Studies on some Ni(II) and Zn(II) diaminomonocarboxylate complexes

L inorg, nucl. Chem. Vol. 43, pp. 1591-1597, 1981 Printed in Great Britain.

0022-19021811071591-.07502.0010 Pergamon Press Ltd.

STUDIES ON SOME Ni(II) AND Zn(II) DIAMINOMONOCARBOXYLATE COMPLEXES ETELKA FARKAS, ARTHUR GERGELY* and ELEONORA KAS Institute of Inorganic and Analytical Chemistry, Lajos Ko~suth University, H-4010 Debrecen, Hungary (Recdved in revisedform 1 May 1979; receivedfor publication 7 October 1980)

Abstract--pH-metric, calorimetric, NMR and spectro-photometric studies were carried out on the Ni(lI) and Zn(II) complexes of 2,3-diaminopropionic acid (dapa), 2,4-diaminobutyric acid (daba), 2,5-diaminopentanoic acid or ornithine (Orn) and 2,6-diaminohexanoic acid or lysine (Lys). It is concluded that the maximum number of coordinated ligands (with the exception of totally deprotonated Orn) is three in the case of Ni(II), but only two in the Zn(II) complexes. The ~o-aminogroups of dapa, daba and Orn take part in the coordination in the Ni(II) complexes. Hence, with Ni(II) both "diamine type" and "glycine-like" complexes are formed. The tendency to "diamineqike" coordination increases from Orn to dapa, so much so that dapa behaves much rather as a C-substituted 1,2-diamino ethane than a substituted glycine. Only Lys is coordinated to the Ni(II) ion exclusively in a "glycine-like" manner. Zn(II) complexes were studied only to a pH value of 9 (because of precipitation). In this pH region the terminal NH2 groups take part in the coordination only in the cases of dapa and daba. With Orn and Lys "glycine-like" parent and mixed hydroxo complexes are formed. INTRODUCTION

The ligands 2,3-diaminopropionic acid (dapa), 2,4diarninobutyric acid (daba), 2,5-diaminopentanoic acid or ornithine (Orn) and 2,6-diaminohexanoic acid or lysine (Lys) are a, to-diaminomonocarboxylic acids. The equilibria in systems containing these ligands and metals are very complicated, because there is a possibility for the formation of both "glycine-like" and "diamine-like" complexes. In addition, in the cases of Ni(II) and Zn(II) bis and tris complexes may occur. As regards the complexes formed between these ligands and the transition metals the Cu(II) complexes have been studied fairly thoroughly[I--6]. Results have also been published for the Ni(II) complexes[5-11]. However, previously it was only known that, because of their ttexidentate character, dapa and daba can coordinate to Ni(II) in various ways [7]. Later Brubaker and Busch[5] studied the Ni(II)-Orn and the Ni(II)-Lys systems spectrophotometrically and identified various complexes. Furthermore, they established that, following the "glycine-like" complexes formed in the low pH range, at higher pH values the to-amino groups of all three ligands take part in the coordination. They assume that the bonding is intramolecular in the cases of daba and Orn but intermolecular in the case of Lys. Experiments performed on the Ni(II)--dapa system pHmetrically[10] and by tH NMR methods, and on the Ni(II)-Orn system by CD[9] and equilibrium methods, confirmed the conclusions of Brubaker and Busch[5]. Recently Brookes and Pettit[6] carried out extensive studies to clarify the equilibrium relations of some 3d transition metal-a, to-diaminomonocarboxylic acid systems. They found the equilibrium systems containing Ni(II) to be especially complicated. Their conclusions on the coordination type of the donor groups agreed with those of Brubaker and Busch[5] but, at the same time, they inferred the formation of many new species too. Up to the present only qualitative findings have been made on the Zn(II)-a, to-diaminomonocarboxylate *Author to whom correspondence should be addressed.

systems[6,8]. It is reported that quantitative evaluation is hindered both by the precipitate occurring at p H - 7 and by the weakness of the interaction between the metal and the ligands. As regards the flexidentate character of these ligands, many questions concerning the Ni(II)-a, to-diaminomonocarboxylate complexes remain to be clarified. Relatively little is known about the corresponding Zn(II) equilibrium systems. The aim of our work is therefore a detailed equilibrium and thermodynamic study of the complexes formed between the ligands dapa, daba, Orn and Lys and the metals Ni(II) and Zn(II). Spectrophotometric and NMR measurements are also performed for a better understanding of the flexidentate character of these ligands. EXPERIMENTAL

REANAL and FLUKA p.a. D,L-amino acids were used after purification by recrystallization from an ethanol-water mixture. The NiCI2 and ZnC12stock solutions were prepared by dissolving NiC12'6 H20 and ZnO in water and dilute HC1 solution, respectively. The concentrations of the metal ions were checked gravimetrically via the oxinates. pH and calorimetric measurements. For the pH-metric titrations samples were used in 1:1, 1:2, 1:3, 1:4 and 1:5 metal ion to ligand ratios at a total ligand concentration 0.003 M. The ionic strength was made up to 0.2 M KCI. In the case of Ni(II) the titrations were performed up to pH 10.5-11.5. In the systems containing Zn(II), measurements were made only up to pH 9.0.-9.5, because of the hydrolysis. Calorimetric experiments in the case of Ni(II) were carried out at the same metal-ligand ratios and concentrations as in the pH-metric titrations. With the Zn(II) amino acid systems the calorimetric data could not be evaluated because of the hydrolysis occurring during the measurements. The pH-metric titrations were made on a Radiometer prim 54 instrument. The calorimetric measurements were carried out on an LKB 8700-2 instrument as described earlier[12]. The data were evaluated by a method already published [13, 14]. Spectrophotometric studies. Visible, UV and near IR measurements were only performed in the case of Ni(II). A Beckman-Acta MIV double-beam recording spectrophotometer was used. The metal-to-ligand ratio was 1:4 at a Ni(lI) concen-

1591

1592

E. FARKASa al.

tration of 0.02M. With Lys, examinations were made at nickel concentrations, of 0.04, 0.06 and 0.08M too. H NMR studies. With Zn(lI) complexes 'H NMR titrations were carried out on a JEOL JMN MH-100instrument in water solution. The pH-dependent chemical shifts of the CH and CH2 protons adjacent to amino groups were followed. The Zn(II) concentration of the samples was 0.10M. The metal-ligand ratio in the case of dapa was 1:2, but with the other ligands 1:4. In the samples not containing metal ions, the ligand concentrations were also 0.20 and 0.40M, respectively. No external standard was used, but the chemical shift referred to the water proton line was measured in internal lock. To maintain the dilution as low as possible a 3.5 M KOH solution was used as a titrant. The experimental conditions and the measurements were as already described[4]. RESULTSANDDISCUSSION Ni(II) complexes Evaluation of the Ni(II)-a, to-diaminomonocarboxylate equilibrium systems is mainly hindered by the difficulty in the identification of the high numbers of species formed. With regard to the results already published [5, 6], spectrophotometric studies were made to identify the Ni(II) complexes. The pH-dependence in the absorption maxima is given in Table I. For comparison, Table 2 contains the absorbance values of the Ni(II) systems of the non flexidentate ligands as glycine, 1,3diaminopropane and diaminoethylene measured at pH 10.5. From the measured colour changes, which are also visible, the following conclusions can be drawn. (i) The colour of the Ni(II)-Lys system is always blue,

independently of the studied metal-ligand ratio and the pH value employed. As can be seen in Tables 1 and 2, the spectral data are similar to those of the glycine complexes. (ii) The colour of the Ni(II)-Orn system at a l : l metal-ligand ratio is blue at each pH value. With metalligand ratios higher than 1 : 1, below pH ~ 8 the colour is likewise blue and the absorption maxima are essentially the same as those for the Ni(II)-glycine system. At pH > 8, however, the colour changes gradually to violet and the absorption maxima shift to higher energies. (iii) The Ni(II)-daba system behaves similarly to that of Ni(II)-Orn, but in this case the band shift begins already at pH 6.5. The absorption maxima at pH > 8 are at even higher energies than those of the NiA3 = type complex of 1,3-diaminopropane. (iv) The Ni(II)-dapa system differs from those of Ni(II)-Orn and Ni(II)-daba only in that the spectral data characteristic of the Ni(II)-l,2-diaminoethane system appeared at a comparatively low pH. On the other hand, the samples with 1:2 metal-ligand ratios were yellow when the alkali added to the samples was exactly equivalent to the ligand (i.e. equivalent to the protons dissociable from the ligand). Though this colour remains unchanged at pH ~ 9.0-9.5, on heating the solution turns violet, but becomes yellow again on cooling. This interesting phenomenon was earlier observed with other complexes of 1,2-diamonethane derivatives[15, 16] and was attributed to an equilibrium between the planar and octahedral forms of the NiA2-type complexes. However,

Table 1. pH-dependent values of the absorption maxima in the case of a 1:4 ratio of Ni(II) to the ligands dapa, daba, Orn and Lys (Vm,~"10-3'cm-') dapa

~a

6.2

10.5

daba

17.4

Orn

Lys

28.1

6.7

-

-

-

10.0

16.7

27.5

9.9

15.6

26.5

7.0

10.6

17.9

28.6

10.1

17. e

e7.5

9.9

15.8

26.7

7.2

-

-

-

7,5

10.8

18.0

28.6

8.0

10.9

18.2

28.7

8°5

-

-

-

9,0

11.2

18.2

28.9

9.6

11.3

18.4

29.0

10.5

11.3

18.4

29.0

10.4

10.5

10.8

17.8

17.9

17.9

27.9

27.9

27.9

i0.0

10.2

10.7

16.0

16.7

17.5

9.8

15.8

26.6

9.9

16.3

27.2

I0.0

16.4

27.4

10.0

16.5

27.4

i0.0

16.5

27.4

27.0

27.5

27.8

Table 2. Absorptionmaxima values in systems with a 1:4 ratio of Ni(II)to the ligands glycine, 1,3-diaminopropane and 1,2-diaminoethaneat pH = 10.5 (~;,,ix"10-3. cm-')

1,2-diaminoethane

ii.4

18.5

28.6

1,3-diamlnopropane

I0.7

17.3

27.6

9.9

16.8

27.6

glycine

Ni(II)- and Zn(II)--diaminomonocarboxylatecomplexes on increase of the metal-ligand ratio of these samples to greater than 1:2, and of the pH to be above pH 9, the colour changes irreversibly to violet. The reason for this is presumably that octahedral MA3[16] and mixed hydroxo complexes are formed. For a better understanding of the phenomena outlined above, additional spectrophotometric measurements were made as a function of temperature at a 1 : 2 metalligand ratio. The character of the spectrum in the visible and the near IR range at 25°C is reminiscent of that of simultaneously present octahedral and planar complexes. As the temperature is raised the ratio of the octahedral Ni(II) complex gradually increases, and above 55°C the spectrum corresponds entirely to that of the octahedral complex. Magnetic susceptibility measurements in solution were carried out by the 'H NMR method to establish the ratio of the complexes. It was found that at 28°C, at a total concentration of the complex 0.02 M, with 0.2 M (KC1) ionic strength about 70% of the Ni(II) is in paramagnetic form. The yellow colour temporarily appears also in the case when the metal-ligand ratio is higher than 1:2 and the pH is above 9.5. This can probably be explained kinetically, for the rate-limiting step is the NiA:->NiA3 process. pH-metric measurements were designed so as to yield as many experimental data as possible in the widest pH-range. Titrations were therefore performed at five different metal-ligand ratios. In the evaluation of the pH-metric data, the assumptions on the species present, drawn on the basis of the spectrophotometric measurements, were also taken into account. The stability data obtained and the earlier published protonation contants of the ligands [4] are listed in Table

1593

3. The data in Table 3 include the protonation constants of the ligands. To be able to make conclusions about the metal-ligand interaction, it is desirable to correct the overall stability constants for the HA protonation constants. Hence, the deprotonation constants of the ligands bound to the metal are given in Table 4. The calorimetric measurements were carried out on all systems examined. Nevertheless, evaluation of the experimental data was not possible in the case of the Ni(II)-daba system because of the very high number of species formed. In the other cases too, we succeeded in obtaining values only for the species formed in relatively high concentrations. The AH and AS values obtained were corrected for the deprotonation heats of the ligands as follows: AHderive d = AHNiAmHn- nAHHA

and ASderived = ASNiAmHn --nASHA.

The derived data are to be found in Table 5. It is obvious from the results given above that in every system species containing three ligands are formed. However, with the exception of daba the ligands do not form all the species theoretically possible. The complexes which might theoretically be present in all the systems are depicted in Fig. 1. Figure 1 also indicates the complexes which are not formed in the appropriate system: in the Ni(II)-dapa system the Ni(AH)32+ and NiA(AH)2 + species (--); in the Ni(II)Orn system the NiA2(AH) and NiA3- complexes (---); and in the Ni(II)-Lys system the NiA + and NiA2 species (-.-.-). On the basis of the data in Tables 1-4 and Fig. 1, the following conclusions can be drawn regarding the Ni(II) complexes.

Table 3. Protonation constants of dapa, daba, Orn and Lys, and stability constants of the nickel complexes: [~qpr =

[NiqApHr] [Ni]q[A]P[a]r log~

~r dapa

daba

0rn

Lys

HA +

011

9.39

10.23

10.52

10.66

~zA ÷

012

16.08

18.46

19.35

19.86

H3A2+

013

17.40

20.35

21.42

22.01

Ni~ AH )2+

iii

13.43

14.89

15.04

15.50

NIA +

ii0

8.13

8.97

6.83

-

Ni( AN )2+

122

26.36

29.15

29.35

30.22

• tAc AN )+

121

21.04

23.12

20.95

20.38

NiA 2

120

15.17

16.34

11.68

-

NI(AN )~+

133

-

43.55

42.71

43.67

132

-

36.61

33.45

33.55

131

26.46

29.49

-

23.18

130

18.35

19.80

-

12.31

• iA2(AN )

•

The

log~HA relates

IO~H ~

to the

@t-amino group,

to all three donor groups.

lOg~H ~

to the two amino groups,

E. FARKAS et aL

1594

Table 4. Deprotonationconstants of the Ni(II) complexes pK dapa

daba

0rn

Lys

NIA + + H +

5.30

5.92

7.21

-

NI(AN )2+ ~

NIA( AH )+ + H +

5.32

6.03

8.40

9.84

NIA(AH)+

NiA 2 + H ÷

5.88

6.78

9.37

-

a+

-

6.94

9.25

10.13

NIA2(AH ) + H +

-

7.22

-

I0.37

8.11

9.69

-

10.87

NI(AH)2+

~I(AN

~

~

)~+ ~

~.~

NIA(AH ): ~

):

+

Table 5. Thermod''namic data on the proton and Ni(II) complexesof the ligands dapa, daba, Orn and Lys ~ S -aerlvea . .J "M'I.K -I

~HderivedKJ'M'l dapa

Orn

Lys

dapa

Orn

Lye

NI(AH )2+

12.9

20.2

18.8

33

18

32

NiA +

26.0

68

50

-

37

23

73

54

63

107

34

-

-

39

r

Ni( AH )2+

-

36.4

43.7

NiA( AH )+

44.7

43.1

36.3

NiA 2

54.5

56.4

-

Ni( AH )~+

-

-

I

54.6

L

comparing the above mentioned derived thermodynamic quantities with the data relating to the Ni(II)--glycine complexes. Hence, in the case of Lys, which contains the chain-terminal amino group in the ~-position, it can be stated that the ~-group does not take part in the coordination and only "glycine-type" complexes are formed. Terminal amino group coordination of fully deproNiA2 ""%./~ '- NiAa(AH) tonated Lys, i.e. polynuclear complex formation, could not be detected even with wide ranges of metal ion concentrations and metal to ligand ratios. : i "- N,A; (ii) The data on the Ni(II)--Orn system show that the Fig. 1. Complexes that might form in the Ni(II)--a, ~o-diamino- fully deprotonated ligand also coordinates to the metal monocarboxylate systems. ~ , Ni(II)-dapa; ---, Ni(ll)-Orn; through the to-amino group. This is manifested primarily -.-.-, Ni(II}--Lys. in the increased stability constants of the NiA ÷ and NiA2 species, which are about one order of magnitude greater (i) The deprotonation constants of the Lys complexes than those of the glycine complexes. (At the same ionic do not differ substantially from the constants relating to strength the appropriate stability constants for the glythe to-amino group of non-coordinated Lys. The totally cine complexes are: log K, = 5.65 and log K2 = 4.75.) deprotonated complexes are formed above pH 9 only, The finding is also supported by decreased pK value of where the non-coordinated ligand would also lose the the terminal NH3 + groups of the bonded ligand comproton. The relative stability constants (pK, A-+-log K) pared to that of the non-coordinated one, and by the and the spectrophotometric data are practically the same spectral data above pH 9, the absorbances being nearly as those of the appropriate glycine complexes (see the same as those in the Ni(II)-l,3-diaminopropane. Tables 1 and 2). A similar conclusion can be reached by Figure 2 shows the pH-dependent percentage concen-

Ni 2.

,- NiAH)2'

" Ni(AH)~"--'~'~

Ni(AH): +

Ni(ll)- and Zn(lI)-diaminomonocarboxylatecomplexes

T

,2+ r2

I0

~_0

~

Ni(AH)2" NRAH)~*

6

4

5

6

7

8

9

I0

II

pH

Fig. 2. pH-dependent percentage concentration distributions of species formed in the Ni(II)-Orn system. CA= 8.10 -3 M; CNi(II) = 2.10-2 M. tration distributions of the species formed in the Ni(II)Orn system. It can be seen in Fig. 2 that only two fully deprotonated ligands can coordinate to the Ni(II). This complex can form both through the deprotonation of Ni(AHh 2÷ and through the loss of a ligand from the Ni(AH)32÷ complexes when a deprotonation process also takes place. The terminal amino group in this sevenmembered chelate ring naturally participates in only a weak interaction, and therefore the stability constant of the NiA2 complex is not too high. (iii) In the Ni(II)-daba system all the feasible complexes are formed in detectable concentrations. The formation constants, suggest that coordination of the ligand to the metal is "glycine-type" in the protonated complexes. On the other hand, in the fully deprotonated species the chain-terminal amino group is also coordinated to the metal. These findings are supported both by the stability constants and by the spectral data: the absorption maxima up to pH 7 agree with those of the Ni(II)-glycine complex but at higher pH values the spectrum resembles that of the Ni(II)-l,3-diaminopropane species.

1595

Formation of the NiA3- complex is supported by the spectrophotometric and pH-metric results. However, the log /3 value of 19.8 is too high compared to the corresponding value for the Ni(II)--l,3-diaminopropane complex (13.4). Further investigations must be made to clarify this surprising phenomenon. The flexidentate character is strongest for daba. This is demonstrated in the various complexes present in the Ni(II)-daba system; i.e. the tendencies forwards "glycine-type" and "amine-type" bonding are almost the same. (iv) In the case of the Ni(II)--dapa system protonated complexes are formed only below pH 6. Therefore, Ni(AH)22÷ and Ni(AHh 2÷ complexes do not occur in detectable concentrations. The deprotonation constants of the complexes are generally lower than 6, and above pH 8 the spectrophotometric data are almost the same as those for the Ni(II)-l,2-diaminoethane system. Also, as already discussed in detail, the Ni(dapah complex similarly to the NiA2 complexes of the Csubstituted ethylenediamine derivatives, occurs in planar ~octahedral equilibrium. All the results obtained for the Ni(II)-dapa system, including the calorimetric measurements, prove the inclination for "amine-type" coordination. This is so definite in the complexes containing the ligands in fully deprotonated form that the N-donor atoms are involved in the first coordination sphere. Zn(II) complexes Because of precipitation study of the Zn(II)-c~, ~odiaminomonocarboxylate systems was carried out only below pH 9. These systems are also very complicated. However, taking into account the results for the Ni(II) complexes, the nature of the ligands studied and the behaviour of other Zn(II) complexes it was possible to identify the species formed. The data are listed in Table 6, and the derived constants, evaluated in the same way as in the case of Ni(II), in Table 7. From the data in Table 6 it follows that the most probable bonding in all Zn(AH)2÷ and Zn(AHh 2÷ complexes is "glycine-like". At the same time it is not easy to explain the deprotonation constants of the Zn(II)-Orn and Zn(II)-Lys systems if they are compared to those of the Ni(II)-dapa and Ni(II)--daba systems. In addition, relating to those of the corresponding Ni(II) complexes the stability of the fully deprotonated Ni(II)-Orn and

Table 6. Stability constants of the Zn(lI) complexesof dapa, daba, Orn and Lys 1og~

qpr dapa

daba

Orn

Lys

Zn( AH )2+

iii

12.59

13.97

14.25

14.72

ZnA +

ii0

6.31

6.70

6.17

6.32

Zn( All )2+

122

24.62

27.55

27.85

28.85

ZnA( AH )+

121

18.43

20.32

19.31

19.67

ZnA 2

120

11.66

12.30

-

-

ZnA2H- I

12 -i

0.94

-

-

-

1596

E. FARKASet aL Table 7. Deprotonationconstants of the Zn(II)complexes dapa

daba

0rn

Lys

Zn( A H )2+ ~

ZnA+

+ H+

6.28

7.27

8.08

8.40

Zn(AH )2+ ~

ZnA( .4.1-I )+ + H +

6.19

7.23

8.54

9.15

ZnAf AH )+ ~

ZnA 2 + H +

6.77

8.02

-

-

-

-

z,,~ 2

~.

, zm~- z+a

1o. 72

+

Ni(II)-Lys complexes are rather high. Hence, the conclusion may be drawn that the ZnA+ and Zn(AH) 2+ complexes in these cases are not deprotonated ones at all, hut are rather hydroxo complexes. This conclusion may be all the more valid, if it is considered that the difference between the proton and the hydroxo ion in the calculations is only one of sign. Therefore, the ZnA÷ species can be substituted by the [Zn(AH)(OH)]+ complex. To acquire more information about the composition of the Zn(II) complexes, IH NMR measurements were carried out. Results which relate to the Zn(II)--dapa and Zn(II)-Lys systems are depicted in Figs. 3 and 4. In the figures curves a denote the results obtained for the ligands only, and curves b those results which relate to the systems also containing Zn(II). (The figures for Zn(II)-Orn and Zn(II)--daba are similar to those for Zn(II)-Lys and Zn(II)-dapa, respectively.) In the case of Orn and Lys, up to pH 9 there is practically no difference between the chemical shifts of the CH2 protons close to the chain-terminal amino group. This means that these amino groups remain protonated below pH 9 and so cannot coordinate to the metal. The conclusion can be drawn, therefore that the ZnA÷ complex might be sub-

-

b

G

CH

,.5o

d

I00

5

6

7

8

9

I0

pH

Fig. 4. Chemical shifts of the CH and CHz protons near to the amino groups. (a) Shifts for dapa in the absence of metal; (b) Shifts in a solution with a 1:4 ratio of Zn(II) to dapa.

stituted by a [Zn(AH)(OH)] ÷ species, and ZnA2H+ by [Zn(AH)2(OH)]+. Hence, "glycine-type" parent and mixed hydroxo complexes are formed in the Zn(II)-Orn and Zn(II)-Lys systems. This is represented in Fig. 5,

200

CH 2

d

150

CH

I00

f

J 5

6

7

8

9

I0

pH

Fig. 3. Chemicalshifts of the CH and CH2 protons near to the amino groups. (a) Shifts for Lys in the absence of metal; (b) Shifts in a solution with a 1:4 ratio of Zn(II) to Lys.

Ni(II)- and Zn(II)--diaminomonocarboxylate complexes

el

IH NMR results suggest that in the pH-range where the corresponding complexes are formed the ~o-amino groups lose protons. Therefore, it can be concluded that ZnA +, ZnA(AH) ÷ and ZnA2 are really deprotonated species. This is to be seen in Fig. 6, where the concentration distributions for the Zn(II)-daba system are presented. Both the deprotonation processes taking place in a relatively low pH range and the high values of the stability constants suggest ~o-amino group coordination in the cases of dapa and daba. For dapa at p H - 9 . 5 a ZnA2H_1- mixed hydroxo complex is also formed (see Table 6). Further studies are necessary to enable us to draw conclusions about the donor groups taking part in the coordination of this complex.

n 2÷

AH)Z÷

1597

2*

8

"_o x

6

4

pH

Fig. 5. pH-dependent percentage concentration distributions of species formed in the Zn(II)-Orn system. CA= 8.10-3 M| Cz.at) = 2.10-3 M.

LO

x t~

8

zolA. .\

7olA.) " /----/,,

6

6

7

8

9

pH

Fig. 6. pH-dependent percentage concentration distributions of species formed in the Zn(II)-daba system. CA= 8.10-3 M; Cz,arj = 2.10-3 M. which illustrates the concentration distributions for the Zn(II)--Orn complexes. On the other hand, in the cases of dapa and daba the

REFERENCES

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