comparative study of the cationic surfactants and their influence on the alkaline hydrolysis of...

Comparative Study of theCationic Surfactants andTheir Influence on theAlkaline Hydrolysis ofAcetylsalicylic AcidBIRENDRA KUMAR,1 KALLOL K. GHOSH,1 P. RODRIQUEZ DAFONTE2

1School of Studies in Chemistry, Pt. Ravishankar Shukla University, Raipur 492 010, India2Departamento de Quimica Fisica, Universidad de Santiago de Compostela, 15782, Santiago de Compostela, Spain

Received 28 February 2010; revised 26 May 2010; accepted 6 June 2010

DOI 10.1002/kin.20515Published online 26 October 2010 in Wiley Online Library (wileyonlinelibrary.com).

ABSTRACT: A kinetic study of the reaction of acetylsalicylic acid (aspirin) with sodium hydroxidehas been studied in the presence of some conventional and novel cationic surfactants. Thepseudo-first-order rate constant increases with the surfactant concentration initially and thendecreases. In comparison to conventional cationic surfactants, i.e., cetyltrimethylammoniumbromide and cetylpyridinium bromide, novel alkyldiethylethanolammonium bromide (R = C16)surfactant accelerated the alkaline hydrolysis significantly. The pseudophase ion-exchangemodel has been applied to fit the experimental results. Activation parameters have also beenevaluated. C© 2010 Wiley Periodicals, Inc. Int J Chem Kinet 43: 1–8, 2011

INTRODUCTION

Acetylsalicylic acid (trade name aspirin) is now aproved effective drug that is used for minor pains,muscle ache, fever, arthritis, blood thinning, and aslong-term low doses to prevent heart attacks [1–7]. Alot of effort has been put into studying the effect ofself-organizing assemblies and its pharmaceutical for-mulation, solubilization, stability, etc. The hydrolysisof aspirin has been studied extensively to understand itsstability in biological systems for pharmaceutical for-mulation [8–14]. The unique role of surfactants in solu-

Correspondence to: Kallol K. Ghosh; e-mail: [email protected]© 2010 Wiley Periodicals, Inc.

bilizing the drugs depends on numerous factors such aschemical structure of the surfactant, chemical structureof the drug, temperature, pH, and ionic strength [15].Furthermore, surfactants have the capacity to modifyrates and equilibria of a chemical reaction and playan important role in the formulation of pharmaceuticalforms by reducing toxicity and increasing their sta-bility. Self-organizing assemblies such as micelle andmicroemulsions are highly efficient drug delivery ve-hicles for hydrophobic and partially hydrophilic drugs[16,17]. Determining a stable environment for drugs isof crucial importance for the design and process of ade-quately implanting pharmaceutical formulae in biolog-ical systems. Therefore, it becomes necessary to collectkinetic data on the influence of cationic surfactants onthe alkaline hydrolysis of aspirin. In comparison to

2 KUMAR, GHOSH, AND DAFONTE

+ H2O NaOH OH

O OH

Salicylic acid

O

O OH

O

Acetylsalicylic acid

O

HO

Acetic acid

+

Scheme 1

reactions in aqueous solutions, reaction kinetics inmicellar solutions may undergo alterations, whichmay be favorable or unfavorable in nature, dependingon micelle charge [18–20], micelle counterion [21],head group size [22,23], and the chain length of thehydrophobic tail of the surfactant [24]. Kinetically,micelles make a highly appealing reaction medium.Because of their structure, they can inhibit chemi-cal reactions by encapsulating a reactant within; con-versely, they can catalyze reactions by having the re-actants concentrated at the interface thus acting asmicroreactors. The alkaline hydrolysis of aspirin hasbeen widely studied in micellar media, whose ef-fects depend on the substrate and surfactant structure.The significant work of Broxton [25–27], Rodenas[28–32], Segovia [33], and Ferrit [34–37] and theirco-workers deserves special mention in this context.Broxton et al. [27] investigated hydrolysis of aspirinin the presence of cetyltrimethylammonium bromide(CTAB) at pH 6–8 and 7–9 and proposed a suitablemechanism. They also studied hydrolysis of aspirinderivatives in micelles. Similarly, Rodenas and Vera[30] developed a new model for the basic hydrolysisof acetylsalicylic acid in CTAB, cetyltrimethylammo-nium chloride (CTACl), and cetyltrimethylammoniumhydroxide (CTAOH), respectively. Recently, alkalinehydrolysis of two salicylic acid derived drugs, i.e., as-pirin and triflusal, has been studied spectrophotomet-rically in the presence of cationic, anionic, nonionic,

and zwitterionic surfactants by Ferrit and co-workers[34,36]. In this work, we have investigated micellar ef-fects on the alkaline hydrolysis of aspirin (Scheme 1)using alkyldiethylethanolammonium bromide (R =C12H25, C14H29, C16H33), CTAB, and cetylpyridiniumbromide (CPB) (Scheme 2). Alkanol amine surfactantscombine the characteristics of the amine and hydroxylgroups in terms of reactivity patterns. The presenceof both amine and hydroxyl groups in alkanolaminesmakes them very versatile for the present investigation.

EXPERIMENTAL

Materials

Aspirin USV limited product (Bangalore, India) wasused as supplied. Cetyl-diethylethanolammonium bro-mide (CDEABr), tetradecyldiethylethanolammoniumbromide (TDEABr), and dodecyldiethylethanolammo-nium bromide (DDEABr) were obtained from Prof.R. M. Palepu (retired), St. Francis Xavier Univer-sity, Antigonish, Canada. CTAB and CPB were pro-cured from Sigma/Aldrich, St. Louis, MO, USA.Sodium hydroxide (NaOH) was obtained from Merck(Mumbai, India) and acetonitrile from s.d. fine chemi-cals (Mumbai, India). All chemicals were used withoutfurther purification, and all the solutions were preparedin triple-distilled water.

R

N +

OH

Br −

C 2H5

C2H 5

R = C 16 H33 , C 14 H29 , C 12H 25

Cetylpyridinium bromide (CPB)

N +

C 16H33

Alkyldiethylethanolammonium bromide

Br −

N +

C 16H 33

Cetyltrimethylammonium bromide (CTAB)

Br −

Scheme 2

International Journal of Chemical Kinetics DOI 10.1002/kin

CATIONIC SURFACTANTS AND THEIR INFLUENCE ON HYDROLYSIS OF ACETYLSALICYLIC ACID 3

Figure 1 (a) Repeat scans every 0.2 min showing the increasing absorbance at 296.5 nm due to the hydrolysis of aspirin inthe absence of micelle. (b) Repeat scans every 0.1 min showing the increasing absorbance at 296.5 nm due to the hydrolysis ofaspirin in the presence micelle. [Color figure can be viewed in the online issue, which is available at wileyonlinelibrary.com.]

Kinetic Measurements

The kinetic studies were performed using a VarianCarry 50 UV–visible spectrophotometer equipped witha Peltier temperature control unit. The reaction wasfollowed by measuring the appearance of the salicy-late ion at the wavelength 296.5 nm at 28◦C. All re-actions were carried out under pseudo-first-order con-ditions in which the sodium hydroxide concentration(0.005 M) was always in larger excess over the sub-strate. The stock solution of aspirin (0.006 M) wasprepared in acetonitrile. The final concentration was[substrate] = 3 × 10−4 M, 5% (v/v) acetonitrile,and [NaOH] = 0.005 M in the reaction mixture.The pseudo-first-order rate constants (kobs) were deter-mined from the plots of log (A0 − At /A∞ − At ) ver-sus time, where A0, At , and A∞ being the absorbancereadings at zero, time, and infinite time, respectively.Figure 1a shows the UV absorption spectra in the ab-sence of micelle, and Fig. 1b shows the UV absorptionspectra in the presence of micelle [CDEABr]. The ab-sorption band centered at λ = 296.5 nm increases withtime.

RESULTS AND DISCUSSION

Effects of Sodium Hydroxide

The reaction was first order with respect to NaOH.Hence, in water the rate equation for the process iskobs = kW [OH−], where kW is the second-order rateconstant. The alkaline hydrolysis of aspirin was stud-ied at 0.001, 0.005, and 0.01 M NaOH concentration inthe absence and presence of CDEABr at 28◦C. The ob-served rate constant values are given in Table I. The cor-responding experimental data are displayed in Fig. 2.The pseudo-first-order rate constant increases linearlywith an increase in [NaOH]. The intercept at the ori-gin of the straight line indicates that the spontaneouswater-catalyzed reaction is insignificant in comparisonwith the OH− ion and surfactant.

Alkaline Hydrolysis in the Presence ofSurfactant

The effect of cationic micelles on the alkaline hy-drolysis of aspirin was studied at fixed [NaOH] and

Table I Observed First-Order Rate Constants for Hydrolysis of Acetylsalicylic Acid (Aspirin) in the Absence andPresence of Cationic Micelles with [NaOH]

Aqueous CDEABr

Substrate [NaOH] (M) kobs 103 (S−1) kw (M−1 S−1) kobs 103 (S−1) kM (M−1 S−1)

ASA 0.001 0.14 0.14 1.50 1.500.005 0.73 0.14 5.93 1.180.01 1.3 0.13 12.3 1.23

Experimental conditions: temperature = 28◦C, [substrate] = 3 × 10−4 M, and medium 5% (v/v) acetonitrile.



Figure 2 Variation of the pseudo-first-order rate constants (kobs) with NaOH in the absence and presence of CDEABr at 28◦C.

[surfactants] between 0.25 and 21 mM at 28◦C. Thevalues of pseudo-first-order rate constant (kobs) in alka-nol amine surfactants and conventional surfactants aregiven in Table II. The reaction rates initially increaseas the surfactant concentration is increased (Fig. 3).The effect is very significant in the CDEABr surfac-tant. The presence of catalyst is easily explained asan electrostatic effect; the cationic nature of the sur-factant favors the presence of OH− in the micellar

medium, accelerating the hydrolysis of the substrateassociated with micelles. The kobs values in the pres-ence of cationic micelles gradually accelerated at thelow surfactant concentration and at high surfactant con-centration show the inhibition effects for hydrolysis ofaspirin. The variation of kobs values of the reactionsdepends on the micellar structure, i.e., hydrophobictail length, counterion, and so on. The rate constantsin a variety of association colloids are slightly higher

Table II Observed First-Order-Rate Constants for Hydrolysis of Acetylsalicylic Acid (Aspirin) in the Presence ofCationic Micelles

103 kobs (s−1)Concentration ofSurfactants (mM) CDEABr TDEABr DDEABr CPB CTAB

0.00 0.73 0.73 0.73 0.73 0.730.25 1.40 0.85 – – –0.50 2.05 0.98 0.90 1.31 0.851.00 4.16 1.45 1.11 1.98 1.311.50 5.50 1.95 1.28 2.21 1.402.00 6.12 2.99 1.56 2.50 –2.60 8.72 3.64 1.85 2.80 1.504.00 11.6 7.20 2.50 – 1.985.00 12.5 8.56 3.20 2.50 1.458.00 15.9 7.65 2.40 1.70 –10.5 15.8 – 1.78 1.23 0.9812.5 14.7 5.30 1.41 – 0.7515.0 14.0 4.20 1.22 – –21.0 11.8 2.50 0.85 0.83 0.65

Experimental conditions: temperature = 28◦C, [substrate] = 3 × 10−4 M, medium 5% (v/v) acetonitrile, and [NaOH] = 0.005 M.



Figure 3 Observed first-order-rate constants for hydrolysis of acetylsalicylic acid (aspirin) in the presence of cationic micelles.[Color figure can be viewed in the online issue, which is available at wileyonlinelibrary.com.]

than in water and increase modestly with increasingmicellar media, involving variation in the surfactanttail groups and changes in the interfacial region. Anal-ysis of kinetic data indicates that the CDEABr showshigher reactivity. The increase in kobs values with in-creasing alkyl chain lengths (R = C16, C14, C12) ofthe surfactants, i.e., with increasing aggregation num-ber of micelle, is mainly due to the increase in theelectrical surface potential of the micelle and partiallydue to an increase in hydrophobicity of the palisadelayer of the micelle. In all cases, the [CDEABr] rateprofile is maximum as compared to [TDEABr] and[DDEABr] surfactant because [CDEABr] correspond-ing to complete solubilization of the substrate in themicellar pseudophase and alkanol group can exert aprotective effect on the positive charge, thus lower-ing the attractive electrostatic interactions between thequaternary ammonium center of the micelle and corre-sponding negative charge of the dissociated acid group.The effect of CTAB is not very significant. The reac-tivity order follows the trend CDEABr > TDEABr >

DDEABr > CPB > CTAB.

Application of the PseudophaseIon-Exchange Model

The experimental results can be explained by meansof the pseudophase model developed by Bunton et al.[38], which considers the total volume of micelles asa separate phase uniformly distributed in the aqueousphase, and that the reaction occurred independently inboth phases, with an equilibrium distribution of sub-strate between both phases according to Scheme 3, in

which S is the substrate, D is the monomeric surfactant,Dn is micellized surfactant, whose concentration isgiven by [Dn] = [D] – CMC (where CMC is the criticalmicelle concentration), n is the number of monomerswithin a micelle, the subscripts M and W denote themicellar and aqueous phase, respectively, and Ks isthe binding constant of the substrate to the micelle ex-pressed in terms of micellized surfactant, where k′

Wand k′

M are the pseudo-first-order rate constants for thereactions in aqueous and micellar pseudophases, givenby

Ks = [SM]

[SW][Dn](1)

k′W = kW[OH−

W] (2)

k′M = kM[OH−

M]

[Dn]= kM mOH (3)

For reactions where two ions compete for the chargedmicelle surface assuming that ions bind to micellesaccording to the exchange model developed forion-exchange resins, the Romsted pseudophase ion-exchange (PPIE) model [39] considered that the neu-tralized fraction of micellar head groups β can be taken

S W + Dn S M

Products

K s

k 'Mk 'W

Scheme 3



as constant when one of the ions in solution bindsvery strongly to a micelle. For the reactive ions (OH−)and the bromide ion as a micelle counterion, the ion-exchange equilibrium can be expressed by

OH−M + Br−W

kOHBr� OH−

W + Br−M (4)

with an equilibrium constant

KOHBr = [OH−

W][Br−M]

[OH−M][Br−W]

Setting

mOH = [OH−M]

[Dn]and mBr = [Br−M]

[Dn]

and β = mOH + mBr according to Scheme 3 andEqs. (1)–(4), the pseudo-first-order rate constant canbe easily derived as

kobs = kW[OH−]T + (kMKs − kW)mOH[Dn]

1 + Ks[Dn](5)

where

[OH−]T = [OH−W] + [OH−

M],

[Br−]T = [Br−W] + [Br−M]

All the concentrations used refer to the total volumeof the solutions. Calculation of mOH proceeds from theequation

mOH2 + mOH

{[OH−]T KOH

Br [Br−]T(K

OH

Br− 1

)[Dn]

− β

}

− β[OH−]T(KOH

Br − 1)[Dn]

= 0 (6)

According to the PPIE model, the OH− and inert coun-terion (Br−) of cationic micelle compete at the micellarsurface level, reaching an ion-exchange equilibrium.Equation (4) provides a representation to this process.The PPIE model explains the variation of the pseudo-first-order rate constant with different cationic mi-celles. When the surfactant is present in small amounts,the relative concentrations of substrate and ionic reac-tant in the micellar Stern layer increase rapidly withsurfactant concentration accelerating the reactions; thiseffect can account for the ascending branch of thecurve. The substrate is in the micellar pseudophase;an increase in the surfactant concentration increases

the unreactive counterions (Br−), which provokes adisplacement of the micellar-bound OH− ions in theproximity of the bound substrate. This effect can ac-count for the descending branch of the experimentalcurve at high surfactant concentrations.

When two ions compete for the micellar surface,the interaction is governed by an ion-exchange equi-librium considering the fraction of micellar surfaceneutralized, β, as a constant with a value of 0.8when the micellar counterion is the bromide ion.The CMC values were obtained from the literature:CDEABr (0.96 × 10−3 M), TDEABr (3.05 × 10−3

M), DDEABr (12.5 × 10−3 M), CPB (0.95 × 10−3 M),and CTAB (0.96 × 10−3 M) [40–43]. The slight ob-served catalysis is due to concentration of the substratewithin the micellar pseudophase for which the bindingconstant Ks is 68–1100 M−1, depending on the hydrox-ide concentration and surfactant head group. Values ofKOH

Br for different substrates lie within the range 2–40approximately. Our values are within this interval, andthis can be explained by the influence of the substrateon the equilibrium constant. The second-order rate con-stant (kM) was calculated with the help of computersimulation. It was necessary to calculate the second-order rate constant (kM) for the reactions occurring inthe micellar pseudophase because this provides a basicunderstanding and knowledge of the exact volume ofthe micellar pseudo-phase where the reaction proceeds.The observed variations in the kinetic parameters weredue to the interaction of hydrophobic nature of thesubstrates and ionic micelles. Both electrostatic andhydrophobic forces are important in the micelle solu-bilization of drugs of cationic micelles. These forceschange the surface polarity of micelles and thus affectthe ion equilibria and binding equilibria of substrates,which strongly affect the rate of reactions. Computersimulation of the variation of the observed rate constant(kobs) with the surfactant at fixed [NaOH], however,has shown that for the best fit, the pseudo-first-orderrate constant in the micellar rate constant and calcu-lated second-order rate constant in the micellar rateconstant (kM) is approximately 100 times smaller thanthat for the reaction in water. The binding constant(Ks) observed for CDEABr is 68 ± 14 M−1 and forCPB is 1100 ± 300 M−1, which depends on the mi-cellar head group size. It has been observed that theKs values are independent of the OH− concentration;these results are different from those obtained in thepresence of cationic micelles. The second-order rateconstant (kM) was calculated with the help of com-puter simulation, and it was observed that kM varied(0.011–0.43) M−1 s−1 with the varying surfactant headgroup size at fixed OH concentrations (Table III). Inthe case of aspirin, Ks values are higher for CPB than



Table III Values of Parameters Calculated for the PPIE Model for the Alkaline Hydrolysis of Acetylsalicylic Acid(Aspirin)

Surfactant kw (M−1 s−1) KOHBr kM (M−1 s−1) KS (M−1)

CDEABr 0.146 10 0.43 ± 0.06 68 ± 14CPB 0.146 4 (1.1 ± 0.1) × 10−2 1100 ± 300

CMC = 0.96 × 10−3 M (CDEABr) and 0.95 × 10−3 M (CPB) β = 0.80.Experimental conditions: tempearture = 28◦C, [substrate] = 3 × 10−4 M, medium 5% v/v) acetonitrile, and [NaOH] = 0.005 M.

CDEABr. The PPIE model has been applied to theexperimental results, and the best fitting results arepresented in Table III and a graphical representationis shown in Figs. 4 and 5. The substrate in the mi-celle would be constant, and surfactant chain lengthwas varied. The effect of this on the alkaline hydroly-sis of aspirin was studied for the best fit ion-exchangemodel with CDEABr and CPB, but other surfactants,TDEABr, DDEABr, and CTAB, are not fitted in thePPIE model. Thus a comparison of a micellar-bindingconstant of drug in the CPB and CDEABr micelles pro-vides the conclusive information on the reactivity pat-terns. The binding constant (Ks) values of aspirin havebeen compared in cationic micelles of the same chainlength, but different head group sizes, with the resultthat the difference between them is higher in cationicmicelle CPB > CDEABr, respectively. According tothese established relationships, it can be deduced thatthe hydrophobic nature of the substrate has a greaterinfluence on cationic micelle.

Figure 4 Influence of surfactant concentration upon thepseudo-first-order rate constant for alkaline hydrolysis ofaspirin in cationic micelles (CDEABr). The solid line is cal-culated by using the PPIE model.

Figure 5 Influence of surfactant concentration upon thepseudo-first-order rate constant for alkaline hydrolysis ofaspirin in cationic micelles (CPB). The solid line is calculatedby using the PPIE model.

To determine various activation parameters (Ea,�H ∗, �S∗, and �G∗), the pseudo-first-order rate con-stants for the alkaline hydrolysis of aspirin in the ab-sence and presence of CDEABr have been studied atdifferent temperatures (28 and 38◦C). The values aregiven in Table IV. This table shows that the observedactivation energy and enthalpy decrease in the micel-lar phase as compared to the aqueous medium. Thehigh negative value of �S∗ in micelle indicates thatthe reagent should be ordered to form the transitionstate.

Table IV Values of Activation Parameters

Ea (kJ �S∗ (J K−1 �H ∗ (kJ �G∗ (kJSystem mol−1) mol−1) mol−1) mol−1)

Aqueous 63.5 −95.7 61.0 90.2CDEABr 41.3 −152 38.7 85.2

CDEABr concentration = 1 × 10−3 M, [NaOH] = 0.005 M, andtemperature = 28 and 38◦C.



CONCLUSIONS

The kinetics of the alkaline hydrolysis of aspirin in thepresence of different cationic surfactants has been stud-ied. The reaction was studied in the presence of alkanolamine surfactants with different carbon chain lengths.The cationic surfactants catalyze the hydrolysis. Theresult fits the PPIE model.

The authors are grateful to Prof. Rama Pande, Head, Schoolof Studies in Chemistry, Pt. Ravishankar Shukla University,Raipur (C.G.), India, for providing laboratory facilities.

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