International Journal of Scientific & Engineering Research, Volume 4, Issue 12, December-2013 912

ISSN 2229-5518

Complexation of Some Transition Metal Ions with Benzooxazole Sulfamethazine: A Potentiometric and Thermodynamic Studies in Aqueous Solution

A. R. El-Shobaky, A.A. El-Bindary*

Department of Chemistry, Faculty of Science, University of Damietta, Damietta 34517, Egypt.

Abstract— The interaction of Mn2+, Co2+, Ni2+, Cu2+, La3+, Ce3+, UO2 2+ and Th4+ ions with benzooxazole sulfamethazine (BOSM) have been studied in aqueous solution. The proton-ligand dissociation constant of benzooxazole sulfamethazine and metal-ligand stability constants of its complexes with metal ions (Mn2+, Co2+, Ni2+, Cu2+, La3+, Ce3+, UO2 2+ and Th4+) have been determined potentiometrically in 0.1 mol dm-3 KCl and 40 % (by volume) DMF–water mixture and at (298, 308 and 318) K. The stability constants of the formed complexes increases in the order Mn2+, Co2+, Ni2+, Cu2+, La3+, Ce3+, UO2 2+ and Th4+. The effect of tem- perature was studied and the corresponding thermodynamic parameters (G, H and S) were derived and discussed. The dissociation process is non-spontaneous, endothermic and entropically unfavourable. The fotmation of the metal complexes has been found to be spontaneous, exothermic and entropically favourable.
Keywords —Benzooxazole sulfamethazine, Stability constants and thermodynamics.
*Corresponding Author: E-mail:; Tel.: 002 01114266996; Fax: 002 0572403868.
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Sulfonamides were the first effective chemotherapeutic agents to be employed systematically for the prevention

and cure of bacterial infections [1]. Sulfonamide and azosul- fonamide derivatives have been found to be biologically ver- satile anticancer, antimalarial and antitubercular drugs [2]. Their metal complexes are highly antibacterial and antifungal [3]. Although potentiometric studies of azo compounds have


Benzooxazole sulfamethazine [4-((benzo[d]oxazol-2- ylthio)diazenyl)-N-(4,6-dimethyl pyrimidin-2-yl)benzenesufon amide] (Fig. 1) was prepared as previously described, using standard procedures [7,11]. The purity was checked by ele- mental analysis, IR and 1H NMR spectra.
been studied extensively [4-6], little attention including azo compounds formed by interaction of benzooxazole and sul- fonamides as ligands has been reported [7]. In continuation of earlier work [8-10], we report here the dissociation constant of BOSM and the stability constants of its complexes with Mn2+,





(Fig. 1)



Co2+, Ni2+, Cu2+, La3+, Ce3+, UO2 2+ and Th4+ at different temper- atures. Furthermore, the corresponding thermodynamic func- tions are evaluated and discussed.
Stock solution of (0.001 mol dm-3) of BOSM was pre- pared by dissolving an accurately weighed amount of the sol- id in DMF (Analar). Metal ion solutions of about (0.0001 mol

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International Journal of Scientific & Engineering Research, Volume 4, Issue 12, December-2013 913

ISSN 2229-5518

dm-3) were prepared from Analar metal chlorides in bidistilled
water and standardized with EDTA [12]. Solutions of 0.001 mol dm-3 HCl and 1 mol dm-3 KCl were also prepared in bidistilled water. A carbonate-free sodium hydroxide solution in 40 % (by volume) DMF–water mixture was used as titrant and standardized against oxalic acid (Analar).
The apparatus, general conditions and methods of cal- culation were the same as in previous work [8-10]. The follow- ing mixtures (i) – (iii) were prepared and titrated potentiomet- rically at 298 K against standard 0.002 mol dm-3 NaOH in a 40
% (by volume) DMF–water mixture:
i) 5 cm3 0.001 mol dm-3 HCl + 5 cm3 1 mol dm-3 KCl + 20 cm3
ii) 5 cm3 0.001 mol dm-3 HCl + 5 cm3 1 mol dm-3 KCl + 15 cm3
The average number of the protons associated with the
ligand (BOSM) at different pH values,nA , was calculated from the titration curves of the acid in the absence and presence of OSM. Applying eq. 1:

(V1 – V2 ) (N° + E° )

nA = Y + (1)

(V° + V1 ) TC°L

where Y is the number of available protons in BOSM (Y=1) and V1 and V2 are the volumes of alkali required to reach the same pH on the titration curve of hydrochloric acid and rea- gent, respectively, V° is the initial volume (50 cm3) of the mix- ture, TC°L is the total concentration of the reagent, N° is the normality of sodium hydroxide solution and E° is the initial concentration of the free acid. Thus, the formation curves (nA vs. pH ) for the proton-ligand systems were constructed and
found to extend between 0 and 1 in thenA scale. This means


DMF + 5 cm3 0.00l mol dm-3 ligand.
iii) 5 cm3 0.001 mol dm-3 HCl + 5 cm3 l mol dm-3 KCl + 15 cm3
DMF + 5 cm3 0.001 mol dm-3 ligand + 5 cm3 0.0002 mol dm-3 metal chloride.
For each mixture, the volume was made up to 50 cm3
with bidistilled water before the titration. For each system three replicate measurements were carried out under nitrogen
that BOSM has one ionizable proton (the enolized hydrogen
ion of the sulphonamide group, pKH) [4]. Different computa- tional methods [15] were applied to evaluate the dissociation constant. Two replicate titrations were performed and the av- erage values obtained are listed in Table 1. The completely protonated form of BOSM has one dissociable proton, that dissociates in the measurable pH range.

Table 1. Table 2. Thermodynamic functions for the dissociation of

atmosphere. These titrations were repeated for temperatures
of 308 K and 318 K. The temperature was controlled to within
± 0.05 K by circulating thermostated water (Neslab 2 RTE 220) through the outer jacket of the vessel. The pH measurements were performed with a HANNA instruments model 211 pH- meter accurate to ± 0.01 units. The term pH is in this work de- fined as –log [H+]. The pH–meter readings in the non–aqueous medium were corrected [13]. The electrode system was cali- brated according to the method of Irving et al. [14].

3. Results and discussion

BOSM in 40 % (by volume) DMF–water mixture and 0.1 mol dm-3 KCl at different temperatures.

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The formation curves for the metal complexes were
obtained by plotting the average number of ligands attached per metal ion (n ) vs. the free ligand exponent (pL), according to Irving and Rossotti [16]. The average number of the reagent molecules attached per metal ion,n, and free ligand exponent, pL, can be calculated using eqs. 2 and 3:

(V3 – V2 ) (N° + E°)

n = (2)

(V° + V2 )nA TC°M



The following general remarks can be pointed out:
(i) The maximum value ofn was ~ 2 indicating the formation of 1 : 1 and 1 : 2 (metal : ligand) complexes only [19].
(ii) The metal ion solution used in the present study was very dilute (2 x 10-5 mol dm-3), hence there was no possibility of formation of polynuclear complexes [20,21].
(iii) The metal titration curves were displaced to the right- hand side of the ligand titration curves along the volume
axis, indicating proton release upon complex formation of the

n = J

H 1

anti log pH

pL = log10

Σn = 0 β n

TCLo - nTCMo

Vo + V3



metal ion with the ligand. The large decrease in pH for the
metal titration curves relative to ligand titration curves point to the formation of strong metal complexes [22,23].
where TC°M is the total concentration of the metal ion present
(iv) For the same ligand at constant temperature, the sta-


in the solution, βH n is the overall proton-reagent stability con-
bility of the chelates increases in the order Mn
, Co
, Ni ,
stant. V1 , V2 and V3 are the volumes of alkali required to reach
Cu2+, La

3+, Ce

3+, UO2 2+

and Th4+

2+ 2+ 2+

the same pH on the titration curves of hydrochloric acid, or- ganic ligand and complex, respectively. These curves were


The dissociation constant (pKH) for BOSM, as well as the stability constants of its complexes with Mn2+, Co2+, Ni2+,
analyzed and the successive metal-ligand stability constants
Cu2+, La3+, Ce3+, UOR2R
and Th4+
have been evaluated at 298 K,
were determined using different computional methods [17,18]. The values of the stability constants (log K1 and log K2 ) are given in Table 2.
Table 2. Stepwise stability constants for ML and ML2 com-
plexes of BOSM in 40 % (by volume) DMF–water mixtures and
0.1 mol dm-3 KCl at different temperatures.
308 K, and 318 K, and are given in Tables 1 and 3, respectively.


298 K

308 K

318 K


log K1

log K2

log K1

log K2

log K1

log K2











































UO 2 2+







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Table 3. Thermodynamic functions for ML and ML2 complexes of
BOSM in 40 % (by volume) DMF–water mixture and 0.1 mol dm-3

log K= (-H / 2.303 R)(1/T ) + ( S / 2.303 R) (5)

From the G and H values one can deduce the entropy

S using the well known relationships 4 and 6:

S = (H-G) / T (6)

All thermodynamic parameters of the dissociation process of BOSM are recorded in Table 1. From these results the fol- lowing conclusions can be made:
(i) The pKH values decrease with increasing temperature, i.e. the acidity of the ligand increases [10].
(ii) A positive value of H indicates that dissociation is ac-
companied by absorption of heat and the process is endo- thermic.
(iii) A positive value of G indicates that the dissociation process is
not spontaneous [26].
(iv) A negative value of S is obtained due to the increased
order as a result of the solvation process.
All the thermodynamic parameters of the stepwise sta- bility constants of complexes are recorded in Table 3. It is known that the divalent metal ions exist in solution as octahe- drally hydrated species [18] and the obtained values of H and S can then be considered as the sum of two contribu- tions: (a) release of H2 O molecules, and (b) metal-ligand bond formation. Examination of these values shows that:
(i) The stability constants (log K1 and log K2 ) for BOSM com- plexes decrease with increasing temperature [9].
(ii) The negative value of G for the complexation process
suggests the spontaneous nature of such processes.
(iii) The H values are negative, meaning that these processes are exothermic and favourable at lower temperature.
The enthalpy (H) for the dissociation and complexa-
tion process was calculated from the slope of the plot pKH or log K vs. 1/T using the graphical reperesentation of van't Hoff eqs. 4 and 5:

G = -2.303 RT log K = H – T S (4)

(iv) The S values for the complexes are positive, confirming that the complex formation is entropically favourable [8].


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