 E
1"6 1'2 0"8 0"4 0"0 1'6 1"2 0"8
I
0"0,50 ] 42
[
I
34
[
i
26
i
50
42
34
26
kK
FIe. 1.--The absorption spectra of pyridine-2-aldoxime and its manganese (l]), iron (N), cobalt (II), nickel (11), copper 013 and zinc 0I) complexes. oxime group). The compositions of the solutions were: cobalt concentration 2 × 10-5-4 x 10--s mole/L, ligand concentration 10-5-2 x 1e-a; pH = 6.5--8.0; NaC10, concentration = 0"3 mole/1. For the spectrophotometric investigations of the complexes solutions of the metal and the ligand of suitable conczntration were mixed; quantitative formation of the complexes was a~ured by choicz of a suitable pH. The cobalt (11), copper (ll) and zinc (H) complexes were extracted by chloroform
142
K. BvgG~t, I. EoY]~a and I. Ru]~
I'O
0.5
o/ log ~',=8'6 log ~z= 17'2
0
I 9
i
i 8
7
- l o g EA:]
Fro. 2.--The formation degree of the cobalt (H)-pyridine-2-aldoxime complexes of CoAs composition plotted as a function of the logarithm of free ligand concentration. ( Q [ ] A mark the points calculated from the data of our measurements; the curve shown represents the best fitting normalized curve.) Cobalt concentration Q 2 x 10-~ M, [] and A 4 × 10-~ M. Pyridine-2-aldoxime concentration 10-5-2 × 10-a M. The p H of the solution O 6.5, [] 7.2, A 8.0. from the aqueous solutions.
The solid complexes were prepared by evaporating the extract in
V/IC//O.
Theptl-metricequilibriummeasurements. These measurements were carried out by BS~RRUM'SCle'17~ method. The formation curves were evaluated by the method of DxqksSSN and SILL'S,{18j based on the use of normalized curves. The points of the (t~ -- log [A-]) formation curve, {~e~calculated from experimental data were fitted to a family of Y -- 0og v)~ normalized curves, where Y=
v + 2v~p 1 + v + _pvs
v = Ka[A-]
and K,
The successivestabilityeomtants (Kt, K O were calculated from the value ofp of the curve with the best fit and from the value of log [A-] at log v = 0. t16~I. B ~ O M , Metal Ammine Formation in Aqueous Solution. Haase, Copenhagen (1941). ~aT~M. CALVZN and N. G. MY~LCmOR,Y. Amer. Chem. Soc. 70, 3270 (1948). tm D. D Y ~ N and L. O. SXLL~N,Acta Chem. $cand. 7, 663 (1953). L. (3. SILL~X%Acta Chem. Stand. 10, 186 (1956).
Problems in the use of organic reagents in chemical analysis--VII
143
The pH-metric titrations were made with an Orion KTS pH-meter with Metrohm glass electrodes. A Wilhelm bridge served as reference electrode. The ionic strength was kept at 0.3 M with sodium perchlorate. The titrations were carried out with 0.1 or 0.01 M sodium hydroxide solutions. The metal concentration was 10-4-5 x 10-4 mole/L, the pyridine-2-aldoximeconcentration was 2.5 x 10-L10 -~ mole/l. ; the acid concentration at the beginning was 5 x 10-8 mole/l. (in the case of titrations with 0.01 M sodium hydroxide 10-' mole/L). The equilibria were rne~ured in a VEB ultrathermostat set on 25.00 ± 0.05 C°. Proton resonance measurements. To decide whether the complexes were dia- or para-magnetic in aqueous solution we used a method based on proton resonance measurements.~I~ The magnetic measurements were accomplished at a complex concentration of 10-ffimole/L, the dissociation of the complex was suppressed with a ligand excess of 5 x 10-~ mole/l, at pH ~ 8. We are indebted to Dr. L. Ko~cz, Department of Atomic Physics L. EOtv6s University, for his valuable help in this measurement. The electrophoretic measurements. These measurements were carried out in a modified Buz~gh apparatus. The complex concentration in the solutions was 0.1-0.1 mole/l. Reagents, Pyridine-2-aldoxime p.a. (Fluka A.G.) Mnm+, Fes+, Co~+, Ni s+, Cu~+ and ZnI÷ solutions of 0-01 mole/1, concentration made from reagents of analytical purity; 0.1 and 0.01 M sodium hydroxide solutions; 0.01 M perchloric acid solution. Sodium perchlorate was prepared from A.R. perchloric acid and A.R. sodium carbonate (both chloride free). Chloroform A.R. (purified according to the Pharm. Hung. V.). RESULTS AND DISCUSSION Comparing the electronic spectra of the ligand with those of the complexes (Fig. 1) shows that in the spectra of the complexes--except that of the manganese (II) complex - - t h e ligand band at 34000 cm -1 shifted to 30000-32,000 cm -1, and in the spectra of the copper (II) and iron (II) complexes a new absorption band appeared at 25000 and 19000 cm -1, respectively. The existence of the manganese (II) complex was indicated by a small shift of the bands of the ligand, and by a significant decrease in the extinction coefficients calculated from the concentration of the ligand. The compositions of the complexes were determined by Job's (continuous variations) method, spectrophotometrically. The absorption of the solutions containing the central atom and the ligand in a suitably varying ratio was measured at a wavelength where under the concentration given, light was absorbed only by the complex. The cobalt (II), copper (II), zinc (II) complexes have the composition of MeA2, the iron (II), nickel (II) complexes have that of MeA3-. Analysis of the cobalt (II), copper (II) and zinc (II) complexes prepared from chloroform extracts, for C, H, N and metal confirmed the composition MeA~. To determine the charge on the complex electrophoretic investigations were carried out. It was found that the cobalt (II), copper (II) and zinc ([I) complexes have no electric charge, while the iron (II) and nickel (II) complexes carry a negative charge. In acid solution (pH = 3-4) of the copper (II) complex a species carrying a positive charge was detected. This complex seems to be the same as the complex of metal:ligand ratio 1 : l, described by KIltSON.~11> Whether the pyridine-2-aldoxime in the complex produces a strong or weak ligand field, can be decided on the base of the magnetic behaviour of the complexes. In this work the mononuclear, regular (non-protonated) complexes present in a dilute aqueous solution were studied, and their magnetic behaviour had to be determined in solutions of similar composition. According to the data in the literature, and partly to the results of our own experiments, in more concentrated aqueous solutions polynuclear complexes form, ~u> in acid solution protonated complexes/s'~°) and in relatively
144
K. BURO~R, I. EGYED a n d I. RuFF TASLE 1.--MAoNZTIC PROPERTIES O¥TRANSITIONAL METALPYRIDINE-2-ALDOXIMBCOMPI.EX.~;IN 10-= M
AQUEOUSSOLUTION
Central metal
MeA= complex
MeAs- complex
Fe II
diamagnetic
Co xx
weakly paramagnetic*
Ni Iz
paramagnetic
Cu ~x
paramagnetic
* The paxamagnetism is smaller than that of the aquocgmplex in solution of the same concentration.
f y....~-.
3.0
2.5
/
< Y/"
- - -
2.o
o 1.5
Co
Zn o
/
/',n/ /
0'51 09
•
r
8
1
7
J
6
I
5
r
4
-log [~A-'] F[(:}. 3.--Bjerrum's formation curves of the pyridine-2-aldoxime complexes of manganese (1I), cobalt (II), nickel (II) and zinc (II). (The circles mark the points calculated from the data of our measurements; the curve shown represents the best fitting normalized curve.)
concentrated alkaline solutions formation of hydroxo mixed complexes ag~ takes place. In dilute solutions magnetic susceptibilities cannot be determined by the usual methods. Thus we had to use a qualitative method based on measuring the proton resonance of water, which makes it possible to decide reliably even in very dilute (10-2 mole/L) solution whether the complex is diamagnetic or paramagnetic. The value of the ca,~R. W. GREENand M. C. K. SVAST],Australian J. of Chem. 16, 356 (1963).
Problems in the use of organic reagents in chemical analysis--VII
145
TAm.~ 2.--TtrE STA~rrY CONSTANTSOFTRANSITIONALMETALPYRIDINE-2-ALDOXIMECOMPLEXES Central metal atom MnxI
log K1
log K~
log K8
5.2 ± 0.2 3.9 ± 0.2
log t2 9.1 ± 0.2
FeZ1 8.4
7-7
5.1
CoXt
8.8 ± 0.1 8.8 ± 0-1 8.6 ± 0.2 8.6 ± 0.2
NiII
9.4 ± 0.2 7.1 ± 0.2 5.5 ± 0.2
Cut~ ZnH
8.9
5.65
5.8 -4- 0.1 5.3 ± 0.1
log 138
16.1
24.85 ± 0.15 21.2
17.6 ± 0.1 17.2 ± 0.2 16.5
Reference
Method of measurement
This work
Pot.
12 13
Spectrophot. Pot.
Pot. This work Spectrophot. 22.0 4- 0.4
This work
Pot.
14.55
11
Pot.
11.1 ± 0.1
This work
Pot.
magnetic moment cannot be determined by this method. But in the case of paramagnetic complexes one can decide, whether the paramagnetism of the aqueous solution of the complex is identical, or smaller than that of the aquocomplexes of the same concentration. The results, summarized in Table 1 show, that the pyridine-2aldoxime in its complexes formed with iron (II) and cobalt (II) is a ligand of strong field, and confirm that the nickel (II) complex is octahedral. The data obtained by spectrophotometric measurements on the composition of the complexes were confirmed by the Bjerrum's formation curves (Fig. 3) calculated from the data of pH-metric titrations. The pH-metric method showed the composition of the manganese (II) complex to be MnA~. The stability constants of the regular complexes of the 3d~3d a° transition metals with the pyridine-2-aldoxime, published in the literature and determined by us are summarized in Table 2. The data show, that (1) the stability constants do not follow the Irving-Williams c~a~ stability rule; (2) in the case of cobalt (II) complex the ratio of the successive stability constants is different from that of the statistical case, and K 1: K S = I ; according to the data of Hanania, tm in the iron (II) complex K1 < K~ < Ks.* The fundamental cause of both these deviations is that the pyridine-2-aldoxime in the complexes of iron (II) and cobalt (II) behaves as a strong-field ligand. This is the reason for the greater stability of these complexes than indicated by the IRXaNGWmLIA~tS rule. The unaccustomed ratio of the successive stability constants may be due to different causes: (1) the strong ligand field discussed above; (2) the donor •r-bond; (3) the hydrogen bridges stabilizing the MeA~ complex. In our case there is no possibility of the formation of hydrogen bridges. It was proved however that pyridine-2-aldoxime in its cobalt (II) and iron (II) complexes is a strong field ligand, and on the basis of the spectrophotometric invesitgations of ISRAELIt~°) the existence of the donor ~r-bond is likely. • The contradietious data of ~ and JtmaR~IS~18~cannot be taken reliable, for in his study he left the acid dissociation constant of the heterocyclic nitrogen of the ligand out of consideration. ito~ y . Iss~ta, Bull. Soc. Chim. Belg. 72, 123 (1963). tsa~H. IRVINOand R. J. P. WmUAMS,Nature, Lond. 162, 746 (1948). 10