Synthesis , Electrochemical , Spectrophotometric and Potentiometric Studies of Two Azo-Compounds Derived from 4-Amino-2-Methylquinoline in Ethanolic-Aqueous Buffered Solutions

Dois compostos-azo, 2-metil-4-(5-amino-2-hidróxi-fenilazo)-quinolina e 2-metil-4-(5-amino2-hidróxi-nitrofenilazo)-quinolina, derivados da 4-amino-2-metilquinolina foram sintetizados. Suas estruturas químicas foram caracterizadas e confirmadas através de análise elementar, espectroscopia no infravermelho (IR), ressonância magnética nuclear (RMN) de H e espectrometria de massas (MS). O comportamento eletroquímico do composto de partida (4-amino-2-metilquinolina) e dos dois azo-derivativos sintetizados foi estudado com um eletrodo de mercúrio em solução tampão universal B-R em diferentes valores de pH (2-11,5) contendo etanol 40% (v/v) usando polarografia dc, voltametria cíclica e coulometria com potencial controlado. Os caminhos de reação dos compostos no eletrodo foram elucidados e são discutidos. As constantes de dissociação (pKa) dos compostos examinados, constantes de estabilidade e estequiometria dos complexos formados em soluções dos compostos com alguns íons de metais de transição (Co(II), Ni(II), Cu(II), La(III) and UO2 ) foram determinadas.


Introduction
Synthetic azo-dyes are among the most explored classes of organic compounds.][20][21][22][23][24][25] Although the electrochemical behavior of some Schiff base compounds derived from 4-amino-2-methyquinoline has been reported, 26 no studies concerning the electrochemical behavior of the latter one or its azo-derivatives are reported in the literature to date.
In the present work, two azo-derivatives from 4-amino-2-methylquinoline were synthesized and characterized.Their electrochemical behavior was investigated at mercury electrodes.Besides, the dissociation constants (pKa) of the investigated compounds, stability constant and stoichiometry of their metal complexes in solution with some transition metal ions were determined.
A set of the Britton-Robinson (B-R) universal buffer of pH values 2-11.5 was prepared in bidistilled water and used as a supporting electrolyte in the presence of 40% (v/v) ethanol.All chemicals used were of analytical grade (BDH or Merck) and were used without further purification.

Physical measurements
Elemental analysis of the synthesized azo-compounds (2 and 3) was carried out using Perkin Elmer 2400 elemental analyzer (Central Laboratory for Searching & Microanalysis, Tanta University, Egypt).Their melting points were measured using a Gallenhamp apparatus and were uncorrected.Infrared (IR) spectra of the solid compounds were recorded on a Jasco FT/IR-4100-A spectrophotometer within the range 4000-400 cm -1 as KBr discs (Microanalytical unit, Kafr El-Sheikh University, Egypt). 1 H nuclear magnetic resonance (NMR) spectra were recorded on a Varian EM 390-90 NMR spectrometer (Micro Analysis Center, Cairo University, Egypt) using DMSO-d 6 (Merck) as a solvent and tetramethylsilane (TMS) as an internal standard at room temperature.The 1 H NMR chemical shifts (d, ppm) are given relative to residual solvent peak.Mass spectral measurements of the synthesized compounds were made on a DI Analysis Shimadzu Qp-2010 Plus (Micro Analysis Center, Cairo University, Egypt).

Electrochemical measurements
A Sargent-Welch Polarograph model 4001 (Fisher, USA) was used in the dc-polarographic measurements.The electrochemical cell used was as described by Meites. 28haracteristics of the capillary used as dropping mercury electrode (DME) were m of 1.2 mg s -1 and t of 3.5 s in a solution of 0.1 mol L -1 KCl at open circuit and a mercury height of 60 cm.A saturated calomel electrode (SCE) was used as a reference electrode.
A computer-controlled Potentiostat/Galvanostat model 273A-PAR (Princeton Applied Research, Oak Ridge, TN, USA) and the electrode assembly model 303A-PAR with the software package 270/250-PAR were used in cyclic voltammetric measurements.A micro-electrochemical cell incorporated of a three-electrode configuration system comprising of a hanging mercury drop electrode (HMDE) as a working electrode (surface area of 0.026 cm 2 ), an Ag/AgCl/KCl s reference electrode and a platinum wire auxiliary electrode was used.
Mercury, gold, platinum and various carbonaceous materials are the most frequently used working electrodes in electrochemical studies.The use of mercury electrodes has been recently discouraged mainly due to environmental reasons (Hg is a toxic element).Mercury has a high hydrogen overvoltage that greatly extends the cathodic potential window up to about -2 V in aqueous solution.
Due to its wide cathodic potential window, mercury was frequently used as working electrodes (e.g., DME and HMDE) for electroreduction of substances of very negative reduction potential in aqueous electrolyte.Besides, it was successfully used as a working electrode for substance that strongly adsorbs onto the surface of solid electrodes.The mercury electrode is far superior to other solid electrodes in this regard since great care must be exercised with such solid electrodes to ensure that their surfaces are not changed by the electrochemical reactions.On the other side, mercury electrode (e.g., DME) also possesses a highly reproducible, readily renewable and smooth surface.
The appropriate concentration of each of the examined compounds in 10 mL B-R universal buffer in the presence of 40% (v/v) ethanol was introduced into the electrochemical cell, and then deoxygenated with pure nitrogen gas for about 10 min and for 30 s in each successive cycle, while a stream of nitrogen gas was kept over the solution during the measurements.
A potentiostat/galvanostat model 173-PAR incorporated with a digital coulometer model 179-PAR (Princeton Applied Research, Oak Ridge, TN, USA) was used for the controlled-potential coulometric measurements at a mercury pool cathode.A micro-coulometric cell incorporated with a Pt wire sealed through the cell bottom for contact with the mercury pool as a working electrode, a reference SCE and a platinum gauze as a counter electrode was used.
Controlled-potential electrolysis of solution (2.5 × 10 -4 mol L -1 ) of each of the investigated compounds was performed in the B-R universal buffer of various pH values containing 40% (v/v) ethanol to determine the total number of electrons consumed in the overall electrode reaction.The applied potential was adjusted to be around the half-wave (E 1/2 ) potential (± 0.1 V vs. SCE) or at the plateau of the limiting current of each of the recorded reduction waves.Prior to measurements, the electrolyzed solutions were deoxygenated by bubbling with pure nitrogen gas.During the measurements, a constant stream of nitrogen gas was passed over the surface of the electrolysis solution.The total charge Q (Coulombs) consumed during the complete electrolysis of the reactant was calculated by electronically integrating the current, after subtracting the background current.The number of electrons (n) transferred per electrolyzed molecule at various pH values was estimated using Faraday's equation: N = Q/nF where N is the number of moles of substance being electrolyzed and F the Faraday's constant (96,485 C mol -1 ).
A pH-meter model HI8014 (Hanna Instruments, Italy) accurate to ± 0.01 pH units was used for the pH measurements.A Mettler balance (Toledo-AB104, Greifensee, Switzerland) was used for weighing the solid materials.Deionized water was obtained from a Purite-Still Plus deionizer connected to an AquaMatic double-distillation water system (Hamilton Laboratory Glass Ltd., Kent, UK).

Spectrophotometric measurements
UV-Visible absorbance spectra of the examined compounds were recorded within the wavelength range 200-800 nm at room temperature using a Shimadzu UV-Vis spectrophotometer model 160A (Kyoto, Japan) with a quartz spectrometric cell (1 cm bath length).The absorbance spectra were scanned for the following mixtures: 5 mL of B-R universal buffer + 4.8 mL H 2 O + 0.2 mL EtOH as blank, while for the sample 5 mL of B-R universal buffer + 4.8 mL H 2 O + 0.2 mL of 2.5 × 10 -3 mol L -1 azo-compound dissolved in ethanol.

Potentiometric measurements
Potentiometric measurements were performed using a pH-meter model HI8014 (Hanna Instruments, Italy) accurate to ± 0.01 pH units.The electrode was standardized before and checked after each titration with standard buffer solution (Fisher, New Jersey, USA).A standard 0.02 mol L -1 NaOH aqueous-ethanolic solution (40% v/v ethanol) was added from a 5 mL total volume micro burette accurate to 0.01 mL and the contents of the titration vessel were stirred using a magnetic stirrer (Sargent-Welch, USA).All titration measurements were carried out at 298 K by circulating water from an Ultra-thermostat (JULABO F10, Seelbach, Germany) through the annular space of a double-walled Pyrex titration cell of 50 mL capacity.Each of the following mixtures was prepared and potentiometrically titrated against a standard 0.02 mol L -1 NaOH aqueous-ethanolic solution (40% v/v ethanol) at 298 K.The volume was made up to 50 mL using bidistilled water before the titration, keeping ethanol content at 40% (v/v) in all titrated solution mixtures: (i) 5 mL of 0.01 mol L -1 HCl + 5 mL of 1 mol L -1 KCl + 20 mL EtOH + 20 mL H 2 O; (ii) 5 mL of 0.01 mol L -1 HCl + 5 mL of 1 mol L -1 KCl + 5 mL of 5 × 10 -3 mol L -1 azo compound + 15 mL EtOH + 20 mL H 2 O; and (iii) 5 mL of 0.01 mol L -1 HCl + 5 mL of 1 mol L -1 KCl + 5 mL of 5 × 10 -3 mol L -1 azo-compound + 5 mL of 2 × 10 -3 mol L -1 metal ion (Co(II), Ni(II), Cu(II), La(III) or UO 2
A constructed QuickBasic language-PC program was used in computing the data resulted from potentiometric measurements to estimate the dissociation constants (pKa) of examined azo-derivatives and the stability constants and stoichiometry of their metal-complexes in solution.

Results and Discussion
Characterization of the synthesized azo-derivatives (2 and 3) As described in the Experimental section, the results provided by the different techniques are in good agreement with the chemical formulae proposed for the synthesized compounds 2 and 3. IR spectra of compounds 2 and 3 are shown in the Supplementary Information (SI) section (Figure S1).In this section, the mass spectra of compounds 2 and 3 (Figure S2) and their fragmentation patterns (Figures S3 and S4) are presented.The IR spectral bands of compounds 2 and 3 at 1650-1662 cm -1 were assigned for the N=N group.The stretching vibration bands, u(OH), of the hydroxyl group were found at 3327-3403 cm -1 .The bands at 2948-2975 cm -1 characterize the u(CH 3 ) stretching vibration.In compound 3, the band at 1390 cm -1 characterizes the nitro group (Figure S1 in the SI section).
The 1 H NMR spectrum of compound (2) exhibited singlet for 3 protons of (CH 3 ) group at d 2.41, singlet for (NH 2 ) at d 4.23, singlet for (OH) at d 5.11, singlet for pyridyl proton attached to C3 at d 7.14, multiplet for 3 aromatic protons at d 6.55-7.32 and multiplet for 4 protons of fused phenyl at d 7.42-8.12.
For compound (3), the 1 H NMR spectrum exhibited a singlet for 3 protons of (CH 3 ) group at d 2.52, singlet for pyridyl proton attached to C3 at d 6.34, multiplet for 3 aromatic protons of substituted phenyl at d 7.3-7.6 and multiplet for 4 protons of fused phenyl at d 7.38-7.96.
The mass spectral assignment of compound (2) showed a strong molecular ion peak [M + ] at m/z (278, 48%) (Figure S2 in the SI section).Decomposition of this ion may take place through two pathways.In the first path, the compound loses CH 3 radical forming ion at m/z (263, 14%), then it loses three H radicals, producing ion with a base peak at m/z (260, 100%).The second path, [M + ] lose C 6 H 6 ON radical, forming ion at m/z (170, 56%), this decomposition step was confirmed by the appearance of peak at m/z (108, 74%), characterized for C 6 H 6 ON radical.The fragment ion at m/z (170, 56%) may lose nitrogen gas followed by five H radicals forming ions at m/z (142, 34%) and (137, 1%), respectively (Figure S3 in the SI section).The mass spectral assignment of compound (3) showed an intense molecular ion peak [M] + at m/z (308, 100%) corresponding to the formula C 16 H 12 O 3 N 4 (Figure S2 in the SI section).The [M + ] may decompose through several fragmentation paths.In the first, the molecule loses C 10 H 8 N 3 moiety, forming ion with a peak at m/z 138, 37%.The decomposed ion (C 10 H 8 N 3 ) with peak at m/z (170, 43%) may lose CH 3 group and N 2 gas forming an ion at m/z 127, which may take hydrogen radical, forming ion with peak at m/z (128, 72%).In the second path, [M + ] lose nitro group forming ion with peak at m/z (262, 59%).The third path, [M + ] may lose CH 3 group and hydroxyl radical, forming ion at m/z 276, in which it takes three hydrogen radicals, forming ion with peak at m/z (279, 64%) (Figure S4 in the SI section).
For 2-methyl-4-(2-hydroxy-5-nitrophenylazo]quinoline (3), the voltammograms showed a main 8-electron irreversible cathodic step (E 1/2 ca.-0.34 to -0.97 V) in solutions of pH values lower than 10 (Figure 1, 1 st wave, curves g, h and i, and Figure 2, curve e), which splits into two steps in solutions of pH values higher than 10 (Figure 1, curve j ,and Figure 2, curve g).An additional irreversible cathodic step was also obtained at more negative potentials (E 1/2 ca.-1.4 V) in solutions of pH values 4-9 (Figure 1, curves h and i, and Figure 2, curve f), its limiting current gradually decreased upon the increase of pH of the medium until complete disappearance at pH values higher than 9 (Figure 1, curve j, and Figure 2, curve g).The half-wave potentials (E 1/2 ) or the peak potentials (E p ) shifted to more negative values on the increase of either the pH of the medium and the scan rate n, indicating the involvement of proton in the electrode processes 29 and the irreversible nature of the reduction processes, 29,30 respectively.
The reduction waves of the examined compounds at the DME were analyzed using the fundamental equation for the irreversible polarographic waves: 28 E d.e.= E 1/2 -(0.0591/αn a ) log (i/(i l -i)) (1) Plots of E d.e vs. log (i/(i li)) for the polarographic waves of investigated compounds at various pH values were straight lines with slope values S 1 (S 1 , mV = 59.1/αn a ) 28 reported in Table 1.Values of αn a (α is the symmetry transfer coefficient and n a is the number of electrons involved in the rate-determining step) at various pH values were determined from slope S 1 (Table 1).Also the E 1/2 vs. pH plots for the polarographic waves (of pH-dependent E 1/2 ) of the examined compounds were straight lines with slope values S 2 reported in Table 1.The number of protons (p) participating in the rate-determining step of the reduction process of the electro-active centers of the investigated compounds was determined by applying the relationship: 29,31 The data reported in Table 1 indicated that the ratedetermining step of the reduction processes involved the consumption of one proton (i.e., p = 1) over the entire pH range.The ratio (p/n a ) may have the value 1 (when value of p and n a are equal) or 0.5 (when value of n a is double that of p).The most probable values obtained for α-parameter (0.40-0.55) indicated that the ratio (p/n a ) equals 0.5, which suggested that the number of electrons n a involved in the rate-determining step of each of the electro-active centers should be double that of the involved protons (p), i.e., n a = 2 and p = 1.
On the other side, linear E p -pH plots of slope values (58-60 mV) at scan rate of 200 mV s -1 for the examined compounds were obtained.From slope values (13-15 mV) of the obtained linear plots of Ep versus ln v (slope, mV = 12.85/αn a ), 32 values of αn a (0.99-0.86) and α (0.43-0.49) were also estimated at na = 2, confirming again the irreversible nature32 of the electrode processes of the examined compounds.Moreover, values of αn a and α were also determined at various pH values using the equation: E p -E p/2 = 0.048/αn a V, 30 where (E p -E p/2 ) is the difference between cathodic peak potential E p and the half-peak E p/2 at half-height of peak current as a function  of scan rate v.The obtained αn a values were found to equal (0.96-1.06) and consequently values of the symmetry transfer coefficient (α) over the pH range were found to equal (0.48-0.53) at n a equals 2. This provided additional support of the irreversible nature 32 of the electrode processes of the examined compounds.

Controlled-potential electrolysis and TLC studies
As shown in Table 2, the total number of electrons (n) transferred per molecule in the reduction process at various pH values for 2.5 × 10 -4 mol L -1 solution of each of the examined compounds (1, 2 and 3) were determined by means of controlled-potential electrolysis as discussed in the Experimental section.
Thin layer chromatographic (TLC) experiments were carried out for monitoring the products of complete electrolysis of the solution of each of the investigated azo-compounds (2 and 3) using benzene/acetone (70:30 v/v) as eluent.This was performed by concentration of the completely electrolyzed solution of the each of examined azo-compounds, and then followed by extraction of the buffer ingredient with ether.The TLC-experiments showed two clear spots compared to a single spot for the starting compound (1), indicating that two products were formed due to the reduction cleavage of the -N=Ndouble bond.The two products were suggested to be the corresponding primary aromatic amines.The presence of the latter one was confirmed by carrying out diazo-coupling reaction 27,33 on the completely electrolyzed solutions of each of the investigated azo-compounds.Recovery of the characteristic colored azo-dye 27,33 in solution confirmed the presence of aromatic amines in the completely electrolyzed solution as a result of the reduction cleavage of -N=Ncenter in each of the examined compounds (2 and 3), which confirmed our suggested reduction mechanisms of the investigated compounds.
On the other hand, the UV-Vis absorbance spectra of solutions (1 × 10 -4 mol L -1 ) of the investigated azo-derivatives in ethanolic-aqueous B-R universal buffer (e.g., pH 7) before controlled potential electrolysis showed a characteristic band at l max 400 nm due to the n-p * transition of -N=Ngroup (e.g., Figure 3; spectrum (a) for compound 3).This band disappeared completely after controlled-potential electrolysis of the examined solutions (e.g., Figure 3; spectrum (b) for compound (3)).It is important to notice that the spectra of solutions of both derivatives ( 2) and ( 3) after controlled-potential electrolysis resemble that of the starting compound (1) (Figure 3; spectrum c).This behavior confirmed the cleavage of the -N=Nbond in both derivatives by controlled-potential electrolysis.

Electrode reaction pathways
Electrode reaction of compound (1)   4-Amino-2-methylquinoline (1) was reduced at the mercury electrode in buffered solutions of pH values 7-11 in a single 2-electron irreversible step which is assigned to the reduction-saturation of the -C=Ndouble bond of its quinoline moiety, 34 as illustrated in Scheme 2.

Electrode reaction of the azo-derivative (2)
The synthesized 2-methyl-4-(5-amino-2-hydroxyphenylazo)-quinoline (2) was reduced at the mercury electrode in buffered solution of pH 2-11.5 in two 2-electron steps (1 st and 2 nd steps), which were attributed to the reduction of the -N=Ndouble bond to the amine stage as illustrated in Scheme 3. The 3 rd reduction step that appeared  at more negative potential in solutions of pH values 4-9 is due to reduction of the -C=Ndouble bond of the quinoline moiety, 34 as illustrated in Scheme 2 of the electrode reaction of the starting compound (1).

Electrode reaction of azo-derivative (3)
T h e s y n t h e s i z e d 2 -m e t h y l -4 -( 2 -h y d r o x y -5-nitrophenylazo)-quinoline (3) was also reduced at the mercury electrode in buffered solution of pH values 2-10 in a single 8-electron step, which is attributed to reduction of either the -N=Ndouble bond to the amine stage and the -NO 2 group to the hydroxylamine stage, via the consumption of four electrons for each centre, 24,35 (Scheme 4, reaction 4).At pH values higher than 10, this reduction step splits into two steps.The 1 st and 2 nd steps were attributed to the reduction-cleavage of the -N=Ndouble bond to the amine stage and reduction of the NO 2 group to the hydroxylamine stage, 35 respectively, via the consumption of four electrons for each center (Scheme 4, reactions 5 and 6, respectively).The reduction step that appeared at more negative potential in solutions of pH values 4-9 is due to reduction of the -C=Ndouble bond of the quinoline moiety 34 as illustrated in Scheme 2 of the electrode reaction of the starting compound (1).

Spectrophotometric studies
UV-Vis absorption spectra of 5 × 10 -5 mol L -1 of each of compounds (2) and (3) were recorded in the B-R universal buffer of various pH values (2-12) within the wavelength range 200-800 nm.Compound 2 showed two absorption bands at l max 438 and 480 nm (n-p * ), with an isosbestic point at l max 462 nm (Figure 4).While compound 3 showed two absorption bands at l max 315 and 400 nm (n-p * ), with an isosbestic point at l max 345 nm (Figure 4).
The first band may be attributed to the non-ionized form, whereas the second one that develops as the pH increases is due to the ionized species of each of the examined Scheme 2. Electrode reaction pathway of the parent compound (1).Scheme 3. Electrode reaction pathway of the azo-derivative (2).pH curves at two different wavelengths were typical Z or S-shaped (Figure 5) which indicated the transformation of the molecule from one form into another one.
The observed high values of pKa of the synthesized azo-derivatives (2 and 3) may be attributed to the location of the ionizable group (-OH) in an ortho-position where it has lower ability to release H + ion. 39,40

Potentiometric studies
The average number of protons associated with the reagent molecule n A at different pH values 41 was calculated compounds as a result of proton dissociation of the hydroxyl group.The isosbestic point indicated the presence of an acid-base equilibrium between non-ionized and ionized species.The ionic form absorbs at longer wavelength indicating that the interamolecular charge transfer is easier in the ionic form than in the non-ionic one, confirming that the interamolecular charge transfer is influenced by the OH group.For 4-amino-2-methylquinoline (1), its absorbance spectra was also recorded under the same experimental conditions and was found to exhibit absorption band at l max 330 nm at pH values lower than 8, whereas, at higher pH values, this band was blue shifted to l max 300 nm.This behavior may be due to that protonation of species takes place on the nitrogen of quinoline moiety in solutions of pH lower than 8.However, with the increase of the pH values of the medium, deprotonation takes place.The changes in the absorption spectra on varying the pH of solutions were used for determination of the dissociation constant (pKa) of the examined compounds (1, 2 and 3).The absorbance-Table 3. Dissociation constants (pKa) of the compounds under investigation in ethanolic-aqueous B-R universal buffer using: the half-height (i), limitingabsorbance (ii), modified limiting-absorbance (iii), isosbestic point (iv) and modified isosbestic point (v) methods  from the titration curves of HCl solution in the absence and in presence of the examined azo-derivative (2 or 3) against 0.02 mol L -1 NaOH aqueous-ethanolic solution (40% v/v ethanol) at 298 K (Figure 6).Thus, the formation curves ( n A vs. pH) for the proton-ligand systems 42 were constructed and found to extend between 0 and 1.0 in the n A scale for both compounds (2 and 3).This means that each of the examined compounds 2 and 3 has one dissociable proton from the enolized hydrogen ion of the hydroxyl group (Table 4).This indicates that the acid-base equilibrium based on the molecular form HL is in complete accordance with the following equilibrium: (4)   As shown in Table 4, compound (2) has a lower acidic character (pK 1 H = 8.75) than compound (3) (pK 1 H = 7.5).This is quite reasonable because the presence of NH 2 group enhances the electron density by its high electrondonating mesomeric and inductive effects, respectively; thereby stronger O-H bond is formed.The presence of electron withdrawing group NO 2 leads to the opposite effect, 39,40 thus the pK 1 H values for the examined compounds follow   the sequence 2 > 3. The results obtained are of the same order when compared to those spectrophotometrically obtained (Table 3).The average number of reagent molecules attached per metal ion n, at different pH values, was calculated from the titration curves of HCl solutions in the absence and the presence of the examined azo-derivative (2 or 3) and the selected metal ions (Co(II), Ni(II), Cu(II), La(III) or UO 2 2+ ) against 0.02 mol L -1 NaOH aqueous-ethanolic solution (40% v/v ethanol) at 298 K (Figure 6).

Compound
The metal-titration curves (curves c-e) were well separated from that of free ligand (curve b), along the axis of added volume of NaOH solution, which is attributed to the release of H + ions as a result of the complexation process (Figure 6).The formation curves for the metal complexes of the two azo-derivative (2 and 3) with the selected metal ions (Co(II), Ni(II), Cu(II), La(III) or UO 2 2+ ) were obtained by plotting the average number of ligands attached per metal ion ( n) vs. the free ligand exponent (pL). 42y analysis of the formation curves, 43,44 the successive stability constant log K 1 and log K 2 of the studied metal complexes ML and ML 2 , respectively, were determined (Table 4) using a constructed QuickBasic language-PC program.As shown in Table 4, the differences observed in the standard deviations (SD) of the obtained data under the same experimental conditions were insignificant, confirming reproducibility and precision of the results.The data obtained can be pointed out based on the following: (i) no precipitate was observed in the titration vessel, indicating that the formation of metal hydroxide is excluded; (ii) the maximum n values for all investigated metal complexes were found to be ca.2, revealing that both ML and ML 2 types of complexes are formed in solution; (iii) for all complexes formed log K 1 were always found higher than those of log K 2 (Table 4), because the vacant sites of the metal ions more freely available for binding of a first ligand than for a second one; 40,45 (iv) the order of stability constants of the metal complexes of the azo-compounds (2 and 3) at 298 K were found as: UO 2 2+ > La(III) < Cu(II) > Ni(II) > Co(II).The sequence of stability of complexes of compounds ( 2) and (3) with Cu(II) > Ni(II) > Co(II) are in agreement with that reported by Irving and Williams. 46,47The UO 2 2+ has higher stability than those of the other metal complexes.This may be attributed to the bonded oxygen atoms which increase the electrostatic attraction between the metal ion and the coordinated ligands and overcome any steric hindrance offered by the oxygen of the oxygenated cations. 40

Table 1 .
Dc-polarographic data for the investigated compounds in ethanolic-aqueous B-R universal buffer solutions containing 40% (v/v) ethanol at 25 o C

Table 2 .
Results of controlled-potential coulometry measurements for the examined compounds