Oxidative Cleavage of β-Lactam Ring of Cephalosporins with Chloramine-T in Alkaline Medium: A Kinetic, Mechanistic, and Reactivity Study

Cephalosporins are β-lactam antibiotics, and the important drugs of this group are cephalexin, cefadroxil and cephradine. In the present research, the kinetics and mechanism of oxidation of cephalexin (CEX), cefadroxil (CFL), and cephradine (CPD) with chloramine-T (CAT) in alkaline medium were investigated at 301 K. All the three oxidation reactions follow identical kinetics with a first-order dependence each on [CAT]o and [substrate]o. The reaction is catalyzed by hydroxide ions, and the order is found to be fractional. The dielectric effect is negative. Proton inventory studies in H 2 O-D 2 O mixtures with CEX as a probe have been made. Activation parameters and reaction constants have been evaluated. Oxidation products were identified bymass spectral analysis. An isokinetic relation was observed with β= 378K, indicating that enthalpy factors control the rate.The rate increases in the following order: CPD > CFL > CEX. The proposed mechanism and the derived rate law are consistent with the observed kinetics.


Introduction
Cephalosporins are an important and large class of bactericidal antimicrobial agents [1]. These are semisynthetic antibiotics that contain a -lactam ring which is fused to a dihydrothiazine moiety. Cephalexin These are widely used antibiotics which differ from each other with diverse group of substituents, namely, phenyl, hydroxyphenyl, and cyclohexadienyl, at the 7th position of the cephem ring. These drugs are widely used to treat respiratory and urinary tract infections, bronchitis, pneumonia, prostatitis, and soft tissues infections that are often caused by sensitive bacteria [2,3]. From the literature survey, it is evident that a lot of attention has been paid on the analytical methods for the determination of these drugs with various reagents [4,5]. But no information is available on the oxidation of these drugs with any reagent from the kinetic and mechanistic aspects. It is also noted that despite the importance of these drugs, relatively little is known about their mode of action at the molecular level. Therefore, the mechanism and rate law of these drugs are obscure. The oxidation-kinetic studies of these drugs provide much information about their mechanistic chemistry.
Sodium N-chloro-p-toluenesulfonamide (p-CH 3 C 6 H 4 SO 2 NClNa⋅3H 2 O), commonly known as chloramine-T (CAT) is of great interest due to its diverse behaviour. The versatile nature of CAT is due to its ability to act as sources of halonium cations, hypohalite species, and nitrogen anions which act both as bases and nucleophiles [6]. It is a potent oxidizing and chlorinating agent in both acidic and alkaline media, with a two electron change per mole, giving p-toluenesulfonamide (PTS) and NaCl. The oxidation potential of PTS/sulfonamide system is pH dependent [7], and it decreases with the increase in pH of the medium (1.139, 0.778, and 0.614 V at pH 0.65, 7.0, and 9.7, resp.).
In the light of the above information and in continuation with our ongoing research on the oxidation kinetics and mechanism of important organic substrates [13][14][15], we report here the results obtained on the oxidation kinetics of three cephalosporin drugs, namely, cephalexin (CEX), cefadroxil (CFL), and cephradine (CPD) with CAT in alkaline medium in order to elucidate the mechanism of these redox systems and also to assess their relative rates. The studies were extended to deduce the appropriate kinetic rate law and also to establish the isokinetic relationship through the computed thermodynamic parameters.

Experimental
2.1. Reagents. The substrates cephalexin, cefadroxil, and cephradine were of analytical grade of purity and were gifted by Orchid Chemicals Pvt. Ltd., Chennai, India. These drugs were used as received. Fresh aqueous solutions of substrates were prepared whenever required. Chloramine T (Merck) was purified by the method of Carrell Morris et al. [16]. An aqueous solution of CAT was prepared, standardized iodometrically and stored in amber-colored, stoppered bottles until further use. The concentrations of stock solutions were periodically checked. All other reagents were of accepted grades of purity. Heavy water (D 2 O 99.4%) was supplied by BARC, Mumbai, India, and was used to investigate the solvent isotope effect. Double distilled water was used throughout the work. The regression coefficient 2 was calculated using fx-100Z scientific calculator.

Kinetic Measurements.
All the kinetic runs were performed under pseudo-first-order conditions with a known excess of [substrate] o over [oxidant] o . The reactions were carried out in glass stoppered Pyrex boiling tubes whose outer surface was coated black to eliminate photochemical effects. Appropriate amounts of the substrate and NaOH solutions and water (to maintain a constant total 50 mL volume) were taken in the tube and thermostated at 301 K for thermal equilibrium. A known amount of CAT solution also thermostated at the same temperature was rapidly added to the mixture in the boiling tube. The mixture was periodically shaken to ensure uniform concentration, and the progress of the reaction was monitored by iodometric determination of unreacted CAT in a measured aliquot (5 mL each) of the reaction mixture at different intervals of time. The course of the reaction was studied for more than two half-lives. The pseudo first-order rate constants ( s −1 ), calculated from linear plots of log [CAT] versus time, were reproducible within 2-6%.

Reaction Stoichiometry and Product
Analysis. Different aliquots of reaction mixtures containing different concentrations of CAT and substrate with constant NaOH concentration (4.4 × 10 −3 mol dm −3 ) were equilibrated at 301 K for 24 h. Iodometric titration of unreacted CAT in the reaction  mixtures showed that one mole of each substrate consumed two moles of the oxidant, and the stoichiometry is given in Figure 6. The reactions in case of all the three substrates with CAT, separately in the stoichiometric ratio under stirred conditions in presence of NaOH (4.4 × 10 −3 mol dm −3 ), were allowed to progress for 24 h at 301 K. The progress of the reaction was monitored by TLC. After completion of the reactions, the products were neutralized with HCl. The solid product formed is filtered off and washed thoroughly with water and dried over sodium sulphate. The organic products were subjected to mass spectral analysis. The molecular ion peaks at 382.0 ( Figure 1 p-Toluenesulfonamide among the reaction product was extracted with ethyl acetate. It was detected by paper chromatography [13]. Benzyl alcohol saturated with water was used as the solvent with 0.5% vanillin in 1% HCl solution in ethanol as spray reagent ( = 0.905). It was also noted that no further oxidation of these products, under prevailing experimental conditions.

Results and Discussion
Our preliminary experimental studies revealed that the oxidation reactions of CEX, CFL, and CPD with CAT were facile in alkaline medium. Therefore, the oxidation of CEX, CFL   and CPD with CAT was kinetically investigated under pseudo first-order conditions with large excess of the substrate over oxidant at 301 K in NaOH medium. The same oxidationkinetic behaviour was observed for all the three substrates which are under the present study.

Effect of p-Toluenesulfonamide (PTS) Concentration on the
Rate of Reaction. Addition of the reduction product of CAT, p-toluenesulfonamide (2.0 × 10 −3 mol dm −3 ), does not have any pronounced effect on the rate of the reaction, indicating that it is not involved in the preequilibrium with the oxidant.

Effect of Halide Ions on the Rate of Reaction. Addition of
Cl − or Br − ions as NaCl or NaBr (5.0 × 10 −3 mol dm −3 ) had no effect on the rate, signifying that no interhalogen or free chlorine is formed in the reaction sequence.

Effect of Ionic Strength of the Medium on the Rate of
Reaction. The effect of ionic strength of the medium on the rate of reaction was carried out at 0.30 mol dm −3 using NaClO 4 solution, with all other experimental conditions held constant. It was found that the added NaClO 4 had negligible effect on the reaction rate, suggesting that a neutral molecule is involved in the rate-determining step. Subsequently, the ionic strength of the reaction mixture was not kept constant for kinetic runs.   Table 3.

Effect of Temperature on the Rate of Reaction.
The reaction rates were determined at different temperatures (291-306 K), keeping the other experimental conditions the same. Based on the Arrhenius plots of log versus 1/T ( 2 > 0.9985), activation energy and other thermodynamic parameters (Δ # , Δ # , Δ # , and log ) were computed for all the three substrates. All these results are summarized in Table 4.

Free Radical Test.
To test for free radical species in the reaction mixture, it was added to acrylamide. Polymerization did not occur, confirming the absence of free radical species in the reaction mixture. Appropriate control experiments were also run simultaneously.
If TsNHCl is the reactive species, retardation of rate by OH − would be expected according to (1), which is contrary to the experimental results. The absence of any significant effect of the added p-toluenesulfonamide on the rate rules out the involvement of OCl − and HOCl ions in the reaction sequence ( (2) and (3)). Further, Bishop and Jennings [6] have calculated the order of the concentrations of the various species present at different pH in 0.05 mol dm −3 solution of CAT, and a comparison with the concentration of species present in alkaline CAT solution would indicate that TsN − Cl is the likely oxidizing species. In the present study, the positive effect of [OH − ] reveals that the anion TsN − Cl is the most probable oxidizing species. Bearing the above facts in mind, the following mechanism (Scheme 1) is proposed for the oxidation of cephalosporins by CAT in alkaline medium.
Here, S is the substrate, and the structure of the intermediate complex species is depicted in Scheme 2 where a detailed mechanistic interpretation for the oxidation of cephalosporins by CAT in alkaline medium is illustrated. In step (i) of Scheme 2, the conjugate acid TsNHCl reacts with OH − in an alkali accelerating step to form an anion TsN − Cl. In the next slow and rds (Step From Step (i) of Scheme 1, or By substituting for [TsNHCl] from (7) into (5) and solving for [TsN − Cl], we get By substituting for [TsN − Cl] from (8) into (4), the following rate law (9) is obtained: The values in parentheses refer to second-order rate constants ( dm 3 mol −1 s −1 ).

Proton Inventory Studies.
The proposed mechanism is supported by the increase in the rate of reaction in D 2 O medium which implies [21] a fast preequilibrium hydroxyl ion transfer (Step (i) of Scheme 1). For a reaction involving a fast equilibrium H + or OH − ion transfer, the rate increases in D 2 O since D 3 O + and OD − ions are stronger acid and stronger base (2∼3 times greater), respectively, than H 3 O + and OH − ions [21,22]. The increase of reaction rate with D 2 O in the present case and the obtained solvent isotope effect (H 2 O) (D 2 O) < 1 conform to the above theory. The magnitude of isotope effect, however, is small, which can be attributed to the fractional-order dependence of rate on Further, the proton inventory studies could shed some light on the nature of the transition state. The variation of rate Here, is the rate constant in a mixed solvent which contains an atom fraction " " of deuterium, and is the rate constant in pure H 2 O. and are the isotopic fractionation factors for isotopically exchangeable hydrogenic sites in the transition state (TS) and reactant states (RS), respectively. Equation (12) allows the calculation of the fractionation factor of TS if the reactant fractionation factor is known.
Proton inventory plot of versus ( Figure 4; 2 = 0.9948; Table 3) in the present case is a straight line. A comparison with the standard plots [25] clearly indicates the involvement of a OH − ion or OH-D exchanges during the reaction sequence. A linear plot is indicative of a simple solvent isotope effect involving single exchangeable hydroxyl ion in the transition state, which is isotopically distinct from water. Hence, the participation of OH − ion in the formation of transition state is inferred. Furthermore, if it is assumed that the reaction proceeds through a single transition state, (12) can be transformed into (13): Equation (13) indicates a linear relation between ( / ) 1/2 and which is shown in (Figure 4; 2 = 0.9948).
The slope ( − 1) = −0.29, from which the fractionation factor is found to be 0.71. Kresge and Allred [26] obtained ISRN Physical Chemistry versus . Experimental conditions are as in Table 3.  Table 4. a value of 0.80 through NMR studies for the isotopic fractionation factor of OH − , which is confirmed by Gold and Grist [27]. Considering the diversity of procedures employed, there is an agreement between the present value and the values reported in the literature.

Dielectric Constant Effect.
The effect of changing solvent composition on the rate of a bimolecular reaction has been discussed by Kirkwood [28]. The theory of Kirkwood was further developed by Laidler and Landskroener [29] and Tanford and Kirkwood [30]. A good account of these and other theories has been given by Entelis and Tiger [31]. Further, for the limiting case of zero angle approach between two dipoles or an ion-dipole system, Amis [32] has shown that a plot of log versus 1/ gives a straight line with a negative slope for the reaction between a negative ion and a dipole or between the dipoles, while a positive slope results in a positive-ion dipole interaction. The negative dielectric effect observed in the present studies (Table 2) clearly supports the negative ion-dipole nature of the rds in the proposed scheme.

Ionic Strength Effect.
The proposed reaction mechanism is also evinced by the observed zero effect of ionic strength on the rate of reaction. The primary salt effect on the reaction rates has been described by Brønsted and Bjerrum theory [33]. According to this concept, the effect of ionic strength ( ) on the rate of a reaction involving two ions is given by the following relationship: where and are the reacting ions, and are the charges on the respective species, and are the rate constants in the presence and in the absence of the added electrolyte, respectively. Equation (14) shows that a plot of log versus 1/2 would be linear yielding a slope 1.02 and an intercept log . As the slope of the line depends on charges of the reacting ions, three special cases may arise: (i) if and have the same charges, will be positive and the rate constant increases with √ ; (ii) if and have opposite signs, will be negative and the rate constant decreases with √ ; and (iii) if either or is uncharged, is equal to zero and is independent of the ionic strength of the solution. In the present case, a negative charge and a neutral molecule are involved in the rate-determining step (Step (ii) of Scheme 2). Hence, variation of the ionic strength of the medium does not alter the rate, and it clearly conforms to the above theory (case (iii)).

Isokinetic Relationship.
Variation in rate within a reaction series may be caused by changes in both enthalpy of activation (Δ # ) and entropy of activation (Δ # ), but these quantities vary in a parallel fashion. These variables are correlated by a linear relationship: Δ # = Δ # + Δ # , where is the isokinetic temperature. It is the temperature at which the rate constant is identical for the reaction series. Table 4 shows that the activation energy is the highest for the slowest reaction and vice-versa, indicating that the reaction is enthalpy controlled. This is verified by calculating the isokinetic temperature ( ) from the slope of linear plot of Δ # versus Δ # ( Figure 5; 2 = 0.9990). The value of was found to be 378 K. Further, the validity of isokinetic relation has been done by the Exner criterion [34] by plotting log (301 K) versus log (291 K) , which gives a straight line ( Figure 5; 2 = 0.9913). The value of was calculated from the equation = 1 2 ( − 1)/( 2 − 1 ), where is the slope of the Exner plot; was found to be 370 K. Generally, in a chemical reaction, when there is a formation of a molecular complex or a linear relationship exhibits between enthalpy, and entropy one might observe what is known as enthalpy-entropy compensation. This can be expressed as where and are constants, usually called as compensation parameters. This type of compensation has been observed in various processes, and many models have been put forward to explain this complex phenomenon [35,36]. According to the transition state theory, the rate constant of the reaction which gives the product from the activated reaction is given as which can be rewritten as follows: By comparing (15) and (18), The enthalpy-entropy compensation points out that the variations of ΔH # in (18) are compensated by variations in ΔS # .
When the experimental temperature < , the reaction rate or equilibrium is mainly by the enthalpy change. In this region, the reaction with the lowest activation energy will react fast, and the interpretation involved in potential energy surfaces can be made. At temperatures above , however, the controlling factor is ΔS # , and interpretations based upon potential energy surfaces would obviously be in error. In general, it is found that the electronic effects are contained in the enthalpy factor and that many solvent effects are due to the entropy factor. Activation energy is the highest for the slowest reaction and vice-versa (Table 4), and the value of obtained is higher than the experimental temperature, which clearly indicates that the reaction is enthalpy controlled in the present case. This result also implies that the family of reactions proceeds through same mechanistic pathways.
3.17. Relative Reactivity of the Substrates. From the rate data (Table 4), the rate of oxidation of cephalosporins was found to follow the order cephradine > cefadroxil > cephalexin. This trend can be explained as in Figure 7.
In Structure I (cephradine), there is no cloud interaction with lone pair of electrons of sulphur atom of the substrate, which makes the sulphur atom more reactive to undergo oxidation at a faster rate. The Structure II (cefadroxil) might get converted into quinonoid forms (IIa) or (IIb) which results in less reactivity towards sulphur, and hence the rate of oxidation is slower than cephradine. In  Structure III (cephalexin), there is an involvement of cloud of the arene with lone pair of electrons of sulphur atom making it less reactive towards the oxidant. Hence, the rate of oxidation of cephalexin is the least when compared to cephradine (Structure I) and cefadroxil (Structure II).

Activation Parameters.
The proposed mechanism is also supported by the moderate values of energy of activation and other thermodynamic parameters. The large negative values of ΔS # indicate the formation of a more ordered, rigid associative transition state in each case. The high positive values of free energy of activation indicate that the transition state is highly solvated. The near constancy of ΔG # shows the operation of an identical mechanism in the oxidation of all the three substrates of the present study.

Conclusions
Based on the experimental results, the following conclusive remarks can be acquired.
Oxidation of all three drugs follows identical kinetics with a rate law: Rate = [CAT][S][OH − ] , and here, < 1. From the inspection of rate data, the rate of oxidation of drugs follows the order: cephradine > cefadroxil > cephalexin.
The stoichiometry of the reaction was found to be 1 : 2, and oxidation products were confirmed by mass spectral analysis. Activation parameters and isokinetic temperature indicates that the reaction is enthalpy controlled, and all the three drugs react with CAT via the same mechanism. A suitable scheme and relevant rate law have been worked out.