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Potentiometric, spectroscopic and thermal studies on the metal chelates of 1-(2-thiazolylazo)-2-naphthalenol
Spectrochimica Acta Part A 61 (2005) 929–936
Potentiometric, spectroscopic and thermal studies on the metal chelates of 1-(2-thiazolylazo)-2-naphthalenol M.M. Omar, Gehad G. Mohamed∗ Chemistry Department, Faculty of Science, Cairo University, Giza, Egypt Received 31 March 2004; accepted 18 May 2004
Abstract The synthesis and characterization of Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), Cd(II), Pd(II) and UO2 (II) chelates of 1-(2-thiazolylazo)2-naphthalenol (TAN) were reported. The dissociation constants of the ligand and the stability constants of the metal complexes were calculated pH-metrically at 25 ◦ C and 0.1 M ionic strength. The solid complexes were characterized by elemental and thermal analyses, molar conductance, IR, magnetic and diffuse reflectance spectra. The complexes were found to have the formulae [M(L)2 ] for M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II); [M(L)X]·nH2 O for M = Cu(II) (X = AcO, n = 3), Pd(II) (X = Cl, n = 0) and UO2 (II) (X = NO3 , n = 0), and [Fe(L)Cl2 (H2 O)]·2H2 O. The molar conductance data reveal that the chelates are non-electrolytes. IR spectra show that the ligand is coordinated to the metal ions in a terdentate manner with ONN donor sites of the naphthyl OH, azo N and thiazole N. An octahedral structure is proposed for Mn(II), Fe(III), Co(II), Ni(II), Zn(II), Cd(II) and UO2 (II) complexes and a square planar structure for Cu(II) and Pd(II) complexes. The thermal behaviour of these chelates shows that water molecules (coordinated and hydrated) and anions are removed in two successive steps followed immediately by decomposition of the ligand molecule in the subsequent steps. The relative thermal stability of the chelates is evaluated. The final decomposition products are found to be the corresponding metal oxides. The thermodynamic activation parameters, such as E∗ , H∗ , S∗ and G∗ are calculated from the TG curves. © 2004 Elsevier B.V. All rights reserved. Keywords: Azo complexes; Stability constants; IR; Conductance; Solid reflectance; Magnetic moment; Thermal analysis
1. Introduction Azo compounds are known to be involved in a number of biological reactions such as inhibition of DNA, RNA and protein synthesis, carcinogenesis and nitrogen fixation [1,2]. Furthermore, they were proved to have biological activity against bacteria and fungi [3,4]. Metal complexes of azo compounds containing heteroaryl ring systems find various applications [5]. Metal complexes of a series of heterocyclic azo compounds, prepared by coupling diazotized 2-aminothiazole with 1,3-dicarbonyl compounds [5], thiouracil [6], thymine [7] and substituted phenolic compounds [8], had been reported. TAN was widely used as a spectrophotometric reagent for the determination of Pd(II) and Co(II) [9], UO2 (II) [10], ∗
Corresponding author. Fax: +002-02-5727556. E-mail address: [email protected] (G.G. Mohamed).
1386-1425/$ – see front matter © 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2004.05.040
and Cu(II) and Zn(II) ions [11,12]. It could be, also, used in the extraction [13] and separation [14–16] of metal ions. Co(II), Ni(II) and Mn(II) ions were determined in water samples via flame atomic absorption spectrometry using TAN as an extracting agent [17,18]. In addition, flow-injection solidphase spectrophotometric technique was used for the determination of Zn(II) and Ni(II) ions using TAN as a chelating agent [19,20]. The conditional stability constants of the interaction of Cu(II), Zn(II) and Co(II) with TAN adsorbed on silica gel had been reported [21]. In continuation of our interest in azo ligands and their metal chelates [22–28], the main target of the present study is to synthesize new azo-metal chelates and determine the coordination capacity of the highly coloured HL (Fig. 1), that incorporates several binding sites. The coordination behaviour of HL towards transition metal ions is investigated by IR, molar conductance, magnetic moment and solid reflectance
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protons attached per ligand ion), n¯ (the average number of ligands attached per metal ion) and pL (the free ligand exponent) values were calculated using equations described by Irving and Rossotti [31]. The method used for calculating successive stability constants was interpolation at half n¯ value [30]. Fig. 1. Structure of the ligand (HL).
spectra measurements. The thermal decomposition of the complexes was also used to infer the structure and the various thermodynamic activation parameters are calculated.
2. Experimental 2.1. Materials and methods All chemicals used were of highest available purity. They include 2-aminothiazole (Sigma), 2-naphthol (Aldrich), sodium nitrite, chloride and hydroxide (BDH), ethyl alcohol (Adwic) and the disodium salt of ethylenediaminetetraacetic acid (Adwic). Cupric acetate dihydrate (Prolabo), uranyl nitrate dihydrate, cobalt and nickel chloride hexahydrates (BDH), zinc, palladium, manganese and cadmium chlorides (Adwic) and ferric chloride hexahydrate (Riedel-de Ha¨en) were used as received. The elemental analyses (C, H, N and S) were made at the Microanalytical Center at Cairo University. IR spectra were recorded on a Perkin-Elmer FTIR type 1650 spectrophotometer. The spectra were recorded as KBr discs. The solid reflectance spectra were measured on a Shimadzu 3101 PC spectrophotometer. The molar magnetic susceptibilities were measured on powdered samples using the Faraday method. The molar conductance measurements were carried out using a Sybron–Barnstead conductometer. A Shimadzu TGA-50H thermal analyzer was used to record simultaneously TG and DTG curves. The measurements were carried out in a dynamic nitrogen atmosphere (20 mL min−1 ) with a heating rate of 10 ◦ C min−1 in the temperature range 20–1200 ◦ C using platinum crucibles. The sample sizes ranged in mass from 1.2 to 3.5 mg. Highly sintered ␣-Al2 O3 was used as a reference. Metal contents were determined by titration against standard EDTA after complete decomposition of the complexes with aqua regia in a Kjeldahl flask several times. Three mixtures were prepared and potentiometrically titrated as previously described [29–31] using digital pHmeter (Orion Research Model 701 A Digital Ionalizer) to determine the proton–ligand and metal–ligand stability constants. Temperature was maintained constant with the help of Techne Circulator C-100 (UK) having accuracy of ±0.1 ◦ C. The pH-meter and electrode were calibrated by using standard buffer solutions, prepared according to NBS specifications [32]. Appropriate correction for converting pH-meter reading in non-aqueous medium was made [33]. The curves were plotted accordingly. The n¯ A (the average number of
2.2. Synthesis of 1-(2-thiazolylazo)-2-naphthol (TAN) 2-Aminothiazole (4.11 g, 10 mmol) was mixed with HCl (11.5 M, 5 mL) and diazotized below 5 ◦ C with NaNO2 (2.07 g, 10 mmol). The resulting diazonium chloride was coupled with an alcoholic NaOH solution (3 g, 25 mL) of 2naphthol (4.53 g, 10 mmol) below 5 ◦ C. The formed solid product was separated by filtration, purified by crystallization from ethanol, washed with diethyl ether and dried in a vacuum over anhydrous calcium chloride. The red azo product is produced in 80% yield (6.92 g). 2.3. Synthesis of metal complexes The metal complexes of HL were prepared by the addition of hot solution (60 ◦ C) of the appropriate metal chloride, nitrate or acetate (1 mmol) in an ethanol–water mixture (1:1, 25 mL) to the hot solution (60 ◦ C) of the azo compound (0.299 g, 2 mmol) in the same solvent (50 mL). The resulting mixture was stirred under reflux for 1 h whereupon the complexes precipitated. They were collected by filtration, washed with a 1:1 ethanol–water mixture and diethyl ether. The analytical data for C, H, N and S were repeated twice.
3. Results and discussion 3.1. Proton–ligand stability constant From the titration curves, it is observed that the ligand curve starts deviating from the free acid curve at about pH 7.5. The deviation increased continuously up to pH 10. It indicated that OH group starts to dissociate at about pH 7.5 and completes its dissociation at about pH 10.0. The proton–ligand formation numbers (¯nA ) were calculated by Irving and Rossotti’s expression [31]. The proton–ligand formation curve was obtained by plotting n¯ A values against pH (Fig. 2). The formation curve was found between 0 and 1. This indicates that the ligand have one dissociable proton from OH group. The pKa value was estimated from formation curve by noting the pH at which n¯ A = 0.5 and was found to be 8.75. The free energy change, G◦ , was found to be 49.84 kJ mol−1 . This positive value indicates the nonspontaneous character of dissociation reaction. 3.2. Metal–ligand stability constants The formation of chelate between the investigated metal ions and the ligand was indicated by: (i) the significant
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Table 1 Metal–ligand stability constants at 25 ◦ C Metal ion
log K1
log K2
log β2
log K1 − log K2
log K1 /log K2
−G◦ (kJ mol−1 deg−1 )
Mn(II) Fe(III) Co(II) Ni(II) Cu(II) Zn(II)
5.40 5.60 5.85 5.90 6.15 5.70
4.12 4.20 4.35 4.50 4.66 4.25
9.52 9.80 10.20 10.40 10.81 9.95
1.28 1.45 1.50 1.40 1.49 1.45
1.31 1.33 1.35 1.31 1.31 1.34
53.18 54.75 56.98 58.10 54.37 54.22
departure, starting from pH 2.5 to 3.0, of the metal complex titration curves from the ligand curves; and (ii) the change in colour from yellow to red violet as pH was raised from 2.5 to 8.0. The values of log K1 , log K2 and log β2 were directly read from the formation curves (¯n versus pL) (Fig. 3). The corresponding values are given in Table 1. The order of stability constants was found to be Cu(II) > Ni(II) > Co(II) > Zn(II) > Mn(II) in accordance with the Irving and Williams order [34] for divalent metal ions of the 3d series. One would expect a bigger difference between log K1 and log K2 values in such a ligand because of possible steric hindrance to the linking of the second ligand to the metal ion. The small difference may be due to trans-structure. The results show that the ratio log K1 /log K2 is positive in all cases. The free energy of formation (G◦ ) accompanying the complexation reaction has been determined at 25 ◦ C. The results are given in Table 1. The negative values of G◦ show that the driving tendency of the complexation reaction is from left to right and the reaction proceeds spontaneously. It is clear from Table 1 that the stability of Cu(II) complex are considerably larger as compared to other metals of the 3d series. Under the influence of the ligand field Cu(II) (3d9 ) will receive some extra stabilization [35] due to tetragonal distortion of octahedral symmetry in their complexes. The Cu(II)
Fig. 2. Proton–ligand formation curve.
complexes will be further stabilized due to the Jahn–Tellar effect [35]. 3.3. Solid metal complexes studies The general reaction for the preparation of the metal complexes of HL is shown below: MCl2 + 2HL → [M(L)2 ] + 2HCl
(1)
where M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II). MX2 + HL + nH2 O → [M(L)X] · nH2 O + HX
(2)
where M = Cu(II) (X = AcO, n = 3), Pd(II) (X = Cl, n = 0) and UO2 (II) (X = NO3 , n = 0). FeCl3 + HL + 3H2 O → [Fe(L)Cl2 (H2 O)] · 2H2 O + HCl (3)
Fig. 3. Metal–ligand formation curves.
M.M. Omar, G.G. Mohamed / Spectrochimica Acta Part A 61 (2005) 929–936
S
– 9.83 (9.57) 12.53 (12.87) 10.22 (10.44) 10.75 (10.44) 14.62 (14.75) 11.56 (11.38) 17.89 (18.12) – – 12.85 (12.55) 11.63 (11.35) 7.00 (7.36) 11.59 (11.33) 11.15 (11.33) 7.05 (7.43) 11.52 (11.21) 10.48 (10.36) 8.32 (8.09) 5.00 (5.46) 16.03 (16.47) 14.52 (14.89) 9.93 (9.66) 15.10 (14.87) 14.62 (14.87) 9.43 (9.76) 14.95 (14.71) 13.20 (13.59) 10.87 (10.62) 9.68 (9.56)
N H
3.72 (3.50) 2.74 (2.84) 3.80 (3.45) 2.80 (2.48) 2.84 (3.48) 4.29 (4.18) 2.49 (2.45) 2.01 (2.27) 2.62 (2.28) 2.00 (1.54) Red (85) Brown (69) Brown (75) Brown (65) Brown (62) Brown (65) Red (71) Brown (70) Red (61) Red (59) HL (L = C13 H9 N3 OS) [Mn(L)2 ] (L = C26 H16 N6 O2 S2 ) [Fe(L)Cl2 (H2 O)]·2H2 O (L = C26 H12 Cl2 FeN3 O4 S) [Co(L)2 ] (L = C26 H16 CoN6 O2 S2 ) [Ni(L)2 ] (L = C26 H16 N6 NiO2 S2 ) [Cu(L)(AcO)]·3H2 O (L = C15 H17 CuN3 O3 S) [Zn(L)2 ] (L = C26 H16 N6 O2 S2 Zn) [Cd(L)2 ] (L = C26 H16 CdN6 O2 S2 ) [Pd(L)Cl] (L = C13 H8 ClN3 O4 PdS) [UO2 (L)(NO3 )] (L = C13 H8 N4 O4 SUO2 )
C
Found (%, calculated) mp (◦ C) Colour (% yield) Compound
Table 2 Analytical and physical data of HL and its metal complexes
A detailed interpretation of the IR spectra of HL and the effect of binding with Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), Cd(II), Pd(II) and UO2 (II) ions on the vibrational frequencies of the free HL ligand are discussed. The IR spectra of the free ligand and its metal chelates were carried out in the 4000–400 cm−1 range (Table 3). The IR spectrum of the ligand shows a broad band at 3500–3050 cm−1 , which can be attributed to the phenolic OH group. This band is still broad in all complexes, which renders it difficult to attribute to the involvement of phenolic OH group in coordination. The involvement of the deprotonated phenolic OH group in chelation is confirmed by the blue-shift of the (C O) stretching band, observed at 1215 cm−1 in the free ligand, to the extent of 5–16 cm−1 in the complexes [24]. Also the participation of the OH group is apparent from the disappearance (in UO2 (II) complex only) or shift in position of the ␦(OH) in-plane bending at 1409 cm−1 in the free ligand by 6–34 cm−1 in the remaining complexes [36]. However, the (N N) stretching band in the free ligand is observed at 1579 cm−1 [22,24]. This band is shifted to higher (6–23 cm−1 ) or lower (34–63 cm−1 ) frequency values upon complexation suggesting coordination via the azo group (M ← N) [24]. The IR spectrum of the ligand revealed a medium band at 1629 cm−1 due to (C N) of the N3 thiazole nitrogen. This band is shifted to higher (27 cm−1 ) or lower (2–35 cm−1 ) frequencies in the complexes indicating that it has been affected upon coordination to the metal ions [24]. The band at 744 cm−1 in the ligand is still in the same position in the complexes indicating the non-involvement of the thiazole S in coordination. In the far-IR spectra of all complexes, the non-ligand bands observed at 422–472 and 473–506 cm−1 region can be assigned to the (M N) stretching vibrations
60.92 (61.17) 55.20 (55.32) 35.60 (35.86) 55.60 (55.22) 55.30 (55.22) 41.61 (41.81) 54.75 (54.64) 50.95 (50.49) 39.83 (39.44) 26.33 (26.62)
3.4. IR spectra and mode of bonding
165 ± 2 >300 >300 >300 >300 >300 >300 >300 >300 >300
M
– 5.30 5.37 5.15 3.0 1.87 Diam. Diam. Diam. Diam.
µeff (B.M.)
The results of the elemental analyses of the metal chelates of TAN, which are recorded in Table 2, are in good agreement with those required by the proposed formulae. In most cases, 1:2 (M:L) solid complexes were isolated and found to have the general formula [M(L)2 ] where M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II). While the Cu(II), Pd(II) and UO2 (II) chelates are of the type 1:1 (M:L) and have the formula [M(L)X]·nH2 O, where X = AcO in case of Cu (n = 3), X = Cl in case of Pd(II) (n = 0) and X = NO3 in case of UO2 (II) (n = 0). In addition, Fe(III) forms a 1:1 (M:L) chelate of the formula [Fe(L)Cl2 (H2 O)]·2H2 O. The HL ligand offers several alternatives to coordination to metals. The results shown in Table 2 prove that HL is indeed monoanionic where the OH proton of 2-naphthol is lost when coordinated to the metal ions, and thus reducing the charge on the complex ion. The presence of Cl, NO3 or AcO ions in the inner sphere of the chelates neutralizes the charge on the complex. Thus, HL is expected to fill three coordination positions and contribute a charge of −1. The non-involvement of Cl, NO3 or AcO ions in the coordination to the metal ions is also confirmed by the molar conductance measurements.
– 14.20 10.25 8.75 11.98 9.32 12.85 13.65 15.63 10.93
Λm (−1 cm2 mol−1 )
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Table 3 IR spectra (4000–400 cm−1 ) of HL and its metal complexes Compound
(C N)
(N N)
(C O)
␦(OH)
(H2 O)
(M O)
(M N) (thiazole)
(M N) (azo)
HL [Mn(L)2 ] [Fe(L)Cl2 (H2 O)]·2H2 O [Co(L)2 ] [Ni(L)2 ] [Cu(L)(AcO)]·3H2 O [Zn(L)2 ] [Cd(L)2 ] [Pd(L)Cl] [UO2 (L)(NO3 )]
1629m 1617s 1656s 1627s 1610sh 1594m 1612m 1613m 1600br 1627s
1579s 1594s 1602w 1585m 1525sh 1525sh 1545m 1516m 1545w 1591s
1215m 1199sh 1202sh 1204sh 1200sh 1203sh 1202sh 1201sh 1202sh 1210m
1409m 1391s 1403sh 1391m 1396s 1375m 1376m 1370m 1368m Dis.
– – 816sh, 47sh – – – – – – –
– 508s 503s 504w 528s 508s 525s 572w 565w 509w
– 437s 473m 473m 473w 485w 500w 514s 506w 474s
– 473w 422w 424w 436s 441w 439m 437m 472w 450w
of the azo and N3 thiazole nitrogen, respectively [6]. Conclusive evidence regarding the bonding of oxygen to the metal ions is provided by the occurrence of bands at 503–565 cm−1 as the result of (M O) [6,24]. Therefore, the IR spectra indicate that HL behaves as monobasic acid and the coordination sites being Ar OH, N N and the N3 atom of the thiazole moiety. 3.5. Molar conductance data The solubility of the complexes in DMF permitted determination of the molar conductivity (Λm ) of 10−3 M solutions at 25 ◦ C and, by comparison, the electrolytic nature for each complex. The low values of the molar conductance data listed in Table 2 indicate that the complexes are nonelectrolytes. 3.6. Magnetic susceptibility and electronic spectra measurements The diffuse reflectance spectra of Mn(II), Fe(III), Co(II), Ni(II) and Cu(II) chelates show three bands at 210–215, 245–254 and 300–311 nm which are attributed to the → ∗ and n → ∗ transitions, respectively, within the HL ligand. The solid reflectance spectrum of Mn(II) complex show three bands at 16,528, 25,000 and 28,571 cm−1 (Table 4) assignable to 4 T1g → 6 A1g , 4 T2g (G) → 6 A1g and 4 T1g (D) → 6 A1g transitions [37], respectively, which lie in the same range as reported for octahedrally coordinated Mn(II) ion. The magnetic moment (5.30 B.M.) is an additional evidence for an octahedral structure [37]. From the diffuse reflectance spectra, it is observed that, the Fe(III) chelate exhibits a band at 22,222 cm−1 , which may be assigned to the 6 A1g → 5 T2g (G) transition in octahedral geometry of the complex [22]. The 6 A1g → 5 T1g transition appears to be split into two bands at 12,594 and 17,482 cm−1 . The observed magnetic moment of Fe(III) complex is 5.37 B.M. Thus, the complex formed has octahedral geometry around the Fe(III) ion [22]. The band observed at 26,954 cm−1 can be attributed to ligand-to-metal charge transfer band.
The diffuse reflectance spectra of Co(II) complex give three bands at 12,820, 15,552 and 17,574 cm−1 . The third region at 25,974 cm−1 refers to the charge transfer band (L → MCT). The bands observed are assigned to the transitions 4 T1g (F) → 4 T2g (F) (ν1 ), 4 T1g (F) → 4 A2g (F) (ν2 ) and 4 T (F) → 4 T (P) (ν ), respectively, suggesting that there is 1g 2g 3 an octahedral geometry around Co(II) ion [38,39]. The magnetic susceptibility measurement (5.15 B.M.) is an indicative of octahedral geometry [39]. The Ni(II) complex reported herein is high spin with a room temperature magnetic moment value of 3.0 B.M.; which are in the normal range observed for octahedral Ni(II) complexes [40,41]. The diffused reflectance spectrum of the Ni(II) complex displays three bands at 12,626 cm−1 (ν1 ), 15,625 cm−1 (ν2 ) and 21,691 cm−1 (ν3 ) assigned to 3 A2g → 3 T2g , 3 A2g → 3 T1g (F) and 3 A2g → 3 T (P) transitions, respectively. The band at 25,839 cm−1 1g refers to the charge transfer band (L → MCT). The electronic spectrum of Cu(II) complex shows bands around 15,873, 17,482 and 21,739 cm−1 . The position of the
Table 4 Electronic bands and d–d transitions of the metal complexes Compound
Band position (cm−1 )
Assignment
[Mn(L)2 ]
16,528 25,000 28,571
4T
1g (G)
4T
2g (G)
12,594 17,482 22,222 26,954
6A
6 A → 5 T (G) 1g 2g Charge transfer
12,820 15,552 17,574 25,974
→ 4 T2g (F) → 4 A2g (F) 4 T (F) → 4 T (F) 1g 2g Charge transfer
12,626 15,625 21,691 25,839
→ 3 T2g → 3 T1g (F) 3 A → 3 T (P) 2g 1g Charge transfer
15,873 17,482 21,739
2A
[Fe(L)Cl2 (H2 O)]·2H2 O
[Co(L)2 ]
[Ni(L)2 ]
[Cu(L)(AcO)]·3H2 O
→ 6 A1g → 6 A1g 4 T (D) → 6 A 2g 1g 1g
→ 5 T1g
4T
1g (F)
4T
1g (F)
3A
2g
3A
2g
2E
1g
g
→ 2 B1g
→ 2 B1g
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bands together with its magnetic moment value (1.87 B.M.) indicates that this complex has a square planar stereochemistry [26]. The Zn(II) and Cd(II) chelates are diamagnetic and in analogy with the previously described Zn(II) and Cd(II) complexes containing ONO donor azo ligands [26] and according to the empirical formulae of these complexes, we propose an octahedral geometry for the Zn(II) and Cd(II) chelates. The diamagnetic nature of the UO2 (II) and Pd(II) complexes indicates octahedral and square planar geometry, respectively, as expected for UO2 (II) and Pd(II) metal ions [37]. 3.7. Thermogravimetric (TG) analysis Thermal data of the complexes are given in Table 5. The correlations between the different decomposition steps of the complexes with the corresponding weight losses are discussed in terms of the proposed formulae of the complexes. The Mn(II) complex with the general formula [Mn(C26 H16 N6 O2 S2 )] is thermally decomposed in three successive decomposition steps. The first estimated mass loss of 20.03% within the temperature range 80–220 ◦ C may be attributed to the loss of C3 H2 N3 S molecule (calculated mass loss = 19.86%). The energy of activation of this step is 45.37 kJ mol−1 . The second and third steps occur within the temperature range 220–950 ◦ C with an estimated mass loss 64.84% (calculated mass loss = 64.54%) are reasonably accounted for the decomposition of the remaining ligand molecules leaving MnS residue, where the activation energies are 44.19 and 39.43 kJ mol−1 for the second and third steps, respectively. The total estimated mass loss 90.69% (total calculated mass loss = 91.32%). The Fe(III) complex, [Fe(C13 H8 N3 OS)Cl2 (H2 O)]·2H2 O, is thermally decomposed in four successive decomposition steps within the temperature range 40–1100 ◦ C. The first decomposition step of estimated mass loss 8.34% within the temperature range 40–140 ◦ C may be attributed to the liberation of two hydrated water molecules (calculated mass loss = 8.28%). The energy of activation was 35.45 kJ mol−1 . The three steps found within the temperature range 160–1100 ◦ C with an estimated mass loss 73.84% (calculated mass loss = 73.33%) which is responsibly accounted for the decomposition of one coordinated water, 2HCl and the ligand molecules with a final oxide residue of (1/2)Fe2 O3 . The activation energies are found to be 63.01, 205 and 370 kJmol−1 for the second, third and fourth steps, respectively. The Cu(II) complex with the general formula [Cu(C13 H8 N3 OS)(AcO)]·3H2 O was thermally decomposed in four successive decomposition steps. The first estimated mass loss of 11.95% within the temperature range 40–130 ◦ C may be attributed to the liberation of three molecules of hydrated water (calculated mass loss = 12.54%). The energy of activation of this step is 71.80 kJ mol−1 . The
remaining steps of decompositions occur within the temperature range 130–1100 ◦ C with an estimated mass loss of 70.20% (calculated mass loss = 68.99%) and activation energies of 175, 244 and 244 kJ mol−1 , which corresponds to the loss of acetate (as CH4 and CO2 gases), and ligand molecule leaving Cu2 O residue with a total estimated mass loss 82.15% (total calculated mass loss = 81.53%). The Zn(II) complex, with the general formula [Zn(C26 H16 N6 O2 S2 )], shows decomposition pattern of four stages. The first and second steps with estimated mass loss of 44.69%, found within the temperature range 40–450 ◦ C corresponding to loss of one ligand molecule; C13 H8 N3 OS (calculated mass loss = 44.48%). The activation energies of these two steps are 76.95 and 158 kJ mol−1 , respectively. The remaining decomposition steps with an estimated mass loss 42.14% which is due to loss of the second ligand molecule leaving ZnO residue occurring within the temperature range 500–1100 ◦ C (calculated mass loss = 41.68%). The sum of the activation energy for the last two decomposition steps are 600 kJ mol−1 . The total estimated mass loss is 86.83% (total calculated mass loss = 86.16%). The uranyl complex with the formula [UO2 (C13 H8 N3 OS) (NO3 )] was thermally decomposed in six successive decomposition steps. The first estimated mass loss of 10.80% (calculated mass loss = 10.58%) within the temperature range 30–100 ◦ C can be attributed to the liberation of nitrate group as NO2 and (1/2)O2 gases. The energy of activation of this step is 19.35 kJ mol−1 . The second to sixth steps occur within the temperature range 110–850 ◦ C with an estimated mass loss 43.18% (calculated mass loss = 43.34%) with an energy of activation sum to 306.61 kJ mol−1 , which is reasonably accounted for the loss of the ligand molecule leaving UO2 as residue with total estimated mass loss 53.98% (total calculated mass loss = 53.92%). 3.8. Kinetic analysis The first stage of dehydration of the complexes was studied in more detail. The kinetic parameters such as activation energy (E∗ ), enthalpy (H∗ ), entropy (S∗ ) and free energy change of decomposition (G∗ ) were evaluated graphically by employing the Coats–Redfern relation [42]: AR log{Wf /(Wf − W)} 2RT = log log 1 − T2 θE∗ E∗ −
E∗ 2.303RT
(4)
where Wf is the mass loss at the completion of the reaction, W is the mass loss up to the temperature T, R is the gas constant, E∗ is the activation energy in kJ mol−1 , θ is the heating rate and (1 − (2RT/E∗ )) ∼ = 1. A plot of the left-hand side of Eq. (4) against 1/T gives a slope from which E∗ was calculated and A (Arrhenius constant) was determined from the intercept. The entropy of activation (S∗ ), enthalpy of activation (H∗ ) and the free energy change of activation (G∗ ) were calculated
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Table 5 Thermoanalytical results (TG, DTG) of metal complexes Complex
TG range (◦ C)
DTGmax (◦ C)
Estimated (%, calculated) Mass loss
Assignment
Metallic residue
MnS
Total mass loss
[Mn(L)2 ]
80–220 220–520 520–950
192 300 650
20.03 (19.86) 18.88 (18.44) 45.96 (46.10)
84.88 (84.40)
Loss of C3 H2 N3 S molecule Loss of C8 H8 molecule Loss of C15 H6 N3 O2 molecule
[Fe(L)Cl2 (H2 O)]·2H2 O
40–140 160–110
75 208, 520, 780
8.34 (8.28) 73.84 (73.33)
82.18 (81.61)
Loss of 2H2 O (coord.) Loss of 2HCl and ligand molecules
(1/2)Fe2 O3
40–130 150–1100
65 198, 278, 750
11.95 (12.54) 70.20 (68.99)
82.15 (81.53)
Loss of 3H2 O (hyd.) Loss of CH4 , CO2 and ligand molecule
Cu2 O
40–450 600–1100
70, 275 670, 850
44.69 (44.48) 42.14 (41.68)
86.83 (86.16)
Loss of C13 H8 N3 OS molecule Loss of C13 H8 N3 S molecule
ZnO
30–100 110–850
46 132, 220, 347, 629, 784
10.80 (10.58) 43.18 (43.34)
53.98 (53.92)
Loss of NO2 and (1/2)O2 Loss of ligand molecule
UO2
[Cu(L)(AcO)]·3H2 O [Zn(L)2 ] [UO2 (L)(NO3 )]
using the following equations: Ah ∗ S = 2.303 log R kT
(5)
According to the kinetic data obtained from the DTG curves, all the complexes have a negative entropy, which indicates that the complexes are formed spontaneously.
H ∗ = E∗ − RT
(6)
3.9. Structural interpretation of solid complexes
G∗ = H ∗ − TS ∗
(7)
The design and synthesis of a new tridentate azo ligand derived from 2-naphthol and 2-aminothiazole for use in square planar or octahedral molecular templates have been successfully demonstrated. The synthesis of the ligand and its complexes proved to be as straightforward as expected, giving high yields of the free ligand and its complexes in simple, onepot reactions. As anticipated, the ligand coordinates equatorially to four- or six-coordinate transition metal ions to give
where k and h are the Boltzman and Plank constants, respectively. The calculated values of E∗ , A, S∗ , H∗ and G∗ for the decomposition steps are given in Table 6. The correlation coefficients of the Arrhenius plots of the thermal decomposition steps were found to lie in the range 0.92–0.97, showing a good fit with the linear function. Table 6 Thermodynamic data of the thermal decomposition of metal complexes Complex
Decomposition temperature (K)
E∗ (kJ mol−1 )
A (s−1 )
S∗ (kJ mol−1 )
H∗ (kJ mol−1 )
G∗ (kJ mol−1 )
80–220 220–520 520–950
45.37 44.19 39.43
2.59 × 105 5.45 × 106 6.11 × 1012
−138.0 −116.0 −184.0
43.78 41.69 38.89
70.20 76.50 76.50
[Fe(L)Cl2 (H2 O)]·2H2 O
40–140 160–250 500–600 750–1100
35.45 63.01 205.0 370.0
1.20 × 105 9.03 × 105 3.99 × 1012 3.88 × 1013
−136.0 −128.0 −83.0 −104.0
34.82 61.37 200.0 250.0
45.01 87.97 205.0 306.0
[Cu(L)(AcO)]·3H2 O
40–130 150–230 250–400 650–1100
71.80 175.0 244.0 244.0
3.21 × 1010 6.98 × 1018 6.89 × 1021 2.31 × 1011
−31.13 −119.0 −174.0 −35.30
71.23 174.0 242.0 238.0
73.28 150.0 194.0 264.0
[Zn(L)2 ]
40–100 200–450 600–700 720–1100
76.95 158.0 212.0 388.0
1.09 × 1011 4.07 × 1014 7.47 × 1010 5.75 × 1018
−21.59 −35.44 −43.49 −106.0
76.37 156.0 207.0 381.0
77.88 146.0 235.0 291.0
30–100 110–180 200–250 330–380 600–660 760–850
19.35 22.92 40.53 48.32 79.84 115.0
4.00 × 106 8.55 × 106 6.22 × 105 1.58 × 106 1.40 × 106 1.82 × 105
−103.0 −105.0 −131.0 −128.0 −134.0 −152.0
18.97 21.82 38.70 45.44 74.61 109.0
23.71 35.74 67.62 89.69 159.0 228.0
[Mn(L)2 ]
[UO2 (L)(NO3 )]
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Fig. 4. Suggested structural formulae of metal complexes.
square planar (Cu(II) and Pd(II)), and octahedral (Mn(II), Fe(III), Co(II), Ni(II), Zn(II), Cd(II) and UO2 (II)) environments, around the metal ion anchor. The proposed general structures of the complexes are shown in Fig. 4. There are additional factors which might affect the structure of these chelates like the possibility of various isomeric structures, bulkiness of the ligand and intermolecular hydrogen bonds, however, but due to instrument limitation, we are not able to explore this further.
References [1] S. Patai, The Chemistry of Hydrazo, Azo and Azoxy Groups. Part 1, Wiley, New York, 1975. [2] A.S.A. Zidan, A.I. El-Said, M.S. El-Meligy, A.A.M. Aly, O.F. Mohammed, J. Therm. Anal. 62 (2000) 665. [3] I.M.A. Awad, Phosphorus Sulfur Silicon 114 (1996) 17. [4] I.M.A. Awad, J. Chem. Technol. Biotechnol. 53 (1992) 227. [5] K. Krishnankutty, D.K. Babu, J. Ind. Chem. Soc. 73 (1996) 379, and references therein. [6] G.G. Mohamed, Spectrochim. Acta Part A 57 (2001) 411. [7] Z.M. Zaki, Spectrochim. Acta Part A 56 (2000) 1917. [8] P.P. Hankare, S.S. Chavan, Indian J. Chem. Sect. A 42 (2003) 540. [9] V.M. Ivanov, O.V. Kuznetsova, O.V. Grineva, J. Anal. Chem. 54 (1999) 233. [10] G.V. Rathaiah, J.K. Charyulu, M.C. Eshwar, J. Radioanal. Nucl. Chem. 99 (1986) 337.
[11] M.H. Gandhi, K.C. Pathak, D.J. Parikh, S.A. Lahiri, Indian J. Chem. Soc. Sect. A 25 (1986) 499. [12] O. Zaporozhets, N. Petruniock, O. Bessarabova, V. Sukhan, Talanta 49 (1999) 899. [13] A.E. Visser, R.P. Swatloski, S.T. Griffin, D.H. Hartman, R.D. Rogers, Sep. Sci. Technol. 36 (2001) 785. [14] E.M. Basova, T.A. Bolshova, V.M. Ivanov, O.A. Shpigun, Zh. Anal. Khim. 48 (1993) 284. [15] J. Peng, C. Deng, S. Liu, Fenxi Huaxue 26 (1998) 303. [16] W. Lee, S.E. Lee, C.H. Lee, Y.S. Kim, Y.I. Lee, Microchem. J. 70 (2001) 195. [17] J. Chen, K.C. Teo, Anal. Chim. Acta 434 (2001) 325. [18] K.C. Teo, J. Chen, Analyst 126 (2001) 534. [19] L.S.G. Teixeira, F.R.P. Rocha, M. Korn, B.F. Reis, S.L.G. Ferreira, A.C.S. Costa, Anal. Chim. Acta 383 (1999) 309. [20] L.S.G. Teixeira, F.R.P. Rocha, M. Korn, B.F. Reis, S.L.G. Ferreira, A.C.S. Costa, Talanta 51 (2000) 1027. [21] O.A. Zaporozhets, N.I. Smyk, V.V. Sykhan, Funct. Mater. 8 (2001) 566. [22] G.G. Mohamed, N.E.A. El-Gamel, F.A. Nour El-Dien, Synth. React. Inorg. Met. Org. Chem. 31 (2001) 347. [23] G.G. Mohamed, F.A. Nour El-Dien, N.E.A. El-Gamel, J. Therm. Anal. 67 (2002) 135. [24] G.G. Mohamed, N.E.A. El-Gamel, F. Teixidor, Polyhedron 20 (2001) 2689. [25] G.G. Mohamed, M.A. Zayed, N.E.A. El-Gamel, Spectrochim. Acta Part A 58 (2002) 3167. [26] G.G. Mohamed, Z.M. Zaki, Synth. React. Inorg. Met. Org. Chem., in press. [27] M.M. Omar, Ann. Chim. 92 (2002) 601. [28] M.M. Omar, J. Therm. Anal. 68 (2002) 1003. [29] M. Calvin, K.W. Wilson, J. Am. Chem. Soc. 67 (1945) 2003. [30] H. Irving, H.S. Rossotti, J. Chem. Soc. (1953) 3397. [31] H. Irving, H.S. Rossotti, J. Chem. Soc. (1954) 2904. [32] R.G. Bates, Determination of pH, Wiley–Interscience, New York, 1973. [33] R.G. Bates, M. Paabo, R.A. Robinson, J. Phys. Chem. 34 (1972) 581. [34] (a) H. Irving, S.P. Williams, Nature (London) 162 (1948) 746; (b) H. Irving, S.P. Williams, J. Chem. Soc. (1953) 3192. [35] L.E. Orgel, An Introduction to Transition Metal Chemistry Ligand Field Theory, Methuen, London, 1966 (p. 55). [36] A.A. Soliman, J. Therm. Anal. 63 (2001) 221. [37] G.M. Abu El-Reash, K.M. Ibrahim, M.M. Bekheit, Bull. Soc. Chim. Fr. 128 (1991) 149. [38] G.G. Mohamed, Z.H. Abd El-Wahab, J. Therm. Anal. 73 (2003) 347. [39] N. Mondal, D.K. Dey, S. Mitra, K.M. Abdul Malik, Polyhedron 19 (2000) 2707. [40] F.A. Cotton, G. Wilkinson, C.A. Murillo, M. Bochmann, Advanced Inorganic Chemistry, sixth ed., Wiley, New York, 1999. [41] D.R. Zhu, Y. Song, Y. Xu, Y. Zhang, S.S.S. Raj, H.K. Fun, X.Z. You, Polyhedron 19 (2000) 2019. [42] A.W. Coats, J.P. Redfern, Nature 20 (1964) 68.
Potentiometric, spectroscopic and thermal studies on the metal chelates of 1-(2-thiazolylazo)-2-naphthalenol M.M. Omar, Gehad G. Mohamed∗ Chemistry Department, Faculty of Science, Cairo University, Giza, Egypt Received 31 March 2004; accepted 18 May 2004
Abstract The synthesis and characterization of Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), Cd(II), Pd(II) and UO2 (II) chelates of 1-(2-thiazolylazo)2-naphthalenol (TAN) were reported. The dissociation constants of the ligand and the stability constants of the metal complexes were calculated pH-metrically at 25 ◦ C and 0.1 M ionic strength. The solid complexes were characterized by elemental and thermal analyses, molar conductance, IR, magnetic and diffuse reflectance spectra. The complexes were found to have the formulae [M(L)2 ] for M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II); [M(L)X]·nH2 O for M = Cu(II) (X = AcO, n = 3), Pd(II) (X = Cl, n = 0) and UO2 (II) (X = NO3 , n = 0), and [Fe(L)Cl2 (H2 O)]·2H2 O. The molar conductance data reveal that the chelates are non-electrolytes. IR spectra show that the ligand is coordinated to the metal ions in a terdentate manner with ONN donor sites of the naphthyl OH, azo N and thiazole N. An octahedral structure is proposed for Mn(II), Fe(III), Co(II), Ni(II), Zn(II), Cd(II) and UO2 (II) complexes and a square planar structure for Cu(II) and Pd(II) complexes. The thermal behaviour of these chelates shows that water molecules (coordinated and hydrated) and anions are removed in two successive steps followed immediately by decomposition of the ligand molecule in the subsequent steps. The relative thermal stability of the chelates is evaluated. The final decomposition products are found to be the corresponding metal oxides. The thermodynamic activation parameters, such as E∗ , H∗ , S∗ and G∗ are calculated from the TG curves. © 2004 Elsevier B.V. All rights reserved. Keywords: Azo complexes; Stability constants; IR; Conductance; Solid reflectance; Magnetic moment; Thermal analysis
1. Introduction Azo compounds are known to be involved in a number of biological reactions such as inhibition of DNA, RNA and protein synthesis, carcinogenesis and nitrogen fixation [1,2]. Furthermore, they were proved to have biological activity against bacteria and fungi [3,4]. Metal complexes of azo compounds containing heteroaryl ring systems find various applications [5]. Metal complexes of a series of heterocyclic azo compounds, prepared by coupling diazotized 2-aminothiazole with 1,3-dicarbonyl compounds [5], thiouracil [6], thymine [7] and substituted phenolic compounds [8], had been reported. TAN was widely used as a spectrophotometric reagent for the determination of Pd(II) and Co(II) [9], UO2 (II) [10], ∗
Corresponding author. Fax: +002-02-5727556. E-mail address: [email protected] (G.G. Mohamed).
1386-1425/$ – see front matter © 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.saa.2004.05.040
and Cu(II) and Zn(II) ions [11,12]. It could be, also, used in the extraction [13] and separation [14–16] of metal ions. Co(II), Ni(II) and Mn(II) ions were determined in water samples via flame atomic absorption spectrometry using TAN as an extracting agent [17,18]. In addition, flow-injection solidphase spectrophotometric technique was used for the determination of Zn(II) and Ni(II) ions using TAN as a chelating agent [19,20]. The conditional stability constants of the interaction of Cu(II), Zn(II) and Co(II) with TAN adsorbed on silica gel had been reported [21]. In continuation of our interest in azo ligands and their metal chelates [22–28], the main target of the present study is to synthesize new azo-metal chelates and determine the coordination capacity of the highly coloured HL (Fig. 1), that incorporates several binding sites. The coordination behaviour of HL towards transition metal ions is investigated by IR, molar conductance, magnetic moment and solid reflectance
930
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protons attached per ligand ion), n¯ (the average number of ligands attached per metal ion) and pL (the free ligand exponent) values were calculated using equations described by Irving and Rossotti [31]. The method used for calculating successive stability constants was interpolation at half n¯ value [30]. Fig. 1. Structure of the ligand (HL).
spectra measurements. The thermal decomposition of the complexes was also used to infer the structure and the various thermodynamic activation parameters are calculated.
2. Experimental 2.1. Materials and methods All chemicals used were of highest available purity. They include 2-aminothiazole (Sigma), 2-naphthol (Aldrich), sodium nitrite, chloride and hydroxide (BDH), ethyl alcohol (Adwic) and the disodium salt of ethylenediaminetetraacetic acid (Adwic). Cupric acetate dihydrate (Prolabo), uranyl nitrate dihydrate, cobalt and nickel chloride hexahydrates (BDH), zinc, palladium, manganese and cadmium chlorides (Adwic) and ferric chloride hexahydrate (Riedel-de Ha¨en) were used as received. The elemental analyses (C, H, N and S) were made at the Microanalytical Center at Cairo University. IR spectra were recorded on a Perkin-Elmer FTIR type 1650 spectrophotometer. The spectra were recorded as KBr discs. The solid reflectance spectra were measured on a Shimadzu 3101 PC spectrophotometer. The molar magnetic susceptibilities were measured on powdered samples using the Faraday method. The molar conductance measurements were carried out using a Sybron–Barnstead conductometer. A Shimadzu TGA-50H thermal analyzer was used to record simultaneously TG and DTG curves. The measurements were carried out in a dynamic nitrogen atmosphere (20 mL min−1 ) with a heating rate of 10 ◦ C min−1 in the temperature range 20–1200 ◦ C using platinum crucibles. The sample sizes ranged in mass from 1.2 to 3.5 mg. Highly sintered ␣-Al2 O3 was used as a reference. Metal contents were determined by titration against standard EDTA after complete decomposition of the complexes with aqua regia in a Kjeldahl flask several times. Three mixtures were prepared and potentiometrically titrated as previously described [29–31] using digital pHmeter (Orion Research Model 701 A Digital Ionalizer) to determine the proton–ligand and metal–ligand stability constants. Temperature was maintained constant with the help of Techne Circulator C-100 (UK) having accuracy of ±0.1 ◦ C. The pH-meter and electrode were calibrated by using standard buffer solutions, prepared according to NBS specifications [32]. Appropriate correction for converting pH-meter reading in non-aqueous medium was made [33]. The curves were plotted accordingly. The n¯ A (the average number of
2.2. Synthesis of 1-(2-thiazolylazo)-2-naphthol (TAN) 2-Aminothiazole (4.11 g, 10 mmol) was mixed with HCl (11.5 M, 5 mL) and diazotized below 5 ◦ C with NaNO2 (2.07 g, 10 mmol). The resulting diazonium chloride was coupled with an alcoholic NaOH solution (3 g, 25 mL) of 2naphthol (4.53 g, 10 mmol) below 5 ◦ C. The formed solid product was separated by filtration, purified by crystallization from ethanol, washed with diethyl ether and dried in a vacuum over anhydrous calcium chloride. The red azo product is produced in 80% yield (6.92 g). 2.3. Synthesis of metal complexes The metal complexes of HL were prepared by the addition of hot solution (60 ◦ C) of the appropriate metal chloride, nitrate or acetate (1 mmol) in an ethanol–water mixture (1:1, 25 mL) to the hot solution (60 ◦ C) of the azo compound (0.299 g, 2 mmol) in the same solvent (50 mL). The resulting mixture was stirred under reflux for 1 h whereupon the complexes precipitated. They were collected by filtration, washed with a 1:1 ethanol–water mixture and diethyl ether. The analytical data for C, H, N and S were repeated twice.
3. Results and discussion 3.1. Proton–ligand stability constant From the titration curves, it is observed that the ligand curve starts deviating from the free acid curve at about pH 7.5. The deviation increased continuously up to pH 10. It indicated that OH group starts to dissociate at about pH 7.5 and completes its dissociation at about pH 10.0. The proton–ligand formation numbers (¯nA ) were calculated by Irving and Rossotti’s expression [31]. The proton–ligand formation curve was obtained by plotting n¯ A values against pH (Fig. 2). The formation curve was found between 0 and 1. This indicates that the ligand have one dissociable proton from OH group. The pKa value was estimated from formation curve by noting the pH at which n¯ A = 0.5 and was found to be 8.75. The free energy change, G◦ , was found to be 49.84 kJ mol−1 . This positive value indicates the nonspontaneous character of dissociation reaction. 3.2. Metal–ligand stability constants The formation of chelate between the investigated metal ions and the ligand was indicated by: (i) the significant
M.M. Omar, G.G. Mohamed / Spectrochimica Acta Part A 61 (2005) 929–936
931
Table 1 Metal–ligand stability constants at 25 ◦ C Metal ion
log K1
log K2
log β2
log K1 − log K2
log K1 /log K2
−G◦ (kJ mol−1 deg−1 )
Mn(II) Fe(III) Co(II) Ni(II) Cu(II) Zn(II)
5.40 5.60 5.85 5.90 6.15 5.70
4.12 4.20 4.35 4.50 4.66 4.25
9.52 9.80 10.20 10.40 10.81 9.95
1.28 1.45 1.50 1.40 1.49 1.45
1.31 1.33 1.35 1.31 1.31 1.34
53.18 54.75 56.98 58.10 54.37 54.22
departure, starting from pH 2.5 to 3.0, of the metal complex titration curves from the ligand curves; and (ii) the change in colour from yellow to red violet as pH was raised from 2.5 to 8.0. The values of log K1 , log K2 and log β2 were directly read from the formation curves (¯n versus pL) (Fig. 3). The corresponding values are given in Table 1. The order of stability constants was found to be Cu(II) > Ni(II) > Co(II) > Zn(II) > Mn(II) in accordance with the Irving and Williams order [34] for divalent metal ions of the 3d series. One would expect a bigger difference between log K1 and log K2 values in such a ligand because of possible steric hindrance to the linking of the second ligand to the metal ion. The small difference may be due to trans-structure. The results show that the ratio log K1 /log K2 is positive in all cases. The free energy of formation (G◦ ) accompanying the complexation reaction has been determined at 25 ◦ C. The results are given in Table 1. The negative values of G◦ show that the driving tendency of the complexation reaction is from left to right and the reaction proceeds spontaneously. It is clear from Table 1 that the stability of Cu(II) complex are considerably larger as compared to other metals of the 3d series. Under the influence of the ligand field Cu(II) (3d9 ) will receive some extra stabilization [35] due to tetragonal distortion of octahedral symmetry in their complexes. The Cu(II)
Fig. 2. Proton–ligand formation curve.
complexes will be further stabilized due to the Jahn–Tellar effect [35]. 3.3. Solid metal complexes studies The general reaction for the preparation of the metal complexes of HL is shown below: MCl2 + 2HL → [M(L)2 ] + 2HCl
(1)
where M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II). MX2 + HL + nH2 O → [M(L)X] · nH2 O + HX
(2)
where M = Cu(II) (X = AcO, n = 3), Pd(II) (X = Cl, n = 0) and UO2 (II) (X = NO3 , n = 0). FeCl3 + HL + 3H2 O → [Fe(L)Cl2 (H2 O)] · 2H2 O + HCl (3)
Fig. 3. Metal–ligand formation curves.
M.M. Omar, G.G. Mohamed / Spectrochimica Acta Part A 61 (2005) 929–936
S
– 9.83 (9.57) 12.53 (12.87) 10.22 (10.44) 10.75 (10.44) 14.62 (14.75) 11.56 (11.38) 17.89 (18.12) – – 12.85 (12.55) 11.63 (11.35) 7.00 (7.36) 11.59 (11.33) 11.15 (11.33) 7.05 (7.43) 11.52 (11.21) 10.48 (10.36) 8.32 (8.09) 5.00 (5.46) 16.03 (16.47) 14.52 (14.89) 9.93 (9.66) 15.10 (14.87) 14.62 (14.87) 9.43 (9.76) 14.95 (14.71) 13.20 (13.59) 10.87 (10.62) 9.68 (9.56)
N H
3.72 (3.50) 2.74 (2.84) 3.80 (3.45) 2.80 (2.48) 2.84 (3.48) 4.29 (4.18) 2.49 (2.45) 2.01 (2.27) 2.62 (2.28) 2.00 (1.54) Red (85) Brown (69) Brown (75) Brown (65) Brown (62) Brown (65) Red (71) Brown (70) Red (61) Red (59) HL (L = C13 H9 N3 OS) [Mn(L)2 ] (L = C26 H16 N6 O2 S2 ) [Fe(L)Cl2 (H2 O)]·2H2 O (L = C26 H12 Cl2 FeN3 O4 S) [Co(L)2 ] (L = C26 H16 CoN6 O2 S2 ) [Ni(L)2 ] (L = C26 H16 N6 NiO2 S2 ) [Cu(L)(AcO)]·3H2 O (L = C15 H17 CuN3 O3 S) [Zn(L)2 ] (L = C26 H16 N6 O2 S2 Zn) [Cd(L)2 ] (L = C26 H16 CdN6 O2 S2 ) [Pd(L)Cl] (L = C13 H8 ClN3 O4 PdS) [UO2 (L)(NO3 )] (L = C13 H8 N4 O4 SUO2 )
C
Found (%, calculated) mp (◦ C) Colour (% yield) Compound
Table 2 Analytical and physical data of HL and its metal complexes
A detailed interpretation of the IR spectra of HL and the effect of binding with Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), Cd(II), Pd(II) and UO2 (II) ions on the vibrational frequencies of the free HL ligand are discussed. The IR spectra of the free ligand and its metal chelates were carried out in the 4000–400 cm−1 range (Table 3). The IR spectrum of the ligand shows a broad band at 3500–3050 cm−1 , which can be attributed to the phenolic OH group. This band is still broad in all complexes, which renders it difficult to attribute to the involvement of phenolic OH group in coordination. The involvement of the deprotonated phenolic OH group in chelation is confirmed by the blue-shift of the (C O) stretching band, observed at 1215 cm−1 in the free ligand, to the extent of 5–16 cm−1 in the complexes [24]. Also the participation of the OH group is apparent from the disappearance (in UO2 (II) complex only) or shift in position of the ␦(OH) in-plane bending at 1409 cm−1 in the free ligand by 6–34 cm−1 in the remaining complexes [36]. However, the (N N) stretching band in the free ligand is observed at 1579 cm−1 [22,24]. This band is shifted to higher (6–23 cm−1 ) or lower (34–63 cm−1 ) frequency values upon complexation suggesting coordination via the azo group (M ← N) [24]. The IR spectrum of the ligand revealed a medium band at 1629 cm−1 due to (C N) of the N3 thiazole nitrogen. This band is shifted to higher (27 cm−1 ) or lower (2–35 cm−1 ) frequencies in the complexes indicating that it has been affected upon coordination to the metal ions [24]. The band at 744 cm−1 in the ligand is still in the same position in the complexes indicating the non-involvement of the thiazole S in coordination. In the far-IR spectra of all complexes, the non-ligand bands observed at 422–472 and 473–506 cm−1 region can be assigned to the (M N) stretching vibrations
60.92 (61.17) 55.20 (55.32) 35.60 (35.86) 55.60 (55.22) 55.30 (55.22) 41.61 (41.81) 54.75 (54.64) 50.95 (50.49) 39.83 (39.44) 26.33 (26.62)
3.4. IR spectra and mode of bonding
165 ± 2 >300 >300 >300 >300 >300 >300 >300 >300 >300
M
– 5.30 5.37 5.15 3.0 1.87 Diam. Diam. Diam. Diam.
µeff (B.M.)
The results of the elemental analyses of the metal chelates of TAN, which are recorded in Table 2, are in good agreement with those required by the proposed formulae. In most cases, 1:2 (M:L) solid complexes were isolated and found to have the general formula [M(L)2 ] where M = Mn(II), Co(II), Ni(II), Zn(II) and Cd(II). While the Cu(II), Pd(II) and UO2 (II) chelates are of the type 1:1 (M:L) and have the formula [M(L)X]·nH2 O, where X = AcO in case of Cu (n = 3), X = Cl in case of Pd(II) (n = 0) and X = NO3 in case of UO2 (II) (n = 0). In addition, Fe(III) forms a 1:1 (M:L) chelate of the formula [Fe(L)Cl2 (H2 O)]·2H2 O. The HL ligand offers several alternatives to coordination to metals. The results shown in Table 2 prove that HL is indeed monoanionic where the OH proton of 2-naphthol is lost when coordinated to the metal ions, and thus reducing the charge on the complex ion. The presence of Cl, NO3 or AcO ions in the inner sphere of the chelates neutralizes the charge on the complex. Thus, HL is expected to fill three coordination positions and contribute a charge of −1. The non-involvement of Cl, NO3 or AcO ions in the coordination to the metal ions is also confirmed by the molar conductance measurements.
– 14.20 10.25 8.75 11.98 9.32 12.85 13.65 15.63 10.93
Λm (−1 cm2 mol−1 )
932
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933
Table 3 IR spectra (4000–400 cm−1 ) of HL and its metal complexes Compound
(C N)
(N N)
(C O)
␦(OH)
(H2 O)
(M O)
(M N) (thiazole)
(M N) (azo)
HL [Mn(L)2 ] [Fe(L)Cl2 (H2 O)]·2H2 O [Co(L)2 ] [Ni(L)2 ] [Cu(L)(AcO)]·3H2 O [Zn(L)2 ] [Cd(L)2 ] [Pd(L)Cl] [UO2 (L)(NO3 )]
1629m 1617s 1656s 1627s 1610sh 1594m 1612m 1613m 1600br 1627s
1579s 1594s 1602w 1585m 1525sh 1525sh 1545m 1516m 1545w 1591s
1215m 1199sh 1202sh 1204sh 1200sh 1203sh 1202sh 1201sh 1202sh 1210m
1409m 1391s 1403sh 1391m 1396s 1375m 1376m 1370m 1368m Dis.
– – 816sh, 47sh – – – – – – –
– 508s 503s 504w 528s 508s 525s 572w 565w 509w
– 437s 473m 473m 473w 485w 500w 514s 506w 474s
– 473w 422w 424w 436s 441w 439m 437m 472w 450w
of the azo and N3 thiazole nitrogen, respectively [6]. Conclusive evidence regarding the bonding of oxygen to the metal ions is provided by the occurrence of bands at 503–565 cm−1 as the result of (M O) [6,24]. Therefore, the IR spectra indicate that HL behaves as monobasic acid and the coordination sites being Ar OH, N N and the N3 atom of the thiazole moiety. 3.5. Molar conductance data The solubility of the complexes in DMF permitted determination of the molar conductivity (Λm ) of 10−3 M solutions at 25 ◦ C and, by comparison, the electrolytic nature for each complex. The low values of the molar conductance data listed in Table 2 indicate that the complexes are nonelectrolytes. 3.6. Magnetic susceptibility and electronic spectra measurements The diffuse reflectance spectra of Mn(II), Fe(III), Co(II), Ni(II) and Cu(II) chelates show three bands at 210–215, 245–254 and 300–311 nm which are attributed to the → ∗ and n → ∗ transitions, respectively, within the HL ligand. The solid reflectance spectrum of Mn(II) complex show three bands at 16,528, 25,000 and 28,571 cm−1 (Table 4) assignable to 4 T1g → 6 A1g , 4 T2g (G) → 6 A1g and 4 T1g (D) → 6 A1g transitions [37], respectively, which lie in the same range as reported for octahedrally coordinated Mn(II) ion. The magnetic moment (5.30 B.M.) is an additional evidence for an octahedral structure [37]. From the diffuse reflectance spectra, it is observed that, the Fe(III) chelate exhibits a band at 22,222 cm−1 , which may be assigned to the 6 A1g → 5 T2g (G) transition in octahedral geometry of the complex [22]. The 6 A1g → 5 T1g transition appears to be split into two bands at 12,594 and 17,482 cm−1 . The observed magnetic moment of Fe(III) complex is 5.37 B.M. Thus, the complex formed has octahedral geometry around the Fe(III) ion [22]. The band observed at 26,954 cm−1 can be attributed to ligand-to-metal charge transfer band.
The diffuse reflectance spectra of Co(II) complex give three bands at 12,820, 15,552 and 17,574 cm−1 . The third region at 25,974 cm−1 refers to the charge transfer band (L → MCT). The bands observed are assigned to the transitions 4 T1g (F) → 4 T2g (F) (ν1 ), 4 T1g (F) → 4 A2g (F) (ν2 ) and 4 T (F) → 4 T (P) (ν ), respectively, suggesting that there is 1g 2g 3 an octahedral geometry around Co(II) ion [38,39]. The magnetic susceptibility measurement (5.15 B.M.) is an indicative of octahedral geometry [39]. The Ni(II) complex reported herein is high spin with a room temperature magnetic moment value of 3.0 B.M.; which are in the normal range observed for octahedral Ni(II) complexes [40,41]. The diffused reflectance spectrum of the Ni(II) complex displays three bands at 12,626 cm−1 (ν1 ), 15,625 cm−1 (ν2 ) and 21,691 cm−1 (ν3 ) assigned to 3 A2g → 3 T2g , 3 A2g → 3 T1g (F) and 3 A2g → 3 T (P) transitions, respectively. The band at 25,839 cm−1 1g refers to the charge transfer band (L → MCT). The electronic spectrum of Cu(II) complex shows bands around 15,873, 17,482 and 21,739 cm−1 . The position of the
Table 4 Electronic bands and d–d transitions of the metal complexes Compound
Band position (cm−1 )
Assignment
[Mn(L)2 ]
16,528 25,000 28,571
4T
1g (G)
4T
2g (G)
12,594 17,482 22,222 26,954
6A
6 A → 5 T (G) 1g 2g Charge transfer
12,820 15,552 17,574 25,974
→ 4 T2g (F) → 4 A2g (F) 4 T (F) → 4 T (F) 1g 2g Charge transfer
12,626 15,625 21,691 25,839
→ 3 T2g → 3 T1g (F) 3 A → 3 T (P) 2g 1g Charge transfer
15,873 17,482 21,739
2A
[Fe(L)Cl2 (H2 O)]·2H2 O
[Co(L)2 ]
[Ni(L)2 ]
[Cu(L)(AcO)]·3H2 O
→ 6 A1g → 6 A1g 4 T (D) → 6 A 2g 1g 1g
→ 5 T1g
4T
1g (F)
4T
1g (F)
3A
2g
3A
2g
2E
1g
g
→ 2 B1g
→ 2 B1g
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bands together with its magnetic moment value (1.87 B.M.) indicates that this complex has a square planar stereochemistry [26]. The Zn(II) and Cd(II) chelates are diamagnetic and in analogy with the previously described Zn(II) and Cd(II) complexes containing ONO donor azo ligands [26] and according to the empirical formulae of these complexes, we propose an octahedral geometry for the Zn(II) and Cd(II) chelates. The diamagnetic nature of the UO2 (II) and Pd(II) complexes indicates octahedral and square planar geometry, respectively, as expected for UO2 (II) and Pd(II) metal ions [37]. 3.7. Thermogravimetric (TG) analysis Thermal data of the complexes are given in Table 5. The correlations between the different decomposition steps of the complexes with the corresponding weight losses are discussed in terms of the proposed formulae of the complexes. The Mn(II) complex with the general formula [Mn(C26 H16 N6 O2 S2 )] is thermally decomposed in three successive decomposition steps. The first estimated mass loss of 20.03% within the temperature range 80–220 ◦ C may be attributed to the loss of C3 H2 N3 S molecule (calculated mass loss = 19.86%). The energy of activation of this step is 45.37 kJ mol−1 . The second and third steps occur within the temperature range 220–950 ◦ C with an estimated mass loss 64.84% (calculated mass loss = 64.54%) are reasonably accounted for the decomposition of the remaining ligand molecules leaving MnS residue, where the activation energies are 44.19 and 39.43 kJ mol−1 for the second and third steps, respectively. The total estimated mass loss 90.69% (total calculated mass loss = 91.32%). The Fe(III) complex, [Fe(C13 H8 N3 OS)Cl2 (H2 O)]·2H2 O, is thermally decomposed in four successive decomposition steps within the temperature range 40–1100 ◦ C. The first decomposition step of estimated mass loss 8.34% within the temperature range 40–140 ◦ C may be attributed to the liberation of two hydrated water molecules (calculated mass loss = 8.28%). The energy of activation was 35.45 kJ mol−1 . The three steps found within the temperature range 160–1100 ◦ C with an estimated mass loss 73.84% (calculated mass loss = 73.33%) which is responsibly accounted for the decomposition of one coordinated water, 2HCl and the ligand molecules with a final oxide residue of (1/2)Fe2 O3 . The activation energies are found to be 63.01, 205 and 370 kJmol−1 for the second, third and fourth steps, respectively. The Cu(II) complex with the general formula [Cu(C13 H8 N3 OS)(AcO)]·3H2 O was thermally decomposed in four successive decomposition steps. The first estimated mass loss of 11.95% within the temperature range 40–130 ◦ C may be attributed to the liberation of three molecules of hydrated water (calculated mass loss = 12.54%). The energy of activation of this step is 71.80 kJ mol−1 . The
remaining steps of decompositions occur within the temperature range 130–1100 ◦ C with an estimated mass loss of 70.20% (calculated mass loss = 68.99%) and activation energies of 175, 244 and 244 kJ mol−1 , which corresponds to the loss of acetate (as CH4 and CO2 gases), and ligand molecule leaving Cu2 O residue with a total estimated mass loss 82.15% (total calculated mass loss = 81.53%). The Zn(II) complex, with the general formula [Zn(C26 H16 N6 O2 S2 )], shows decomposition pattern of four stages. The first and second steps with estimated mass loss of 44.69%, found within the temperature range 40–450 ◦ C corresponding to loss of one ligand molecule; C13 H8 N3 OS (calculated mass loss = 44.48%). The activation energies of these two steps are 76.95 and 158 kJ mol−1 , respectively. The remaining decomposition steps with an estimated mass loss 42.14% which is due to loss of the second ligand molecule leaving ZnO residue occurring within the temperature range 500–1100 ◦ C (calculated mass loss = 41.68%). The sum of the activation energy for the last two decomposition steps are 600 kJ mol−1 . The total estimated mass loss is 86.83% (total calculated mass loss = 86.16%). The uranyl complex with the formula [UO2 (C13 H8 N3 OS) (NO3 )] was thermally decomposed in six successive decomposition steps. The first estimated mass loss of 10.80% (calculated mass loss = 10.58%) within the temperature range 30–100 ◦ C can be attributed to the liberation of nitrate group as NO2 and (1/2)O2 gases. The energy of activation of this step is 19.35 kJ mol−1 . The second to sixth steps occur within the temperature range 110–850 ◦ C with an estimated mass loss 43.18% (calculated mass loss = 43.34%) with an energy of activation sum to 306.61 kJ mol−1 , which is reasonably accounted for the loss of the ligand molecule leaving UO2 as residue with total estimated mass loss 53.98% (total calculated mass loss = 53.92%). 3.8. Kinetic analysis The first stage of dehydration of the complexes was studied in more detail. The kinetic parameters such as activation energy (E∗ ), enthalpy (H∗ ), entropy (S∗ ) and free energy change of decomposition (G∗ ) were evaluated graphically by employing the Coats–Redfern relation [42]: AR log{Wf /(Wf − W)} 2RT = log log 1 − T2 θE∗ E∗ −
E∗ 2.303RT
(4)
where Wf is the mass loss at the completion of the reaction, W is the mass loss up to the temperature T, R is the gas constant, E∗ is the activation energy in kJ mol−1 , θ is the heating rate and (1 − (2RT/E∗ )) ∼ = 1. A plot of the left-hand side of Eq. (4) against 1/T gives a slope from which E∗ was calculated and A (Arrhenius constant) was determined from the intercept. The entropy of activation (S∗ ), enthalpy of activation (H∗ ) and the free energy change of activation (G∗ ) were calculated
M.M. Omar, G.G. Mohamed / Spectrochimica Acta Part A 61 (2005) 929–936
935
Table 5 Thermoanalytical results (TG, DTG) of metal complexes Complex
TG range (◦ C)
DTGmax (◦ C)
Estimated (%, calculated) Mass loss
Assignment
Metallic residue
MnS
Total mass loss
[Mn(L)2 ]
80–220 220–520 520–950
192 300 650
20.03 (19.86) 18.88 (18.44) 45.96 (46.10)
84.88 (84.40)
Loss of C3 H2 N3 S molecule Loss of C8 H8 molecule Loss of C15 H6 N3 O2 molecule
[Fe(L)Cl2 (H2 O)]·2H2 O
40–140 160–110
75 208, 520, 780
8.34 (8.28) 73.84 (73.33)
82.18 (81.61)
Loss of 2H2 O (coord.) Loss of 2HCl and ligand molecules
(1/2)Fe2 O3
40–130 150–1100
65 198, 278, 750
11.95 (12.54) 70.20 (68.99)
82.15 (81.53)
Loss of 3H2 O (hyd.) Loss of CH4 , CO2 and ligand molecule
Cu2 O
40–450 600–1100
70, 275 670, 850
44.69 (44.48) 42.14 (41.68)
86.83 (86.16)
Loss of C13 H8 N3 OS molecule Loss of C13 H8 N3 S molecule
ZnO
30–100 110–850
46 132, 220, 347, 629, 784
10.80 (10.58) 43.18 (43.34)
53.98 (53.92)
Loss of NO2 and (1/2)O2 Loss of ligand molecule
UO2
[Cu(L)(AcO)]·3H2 O [Zn(L)2 ] [UO2 (L)(NO3 )]
using the following equations: Ah ∗ S = 2.303 log R kT
(5)
According to the kinetic data obtained from the DTG curves, all the complexes have a negative entropy, which indicates that the complexes are formed spontaneously.
H ∗ = E∗ − RT
(6)
3.9. Structural interpretation of solid complexes
G∗ = H ∗ − TS ∗
(7)
The design and synthesis of a new tridentate azo ligand derived from 2-naphthol and 2-aminothiazole for use in square planar or octahedral molecular templates have been successfully demonstrated. The synthesis of the ligand and its complexes proved to be as straightforward as expected, giving high yields of the free ligand and its complexes in simple, onepot reactions. As anticipated, the ligand coordinates equatorially to four- or six-coordinate transition metal ions to give
where k and h are the Boltzman and Plank constants, respectively. The calculated values of E∗ , A, S∗ , H∗ and G∗ for the decomposition steps are given in Table 6. The correlation coefficients of the Arrhenius plots of the thermal decomposition steps were found to lie in the range 0.92–0.97, showing a good fit with the linear function. Table 6 Thermodynamic data of the thermal decomposition of metal complexes Complex
Decomposition temperature (K)
E∗ (kJ mol−1 )
A (s−1 )
S∗ (kJ mol−1 )
H∗ (kJ mol−1 )
G∗ (kJ mol−1 )
80–220 220–520 520–950
45.37 44.19 39.43
2.59 × 105 5.45 × 106 6.11 × 1012
−138.0 −116.0 −184.0
43.78 41.69 38.89
70.20 76.50 76.50
[Fe(L)Cl2 (H2 O)]·2H2 O
40–140 160–250 500–600 750–1100
35.45 63.01 205.0 370.0
1.20 × 105 9.03 × 105 3.99 × 1012 3.88 × 1013
−136.0 −128.0 −83.0 −104.0
34.82 61.37 200.0 250.0
45.01 87.97 205.0 306.0
[Cu(L)(AcO)]·3H2 O
40–130 150–230 250–400 650–1100
71.80 175.0 244.0 244.0
3.21 × 1010 6.98 × 1018 6.89 × 1021 2.31 × 1011
−31.13 −119.0 −174.0 −35.30
71.23 174.0 242.0 238.0
73.28 150.0 194.0 264.0
[Zn(L)2 ]
40–100 200–450 600–700 720–1100
76.95 158.0 212.0 388.0
1.09 × 1011 4.07 × 1014 7.47 × 1010 5.75 × 1018
−21.59 −35.44 −43.49 −106.0
76.37 156.0 207.0 381.0
77.88 146.0 235.0 291.0
30–100 110–180 200–250 330–380 600–660 760–850
19.35 22.92 40.53 48.32 79.84 115.0
4.00 × 106 8.55 × 106 6.22 × 105 1.58 × 106 1.40 × 106 1.82 × 105
−103.0 −105.0 −131.0 −128.0 −134.0 −152.0
18.97 21.82 38.70 45.44 74.61 109.0
23.71 35.74 67.62 89.69 159.0 228.0
[Mn(L)2 ]
[UO2 (L)(NO3 )]
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M.M. Omar, G.G. Mohamed / Spectrochimica Acta Part A 61 (2005) 929–936
Fig. 4. Suggested structural formulae of metal complexes.
square planar (Cu(II) and Pd(II)), and octahedral (Mn(II), Fe(III), Co(II), Ni(II), Zn(II), Cd(II) and UO2 (II)) environments, around the metal ion anchor. The proposed general structures of the complexes are shown in Fig. 4. There are additional factors which might affect the structure of these chelates like the possibility of various isomeric structures, bulkiness of the ligand and intermolecular hydrogen bonds, however, but due to instrument limitation, we are not able to explore this further.
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