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Multi-wavelength spectrophotometric determination of acidity constants of somesalicylaldimine derivatives

Maryam Bordbar a,⁎, Ali Yeganeh Faal b, Mohammad Mahdi Ahari-Mostafavi a,Mehrnaz Gharagozlou c, Razieh Fazaeli d

a Department of Chemistry, University of Qom, Qom, Iranb Department of chemistry, Payame Noor University, Iranc Department of Nanomaterials and Nanotechnology, Institute for Color Science and Technology, Tehran, Irand Department of Chemistry, Shahreza Branch, Islamic Azad University, 86145-311, Iran

a b s t r a c ta r t i c l e i n f o

Article history:Received 1 June 2012Received in revised form 14 October 2012Accepted 31 October 2012Available online 21 November 2012

Keywords:Schiff baseSalicylaldimineAcidity constantSQUAD

The acidity constants of 9 synthesized derivatives of Schiff base in dimethylformamide/water and ethanol/water(25:75 v/v) at 25 °C and an ionic strength of 0.1 M have been determined spectrophotometrically. All of thespectrophotometric data as pure spectra and distribution diagrams calculated with the SQUAD and MCR-ALSas hard modeling and soft modeling methods, respectively. Also the influence of substituents in the molecularstructure on the ionization constants is discussed.

© 2012 Elsevier B.V. All rights reserved.

1. Introduction

Schiff bases are very important tools for the chemists as thesehave been extensively studied as they possess many interesting fea-tures, including photochromic and thermochromic properties [1],proton transfer tautomeric equilibria [2], biological and pharmacolog-ical activities [3–6], as well as suitability for analytical applications[7]. Among Schiff bases, salicylaldimine Schiff bases have been exten-sively studied as they possess many interesting features [8–10].

Therefore, their successful application requires a detailed study oftheir characteristics. Also the predication of acidity constants of or-ganic reagents is important for understanding and quantifying chem-ical phenomena such as reaction rates, biological activity, biologicaluptake, biological transport and environmental fate [11]. Protontransfer reactions in Schiff bases have been studied extensively bothexperimentally and theoretically in the last three decades [12–14].

There have been several methods for the determination of acidityconstants, including the use of potentiometric titration, spectropho-tometry, capillary electrophoresis, and so on. Spectroscopic methodsare in general highly sensitive and as such are suitable for studyingchemical equilibria in solution. When the components involved in

the chemical equilibrium have distinct spectral responses, theirconcentrations can be measured directly, and the determination ofthe equilibrium constant is simple [15]. In many cases, the spectral re-sponses of two and sometimes even more components overlap con-siderably and their analysis is no longer straightforward. Therefore,to overcome this problem we have to employ graphical and compu-tational methods. The most relevant reports are on SPECFIT [16],SQUAD [17] and HYPERQUAD [18].

In this study the analysis is readily performed with the computerprogram SQUAD as the hard modeling method. In comparison, multi-variate curve resolution-alternating least squares (MCR-ALS) are usedas soft-modeling methods. MCR-ALS is one of the powerful tools forobtaining information about how the concentration of the species in-volved in the reactions evolves [19,20].

The MCR-ALS method is based on factor analysis. Ideally, the num-ber of significant factors should be the same as the number of chem-ical species involved in the reaction. The physical changes associatedwith the unknown equilibriums also that are reflected in the numberof acid–base species, which can be detected from the spectroscopicmonitoring of a titration experiment, were determined. Among thecomputational and statistical methods used to solve mixture analysisproblems, the FA, principal component analysis (PCA) [21,22] andsingular value decomposition (SVD) [23] techniques play a key role,especially in the estimation of the number of species contributing sig-nificantly to the experimental data variance. In this work the effectsof different substitutes were studied on the dissociation constantsand pure spectrum of these salicylaldimine Schiff bases.

Journal of Molecular Liquids 178 (2013) 70–77

⁎ Corresponding author. Tel.: +98 251 2906448; fax: +98 251 2916449.E-mail address: m.bordbar@qom.ac.ir (M. Bordbar).

0167-7322/$ – see front matter © 2012 Elsevier B.V. All rights reserved.http://dx.doi.org/10.1016/j.molliq.2012.10.039

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2. Experimental

2.1. Apparatus and materials

Absorption spectra were obtained with a Cary-100 UV–Vis spectro-photometer by using 1 cm path length glass cells and themeasurementswere performed at 25±0.1 °C. pH measurements were made with aMetrohm 729 pH-meter by using a combined glass electrode. A stockof L1–L6 solutions and L7–L8 containing 10 mg mL−1 of salicylaldiminederivatives was prepared by dissolving 0.001 g of these reagents in10 mL ethanol and DMF respectively, and adjusting the volume to 10.0.

The protonation constants and spectral profiles of salicylaldiminederivatives were obtained in a DMF/water and ethanol/water mixture(25:75 v/v) by the mixing of stock solutions and water. The recordingpH values in the binary ethanol/water and DMF/water solvents werecorrected by using the following equation:

pH� ¼ pH Rð Þ−δpH

where pH* is the corrected pH, pH(R) is the pH-meter readingobtained in a the binary solvent, and δ is the correction term [24].

The L1–L9 tetradentate Schiff base ligands are N,N′-bis(salicylidene)-1,2-cyclohexanediamine (L1), N,N′-bis(4-bromosalicylidene)-1,2-cyclohexanediamine (L2), N,N′-bis(4-nitrosalicylidene)-1,2-cyclo-hexanediamine (L3), N,N′-bis(salicylidene)ethylenediamine (L4),N,N′-bis(4-bromosalicylidene)ethylenediamine (L5), N,N′-bis(4-nitro-salicylidene)ethylenediamine (L6), N,N′-bis(6-metoxysalicylidene)

diethylenetriamine (L7), N,N′-bis(4-bromosalicylidene)diethylene-triamine (L8), N,N′-bis(4-nitrosalicylidene) diethylenetriamine (L9)were synthesized in our laboratory as shown in Fig. 1.

DMF, HCl, KCl and NaOHwere of pro-analysis grade from E. Merck.These chemicals were used without further purification. All of the so-lutions were prepared fresh daily. Due to that the salicylaldimines areunstable in solution, and are involved in various equilibriums likeketo-enol or in ring-chain tautomeric interconversion [25,26], allsolutions were allowed to remain in thermostated sample compart-ments under stirring for a minimum of 45 min before the spectrawere collected. All calculations were performed in MATLAB 7.5(Math Works, Cochituate Place, MA).

2.2. Equilibrium measurements

The acidity constants were evaluated from the computer fitting ofthe absorbance–pHdata to the equations that resulted from substitutingthe pH andabsorbance values in themass balances [27,28]. The resultingequations for diprotic acids are given in Eq. (1):

A ¼ A0 þ A1 Hþh i=K2

� �þ A2 Hþh i2

=K1K2

� �� �= 1þ Hþh i

=K2 þ Hþh i2=K1K2:

� �ð1Þ

In these equations, A is the observed absorbance at each titra-tion point, A0, A1 and A2 are the absorbances of the basic form,monoprotonated form and diprotonated form, respectively and K1 andK2 are the first and the second acidity constant.

N

OH

N

HO

N

OH

N

HO

L1 : R : HL2 : R : BrL3 : R : NO2

L4 : R : HL5 : R : BrL6 : R : NO2

R RR R

N

OH

N

HOR

R

NH

L7 : R : o-OCH3L8 : R : p-BrL9 : R : p-NO2

Group Group GroupA B C

Fig 1. Schematic representation of the ligands and labels.

Fig. 2. Absorption spectra of L6 at different pH values: (1) 1, (2) 1.5, (3) 3, (4) 2.5, (5) 3, (6) 3.5, (7) 3.8, (8) 4, (9) 4.2, (10) 4.5, (11) 4.75, (12) 5, (13) 5.25, (14) 5.5, (15) 5.75, (16) 6,(17) 6.25, (18) 6.5, (19) 7, (20) 7.5, (21) 8, (22) 9, (23) 10, (24) 11, and (25) 12 in water/ethanol (75:25% v/v).

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The A0, A1, A2, K1and K2 values were calculated by computerfitting of the absorbance–pH data to either Eq. (1) by using SQUADand MCR-ALS as hard modeling and soft modeling methods, respec-tively. In order to prevent the protonation of amine groups (in L7–L9)the pH of the solutions was kept upper than that of the pH of theligands.

2.3. Synthesis of ligands

2.3.1. Synthesis of L1–L3A solution of 1,2-diaminocyclohexane (5 mmol) in 20 mL

dichloromethane was added to a 250-mL three-necked round bot-tomed flask. Salicylaldehyde derivatives (10 mmol) were dissolvedin 30 mL of dichloromethane and placed in the addition funnel. Thesolution of aldehyde was added to the stirred solution of diamineover 15 min. An exothermic reaction occurs; and the reaction mixturewas gently heated upon complete disappearance of the water–dichloromethane azeotrope. The resulting mixture then was allowedto cool slowly to room temperature and stirred for 15 min. During theconcentration and cooling period an orange yellow solid precipitated.The evaporation of the solvent afforded the crude Schiff base ligandsas orange viscous liquids which upon further drying afforded powdersin nearly 100% yield with a trace of excess salicylaldehyde derivates.The crude product was refluxedwith 20 mL of absolute ethanol, cooled,filtered and vacuum dried.

2.3.2. Synthesis of L4–L60.1 mol salicylaldehyde derivatives was dissolved in 45 mL ethanol

and heated to 40 °C, and then added dropwise to a solution of 0.05 molethylenediamine in ethanol under vigorous stirring. The mixture solu-tion turned light yellow, and for a while yellow sheet-like crystals pre-cipitated out. After 20 min, the mixture was cooled and the precipitatewas collected. The yellow solid was re-crystallized in ethanol and driedat room temperature under a vacuum.

2.3.3. Synthesis of L7–L9Salicylaldehyde derivatives (0.1 mol) were dissolved in 75 mL of

methanol, and to this was added a solution of diethylenetriamine(0.05 mol) in 25 mL of methanol. The reaction mixture thus obtainedwas refluxed on a water bath for 1 h. After reducing the volume of thesolvent to ca. 50 mL, the content was transferred into a beaker and theexcess solvent was evaporated under the current of air where a viscousyellow-red oil was obtained. This was further dried in a vacuum.

All of the known products were identical with authentic samplesby melting points, TLC and NMR determinations.

N,N′-bis(salicylidene)-1,2-cyclohexanediamine (L1)Mp 119–120 °C; 1H NMR (DCCl3) δ 1.48–1.99 (m, 8H); 3.33–3.36(m, 2H), 6.81–6.84 (m, 2H), 6.88–6.96 (m, 2H), 7.18(dd, J=1.2 Hz,7.6 Hz, 2H), 7.25–7.31 (m, 2H), 8.29 (s, 2H), 13.36 (s, 2H),

N,N′-bis(4-bromosalicylidene)-1,2-cyclohexanediamine (L2)Mp 189–191 °C; 1H NMR (DCCl3) δ 1.50–1.60 (m, 2H); 1.73–1.75(m, 2H); 1.91–1.97 (m, 4H); 3.31–3.36 (m, 2H), 6.81–6.87 (m, 2H),7.28–7.40 (m, 4H); 8.20(s, 2H), 13.26 (s, 2H)

N,N′-bis(4-nitrosalicylidene)-1,2-cyclohexanediamine (L3)Mp 217–219 °C; 1H NMR (DCCl3) δ 1.76–1.85 (m, 2H); 2.00–2.19(m, 2H); 2.22–2.25 (m, 4H); 3.42–3.54 (m, 2H), 6.96–7.00 (m, 2H),7.74–7.80 (m, 4H); 8.45(s, 2H), 13.45 (s, 2H)

N,N′-bis(salicylidene)ethylenediamine (L4)Mp 127–129 °C; 1H NMR (DCCl3) δ 3.96 (s, 4H); 6.87–6.98 (m, 4H);7.24–7.34 (m, 4H), 8.38 (s, 2H), 13.25 (s, 2H)

N,N′-bis(4-bromosalicylidene)ethylenediamine (L5)Mp 195–196 °C; 1H NMR (DCCl3) δ 3.96 (s, 4H); 6.86–6.89 (d, J=8.8 Hz, 2H); 7.29–7.41 (m, 4H), 8.31 (s, 2H), 13.18 (s, 2H)

N,N′-bis(4-nitrosalicylidene)ethylenediamine (L6)Mp 275–278 °C; 1H NMR (DCCl3) δ 4.03 (s, 4H); 6.77–6.79 (d, J=9.6 Hz, 2H); 8.10 (dd, J=8.8 Hz, 2.8 Hz, 2H), 8.45 (d, J=2.8 Hz,2H), 8.79 (s, 2H), 14.17 (s, 2H)

N,N′-bis(6-metoxysalicylidene)diethylenetriamine (L7)Mp 170–172 °C; 1H NMR (DCCl3) δ 2.01 (t, J=5.2,4 H); 3.64(t, J=5.2, 4H); 3.92 (s, 6H); 6.90–7.04 (m, 4H); 7.21 (t, J=3.6, 4H);8.64 (s, 2H), 13.22 (s, 2H)

N,N′-bis(4-bromosalicylidene)diethylenetriamine (L8)Mp 224–226 °C; 1H NMR (DCCl3) δ 2.70 (t, J=5.6, 4H); 3.80(t, J=5.6, 4H); 6.97–7.00 (d, J=8.8 Hz, 2H); 7.25–7.40 (m, 2H);7.40–7.48(m, 2H); 8.59 (s, 2H), 13.06 (s, 2H)

N,N′-bis(4-nitrosalicylidene)diethylenetriamine (L9)Mp 236–238 °C; 1H NMR (DCCl3) δ 2.89 (t, J=5.6, 4H); 3.64(t, J=5.6, 4H); 6.52–6.55 (d, J=9.6 Hz, 2H); 7.94–7.97 (m, 4H); 8.34(s, 2H), 13.65 (s, 2H)

3. Results

The absorption spectra that are obtained for the titration of a4.16×10−5, 5.39×10−5, 7.44×10−5, 7.45×10−5, 5.58×10−5,7.04×10−5, 4.996×10−5, 4.27×10−5 and 4.85×10−5 M of somesalicylaldimine derivative L1–L9 solutions respectively in water/ethanol and water/DMF mixture by a standard solution of 3 M NaOHto adjust the pH values at 200–600 nmwere recorded. In order to pre-vent the formation of cationic species of the Schiff base (i.e., H3L+ andH4L2+) the pH of the solutions was kept upper than that of the pKa1

value. The 3D absorbance–response surfaces representing the mea-sured multiwavelength absorption spectra of the Schiff base ligands(L6), on the dependence of pH at 25 °C are plotted in Fig. 2 and

N N

OH

OH

N N

OO

HH

Fig. 3. Tautomerism between enol-imine and keto-amine forms.

0

0.2

0.4

0.6

0.8

1

1.2

1.4

1.6

1.8

200 250 300 350 400 450

Wavelength

Abs

orba

nce

Fig. 4. Spectra of L1 in ethanol/water mixture (25:75 v/v) during time.

72 M. Bordbar et al. / Journal of Molecular Liquids 178 (2013) 70–77

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represent the input data of the regression program SQUAD. A compari-son of both the SQUAD and MCR-ALS program treatments, along withthe proposed strategy for efficient experimentation in deprotonationconstant determination, is presented. Even though the actual SQUADversion used has a limited dimension and input can contain 100 spectraonly, an efficient spectra sample 100×55 (ns×nw) was used (Fig. 2a)for regression analysis. As the changes in the spectra are small withindeprotonation in the range of 240–300 nm, however, both of the vari-ous deprotonated species L and LH exhibit partly similar absorptionbands. When such small changes in the absorbance spectra are avail-able, a very precise measurement of the absorbance is then necessaryfor the reliable detection of the deprotonation equilibrium studied.

One of the most important behaviors of the salicylaldimine Schiffbase in solution and in solid state is tautomerism. Tautomerism insalicylaldimine Schiff bases was investigated by using spectroscopyand X-ray crystallography techniques [29,30]. These Schiff baseswith the OH group in ortho position to the imino group are of interestmainly due to the existence of either OH....N or O....HN type of hydro-gen bond and tautomerism between enolimine and keto-amine forms[25], Fig. 3.

As it was shown in Fig. 4 because of the tautomerism the UV–visspectra of these compounds have changed during time. For the con-trol of this instability, all solutions were allowed to remain in roomtemperature under vigorus stirring for a minimum of 45 min before

the spectra were determined, because after 45 min the spectra ofthese Schiff base didn't change.

After collecting the absorption spectra of the synthesized Schiffbase in the solvent mixture at different pH values and performingPCA, the PCA results showed that there are three components inthe whole pH range. Then the matrix of the absorption data wasprocessed by the SQUAD program to obtain the acidity constantsand spectral profiles of the components which participate in theabsorbance data matrix. Distribution diagrams of the different speciesof salicylaldimine derivatives (L1–L9) are shown in Fig. 5. The acidityconstants of all the salicylaldimine derivatives studied were eval-uated by computer fitting of the corresponding absorbance–pH datato the appropriate equations (i.e., Eq. (1) for diprotic acids). Alsothe pure spectra of the salicylaldimine species in water/ethanol andwater/DMF were calculated by the program fitting of SQUAD andthat the corresponding results for L1–L9 are shown in Fig. 6.

In this work about MCR-ALS firstly the number of acid–base speciesdetected from the singular value decomposition (SVD) as a spectro-scopic monitoring method [31] and the pure spectra as the initial esti-mated data were predicted by the evolving factor analysis [32]. Alsothe multivariate curve resolution (MCR-ALS) has been implementedin the MATLAB and it is available in the internet [33].

Also the estimated acidic constant values for L1–L9 by the twomethods SQUAD and MCR-ALS are shown in Table 1.

Fig 5. Distribution diagrams of different species of salicylaldimine derivatives L1–L9: H2L (solid line), HL (dash line) and L (solid–dash line) in water/DMF and water/ethanol.

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4. Discussion

In this study, the structures of the newly synthesized N,N′-bis(salicylidene)-1,2-cyclohexanediamine (L1), N,N′-bis(4-bromo-salicylidene)-1,2-cyclohexanediamine (L2), N,N′-bis(4-nitro-salicylidene)-1,2-cyclohexanediamine (L3), N,N′-bis(salicylidene)ethylenediamine (L4), N,N′-bis(4-bromosalicylidene)ethylenediamine

(L5), N,N′-bis(4-nitrosalicylidene)ethylenediamine (L6), N,N′-bis(6-metoxysalicylidene)diethylenetriamine (L7), N,N′-bis(4-bromo-salicylidene)diethylenetriamine (L8), N,N′-bis(4-nitrosalicylidene)diethylenetriamine (L9) were identified by using 1H NMR, meltingpoint and UV spectral data, and these obtained spectral valueswere seen to be compatible with literature reports [34–41].

4.1. 1H NMR spectra

In the 1H NMR spectra of the ligands, the phenolic proton wasseen at 13–14.5 ppm. These high frequency orthophenolic hydrogensin all the Schiff bases are attributed to the presence of intramolecularhydrogen bonding and electron-withdrawing effect of the substitu-ents on the aromatic phenolic ring. The azomethine protons appearat 8–9 ppm. The phenol ring proton signals resolve in the range of6–8 ppm and the aliphatic proton in L1–L9 due to different chemicalenvironments was seen at 1–4 ppm. A singlet was seen at 3.92 ppmthat was assigned to L7 that has the OMe group.

4.2. Electronic spectra

With the aim to obtain information about the type of the electronictransition and in order to investigate the different effects, a series of

Fig 6. Pure spectra of L1–L9 species: H2L (solid line), HL (dash line) and L (solid–dash line) in water/DMF and water/ethanol.

Table 1The estimated acidic constant values for Schiff Base derivatives by two methods.

No. SQUAD MCR-ALS

pKa1 pKa2 pKa1 pKa2

L1 6.92 7.45 6.88 7.44L2 6.66 9.3 6.25 9.28L3 5.04 6.39 5.1 6.45L4 6.75 8.28 6.74 8.35L5 6.39 9.03 6.42 9.00L6 5.24 10.31 5.24 10.38L7 8.19 10.54 8.11 10.59L8 7.37 8.89 7.31 8.91L9 5.36 9.69 5.44 9.73

74 M. Bordbar et al. / Journal of Molecular Liquids 178 (2013) 70–77

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the Schiff bases with several bridge lengths were prepared and theelectronic spectra of the ligands were measured in ethanol/water andDMF/water mixture solvents (Fig. 7). Table 2 presents the color, meltingpoint and absorption spectra characterization of the L1–L6 and L7–L9 li-gands in ethanol/water andDMF/watermixture (25:75 v/v) respectively.

Generally, the recorded spectra of the ligands displayed threemain absorption bands (Fig. 7). The first band due to the n→л* tran-sition of the C_N appear between 300 and 450, 280 and 450 and 320and 450 in groups A, B and C respectively and in the presence of theelectron-withdrawing substituent an apparent blue shift is observed.Also the bands between 235and 280 nm are assigned as a π→π*transition involving C_N. The absorption band between 200 and235 nm is most probably related to the π–π* transition of the pheno-lic chromophore [37].

4.3. Estimation of acidity constants

The acidity constant of the various Schiff bases were calculated byusing the SQUAD program [17], designed to calculate the best valuefor the acidity constants of the proposed equation model (Eq. (2))by employing a non-linear least-squares approach. For the investiga-tion of validated results we performed this protocol by MCR-ALS asthe soft modeling method and that the comparisons of the resultsare shown in Table 1.

H2L þH2O↔HL−

HL− þH2O↔L2−:

ð2Þ

Generally, for all compounds studied by increasing the pH of theme-dium within pHs 2.0–12.0, the absorbance band at the 350–450 rangeband increases until a more or less constant value is reached (Fig. 8).This behavior can be ascribed to the expected easier electron transfer

(n–π*) on the increase in the pH of the medium as explained above.The recorded spectra of all compounds exhibit a clear isosbestic pointwithin the high pH range revealing the existence of equilibriumbetweenthe neutral and ionic forms of these compounds in such media (Fig. 8).

Fig. 9 shows the plot of the absorbance–pH data for the L1–L9 sys-tem, and the trend of increasing and decreasing absorbance values atthe selected wavelengths is easily observed. Each figure comprising aclear inflection, indicates typical dissociation processes. In all com-pounds the absorbance–pH curve at the selected wavelength displayedone or two inflections.

One of the outputs of the employedmultivariate data analysismethodis the concentration profiles of the acid–base forms of the studied weakacids. These are shown in Fig. 5. Another output of the employed multi-variate data analysis method is the pure spectra of the different acid–base species, which are given in Fig. 6. One can see that for the acidicformof themost Schiff bases no clear peakmaxima are observedwhereasthe basic forms represent a clear peakmaximum at about 380 nmwhichis dependent on the substitution pattern on the Schiff bases.

In general, the factors that influence the acidity of the compoundsare the inductive effects, steric effects, solvent effects, and hydrogenbonding and resonance effects. In this work, we observed that withregard to the acidity of the Schiff bases in water/ethanol and water/DMF, the inductive, steric and hydrogen bonding effects are strong.These are discussed below in detail.

4.3.1. The electronic effect of para substituted Schiff base ligandsConsidering the structures of the Schiff bases L1–L9 we can de-

duce that they have two deprotonation sites (two OH groups of thephenol rings).

The acidity of the Schiff bases is influenced by the substituents onthe aromatic phenolic ring. This order of increasing acidity seems log-ical when one thinks about the inductive electron-donating effect of

0

0.5

1

1.5

2

2.5

3L1

L2

L3

0

0.5

1

1.5

2

2.5L4

L5

L6

0

0.5

1

1.5

2

2.5

200 250 300 350 400 450 500 200 250 300 350 400 450 500 200 250 300 350 400 450 500

L7

L8

L9

Abs

orbn

ace

Wave length

Fig. 7. Electronic spectra of ligands were measured in ethanol/water and DMF/water mixture.

Table 2UV–vis spectral data of ligands.

Ligand Formula MW colour m. p.(°C)

π–π*(benzenring)

π–π*C_N

n–π*

λ (nm)

L1 C20H22N2O2 322.4 Fulvous 119–120 210 256 323L2 C20H20Br2N2O2 480.19 Yellow 189–191 223 253 340L3 C20H20N4O6 412.4 Orange 217–219 224 247 357L4 C16H16N2O2 268.31 Yellow 127–129 213 256 324L5 C16H14Br2N2O2 426.1 Yellow 195–196 222 252 341L6 C16H14N4O6 358.3 Yellow 275–278 – 250 309L7 C20H25N3O4 371.42 Dark

Yellow170–172 204 263 352

L8 C18H19Br2N3O2 469.17 Orange 224–226 – 253 340L9 C18H19N5O6 401.36 Dark

Yellow236–238 – 252 307

0

0.5

1

1.5

2

2.5

250 300 350 400 450

Wavelength

Abs

orba

nce

Fig. 8. Absorption spectra of the L9 (4.99×10−5 M) in different pH-values (4–12).

75M. Bordbar et al. / Journal of Molecular Liquids 178 (2013) 70–77

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the OMe group and the strong electron-withdrawing effect of the NO2

and Br groups in the para position of the phenol ring.About pKa1 when we arrange the studied Schiff bases and their re-

duced derivatives in an increasing acidity order for the deprotonationprocess we get the following sequence:

NO2 > Br > H > OMe:

About the second deprotonation (pKa2), according to the resultsfounded, hydrogen bonding seems to be more effective than the otherfactors and the formation of hydrogen bonding is dependent to theeffects of different substitutes and geometry of molecules, directly.The results showing the presence of the strong electron-withdrawinggroups in the para position of the phenol ring lead to an increase inthe hydrogen bonding effect. In detail about group A (L4–L6), thelower acidities can be attributed to the strong hydrogen bondingdue to the presence of the nitro group in para position of the phenolring. It is observed that the decrease in the electronegativity of L4–L6molecules gives rise to the increase of acidity (pKa2) as follows(Scheme 1):

H > Br > NO2:

4.3.2. Effect of the diamine bridge lengthAlso group C behaves in a similar manner to group B but because

of the distance between the two hydroxyl groups, the intramolecularhydrogen bonding ability of the molecules L7 and L8 is very weak thanthat of L9, while in group A it seems that because the rigidity of mol-ecules is not promising for hydrogen bonding to occur, the pKa1 andpKa2 values are closer to each other. The electronic absorption spectraand obtained acidity constants for the deprotonation process aredepicted in Table 1.

5. Conclusion

In the present study, the acid dissociation constants of a series ofsubstitutive derivatives of the salicylaldimine derivatives have beencarried out by using ultraviolet visible absorption spectroscopy in allpH regions. This study showed that the proposed program (SQUAD)is useful in obtaining acidic constants and concentration and spectralprofiles of the components which exist in a suitable pH range for adiprotic Schiff base and the results are in good agreement with the re-sults obtained by the MCR-ALS method. So we proposed this methodto determine the equilibrium characteristics of the diprotic acids in ashort time with good accuracy.

Fig. 9. Change of absorbance vs pH at λmax for L1–L9 (L1:255 nm; L2,L3,L5,L6,L7,L9:390 nm; L4:378 nm and L8:395 nm).

76 M. Bordbar et al. / Journal of Molecular Liquids 178 (2013) 70–77

Author's personal copy

Also the structural effects of these compounds as substituent ef-fect and bridge length between the two hydroxyl groups on acidityof the model were investigated.

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N N

OH

OH

N N

O-OH

H

N N

-OO-

NH HN

OO

N HN

OOH

H

N N

-OOH

H

pKa1pKa2

R R R R

R

RR

R R

R

RR

Scheme 1. Possible deprotonation patterns for investigated Schiff bases.

77M. Bordbar et al. / Journal of Molecular Liquids 178 (2013) 70–77