Kinetics, reactivity, initial-transition state analysis and thermodynamic parameters of base-catalyzed hydrolysis of coumalic acid in solvents with different polarities

Ezz A. Abu-Gharib; Rafat M. EL-Khatib; Lobna A.E. Nassr; Ahmed M. Abu-Dief

doi:10.1016/j.arabjc.2012.12.040

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Original article

10 (

1_suppl

); S988-S995

doi:

10.1016/j.arabjc.2012.12.040

Kinetics, reactivity, initial-transition state analysis and thermodynamic parameters of base-catalyzed hydrolysis of coumalic acid in solvents with different polarities

Ezz A. Abu-Gharib, Rafat M. EL-Khatib, Lobna A.E. Nassr, Ahmed M. Abu-Dief^⁎

Chemistry Department, Faculty of Science, Sohag University, 82534 Sohag, Egypt

⁎Corresponding author. Tel.: +20 1064162700; fax: +20 934601159. ahmed_benzoic@yahoo.com (Ahmed M. Abu-Dief)

Received: 2012-8-25, Accepted: 2012-12-30, Published: 2013-02-01

Disclaimer:
This article was originally published by Elsevier and was migrated to Scientific Scholar after the change of Publisher.

Peer review under responsibility of King Saud University.

Abstract

Base-catalyzed hydrolysis of coumalic acid (COU) in binary aqueous-methanol and aqueous-acetone mixtures has been studied kinetically at a temperature range from 283 to 313 K. Moreover, the change in the activation energy barrier of COU from water to water–methanol and water–acetone mixtures is estimated from the kinetic data. Solvent effects on reactivity trends have been analyzed into initial and transition state components by using transfer chemical potentials of the reactants and kinetic data. The transfer chemical potentials $δ_{m} μ^{θ}$ for COU⁻ anion are derived from solubility data from its calcium, cerium and lanthanum salts. The decrease in rate constant of the base hydrolysis reaction of COU as the percentage of methanol or acetone increases is dominated by transition state (TS). The base hydrolysis reaction of COU follows a rate law with k_obs = k₂[OH⁻] and the reaction mechanism was suggested. The high negative values of entropy of activation support the proposal mechanism, i.e. the investigated reaction takes place via the formation of an intermediate complex. Thus, the ring opening of the intermediate complex would be the rate controlling step.

Keywords

Base-catalyzed hydrolysis

Activation energy barrier

Transfer chemical potential

Initial state

Transition state

Isokinetic temperature

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1 Introduction

Coumalic acid is classified under the family of chromen-2-one derivatives. Chromen-2-ones have attracted an intense interest in recent years because of their diverse pharmacological properties. Among these properties, their cytotoxic effects were most extensively examined, and their broad range of effects on the tumors as shown by various in vitro and in vivo experiments and clinical studies were discussed (Kostova, 2005). Chromen-2-one derivatives have also found applications as fluorescent dyes (Brady and Shiek, 1984; Camur and Bulut, 2008), antitumor agents (Rosskoph et al., 1992), antioxidants (Kaneko et al., 2003), anti-inflammatory agents (Melagraki et al., 2009), antineoplastic agents (Lin et al., 2007), immunomodulant agents (Reddy et al., 2004) antifungals, anticoagulants, antibacterial, antimicrobial and insecticides. Moreover, they are used as proliferators of HIV and particular human malignant cell lines in vitro, as well as affecting tumor activity against several in vivo tumor types (El-Khatib and Nassr, 2007). Coumalic acid or its salt which is included in a tea beverage product is used to inhibit non-enzymatic browning of the tea beverage product.

Rate coefficient for the reaction between organic and inorganic compounds with hydroxide ions in aqueous solutions is sensitive to the nature and molar fraction of organic solvent. The dependence of the rate constants on the solvent comprising binary aqueous mixtures depends on the composition of the solvent (Blandamer et al., 1986; Marcus, 2007; Abu-Gharib et al., 2011a,b ).

Our work treats the effect of medium on the kinetics of base hydrolysis of COU in aqueous-methanol and aqueous-acetone mixtures. The influence of solvent composition on the rate constant was analyzed in terms of initial and transition states. This approach provides complementary information on the role of the solvation and how transition state modeling can assist in interpretation. Also the activation parameters of the mentioned reactions were calculated by least squares of Arrhenius and Eyring plots.

2 Experimental

2.1

2.1 Materials

All materials, sodium hydroxide (99%), sodium chloride (98%), sodium nitrate (97%), methanol, acetone were obtained from BDH. The stock solutions of NaOH and NaCl and NaNO₃ were prepared by dissolving the calculated amounts of AnalaR samples in bidistilled water. Coumalic acid (99%) was obtained from sigma.

2.2

2.2 Kinetic measurements

The kinetics of the base hydrolysis of coumalic acid was measured by following the dependence of absorbance on time using 10 mm silica cells using the thermostated cell compartment of JASCO model V-530 spectrophotometer thermally controlled at 25 °C ± 0.1 °C and at different temperatures in ultrathermostate (CRIOTERM 190). Chemical reactions were monitored in solutions held at constant ionic strength using appropriate amounts of sodium chloride for methanol and sodium nitrate for acetone over at least three half-life times. Kinetic measurements were performed so that the reaction follows pseudo-first-order kinetics in [COU], where [NaOH] is presented in large excess over [COU]. Rate constants were calculated from the dependence of absorbance on time at 319 nm. It was confirmed that there is no interference from other reagents at the selected wavelength absorption maxima for the investigated compound.

The kinetic of base-catalyzed hydrolysis was carried out on the coumalic acid but solubility measurements were carried out on its salts.

2.3

2.3 Solubility measurements

The sparingly soluble calcium, cerium and lanthanum salts of COU were prepared by double decomposition in aqueous-alcohol mixtures. Solubilities were measured by agitating a generous excess of solid with the appropriate solvent, aqueous-methanol and aqueous-acetone mixtures, in a thermostated vessel for 7 h. After allowing time for undissolved solids to settle by using centrifuge, an aliquot of saturated solution was removed and the solution diluted as necessary. Concentrations were measured spectrophotometrically at the wavelength of maximum absorbance (Blandamer et al., 1986; Abu-Gharib et al., 2011a,b ). All aliquots of equilibrated solutions were diluted with the solvent methanol or acetone and the dependences of λ_max on solvent composition, were found to be very small for COU, and thus ignored in calculation of solubilities. The spectra were measured at different sampling times to verify if a constant concentration has been achieved.

3 Results and discussion

The absorption band of COU at λ_max = 236 nm can be attributed π–π^∗ transition in the whole electronic system. Moreover, the absorption band at λ_max = 290 nm is due to n–π^∗.

3.1

3.1 Kinetic results and reaction mechanism

The repeated spectral scans of COU (cf. Fig. 1) in the presence of NaOH are characterized by the presence of isosbestic points which imply that the absorbance at those wavelengths remains constant throughout the whole base hydrolysis reaction. It can be realized from repeated spectral scans that the initial action on COU by NaOH takes place in one stage and leads to the rate determining opening of the pyrone ring and the formation of a salt of (2Z)-4-(hydroxymethylene)-2-pentenedioic acid at λ_max = 312 nm (El-Khatib and Nassr, 2007; Abu-Gharib et al., 2011a,b ). This compound is immediately converted into a salt of (2E)-4-(hydroxymethylene)-2-pentenedioic acid at λ_max = 319 nm as shown in the following scheme.

Repeated spectral scan of COU in aqueous-methanol mixture at [OH−] = 0.03 M, [COU] = 1 × 10−4 M, I = 0.05 M and 298 K with interval time = 2 min.

The observed first-order rate constants (k_obs) as a function on [OH⁻] and in the presence of different water–methanol and water–acetone ratios were calculated from the dependence of absorbance on time at λ_max for the investigated compounds using Microcal Origin program version 7.5 (cf. Tables 1 and 2).

Table 1 The observed first-order rate constant (10⁴ k_obs, s⁻¹), second order rate constant (10³ k₂, dm³ mol⁻¹ s⁻¹) and free energy of activation (δ_mΔG^# kJ mol⁻¹) values for the base hydrolysis of COU in various ratios (v/v) of MeOH in the presence of different [OH⁻] at 298 K and I = 0.05 M (NaCl).

[OH⁻], mol dm⁻³	MeOH% (v/v)
	0	10	20	40	50	60
	10⁴ k_obs, s⁻¹
0.02	14.00	10.50	8.32	5.73	4.22	3.20
0.025	17.80	13.30	10.52	6.60	5.13	3.67
0.03	21.00	16.60	12.60	8.30	6.30	4.42
0.035	24.10	18.50	14.43	9.20	7.00	5.00
0.04	28.00	21.50	16.50	11.10	8.23	5.80
0.045	31.10	23.70	18.70	12.30	8.90	6.40
k₂, dm³ mol⁻¹ s⁻¹10³	70.00	53.14	41.00	26.90	19.43	13.13
δ_mΔG^#, kJ mol⁻¹		0.68	1.33	2.37	3.18	4.15

Maximum error is 1%.

Table 2 The observed first-order rate constant (10⁴ k_obs, s⁻¹), second order rate constant (10³ k₂, dm³ mol⁻¹ s⁻¹) and free energy of activation (δ_mΔG^#, kJ mol⁻¹) values for the base hydrolysis of COU in various ratios (v/v) of acetone in the presence of different [OH⁻] at 298 K and I = 0.05 M (NaNO₃).

[OH⁻], mol dm⁻³	Acetone% (v/v)
	0	10	20	40	50	60
	10⁴ k_obs, s⁻¹
0.02	14.60	13.20	11.60	8.25	6.40	5.30
0.025	18.10	16.70	14.70	10.10	8.20	6.00
0.03	21.60	20.30	17.10	12.30	9.70	7.25
0.035	25.10	23.40	20.20	13.90	11.12	8.44
0.04	29.30	26.10	22.40	16.20	12.80	9.50
0.045	33.30	29.50	25.30	17.40	14.14	10.50
10³k₂, dm³ mol⁻¹ s⁻¹	72.90	64.90	54.60	37.80	31.00	22.60
δ_mΔG^#, kJ mol⁻¹		0.29	0.72	1.63	2.12	2.90

Maximum error is 1%.

Tables 1 and 2 shows that the values of k_obs decrease as the proportion of methanol or acetone increases. This may be due to the salting out of the hydrophilic OH⁻ and stabilization of the studied compound.

The observed enhancement of the rate constant values in acetone co-solvent compared with the corresponding values in methanol co-solvent would be ascribed to the destabilization of the hydrophilic OH⁻ ions in acetone more than in methanol.

The dependence of k_obs on base concentration is linear for COU in methanol and acetone without significant intercept (cf. Figs. 2 and 3) hence the hydrolysis follows the rate law with

(1)

k_{obs} = k_{2} [{OH}^{-}]

where k₂ is the second order rate constant at [OH⁻] ≫ [compound]. This equation indicates that the second order process is dominant in the present solvent mixtures. Values of k₂ are cited in Tables 1 and 2.

Plots of the observed first order rate constants of the reaction between NaOH and COU as a function of [NaOH] in different water–methanol mixtures at [COU] = 1 × 10−4 M, I = 0.05 M and 298 K.

Plots of the observed first order rate constants of the reaction between NaOH and COU as a function of [NaOH] in different water–acetone mixtures at [COU] = 1 × 10−4 M, I = 0.05 M and 298 K.

The change in the activation barrier δ_mΔG^# is evaluated and reported in Tables 1–4 from the ratio of rate constants of the base hydrolysis in the aqueous-solvent (k₂_S) to the corresponding values in the aqueous solution (k₂_W) according to the following relation [11]:

(2)

δ_{m} Δ G^{#} = - RT \ln (\frac{k_{2 s}}{k_{2 w}})

Table 3 Determination of transfer chemical potential of (COU^-) from solubility of its Ca²⁺ and La³⁺ sparingly soluble salts in aqueous-methanol mixtures at 298 K.

Compound	MeOH%	¹⁰⁴ Solubility.	δ_mμ^θ (salt) kJmol⁻¹	δ_mμ^θ (cation) kJmol⁻¹	δ_mμ^θ (anion) kJmol⁻¹
La(COU)₃ Є = 12465	0	40.90
Mol L⁻¹ cm⁻¹	20	17.93	8.17	3.70	1.49
λ_max = 234 nm	40	8.40	15.96	7.29	2.80
	60	4.93	21.00	10.56	3.48
Ca(COU)₂	0	43.45
Є = 12472	20	21.47	5.24	2.20	1.52
Mol L⁻¹ cm⁻¹	40	11.43	9.93	4.29	2.82
λ_max = 235 nm	60	8.20	12.42	5.42	3.50
Average	20				1.27
δ_mμ^θ(COU^-)	40				2.60
kJ mol⁻¹	60				3.49

Table 4 Determination of transfer chemical potential of (COU^-) anion from solubility measurements of its Ce³⁺ and La³⁺ sparingly soluble salts in aqueous-acetone mixtures at 298 K.

Compound	MeOH%	10⁴ Solubility	δ_mμ^θ (salt) kJ mol⁻¹	δ_mμ^θ (cation) kJ mol⁻¹	δ_mμ^θ (anion) kJ mol⁻¹
La(COU)₃	0	40.23
Є = 12465	20	6.78	17.65	−2.90^(a)	6.85
Mol L⁻¹ cm⁻¹	40	5.40	19.90	−15.80	11.90
λ_max = 234 nm	50	3.76	23.50	−26.60	16.70
Ce(COU)₃	0	38.25
Є = 12470	20	6.15	18.15	−2.1^(a)	6.75
Mol L⁻¹ cm⁻¹	40	4.15	22.00	−13.1	11.70
λ_max = 234 nm	60	2.67	26.40	−23.4	16.60
Average					6.80
δ_mμ^θ (COU⁻)					11.80
kJ mol⁻¹					16.65

It is observed that the values of δ_mΔG^# increase with the increasing of methanol or acetone content and these match with the decreasing of k_obs and k₂ values as methanol or acetone content increases.

On applying the steady-state approximation for the concentration of the intermediates B, C and D on the suggested reaction mechanism of COU, the total concentration of the compound: [A]T = [B] + [C] + [D], where [A] is very small and neglectable because it reacts very fast with the base to form the intermediate B. Thus, the rate equation can be formulated as:

(3)

Rate = \frac{- d [compound]}{dt} = \frac{k_{b} k_{d} [{OH}^{-}] {[A]}_{T}}{k_{b} k_{c} + k_{c} k_{d} + k_{b} k_{d} [{OH}^{-}]}

where k_{obs} = \frac{k_{b} k_{d} [{OH}^{-}]}{k_{b} k_{c} + k_{c} k_{d} + k_{b} k_{d} [{OH}^{-}]}

3.2

3.2 Michaelis–Menten kinetics

The base hydrolysis reaction of COU exhibits Michaelis–Menten kinetics (cf. Fig. 4). This suggests that the reaction occurs via the formation of intermediates as shown in the suggested mechanism (cf. Scheme 1). From Lineweaver–Burk plots, Michaelis–Menten constants (K_M) were calculated and found to be 0.43 mol dm⁻³ s⁻¹ for COU. The large values of K_M indicate the formation of intermediates during the reaction progress (Michaelis and Menten, 1918; Nassr, 2010; Awad et al., 2008 ).

Lineweaver–Burk plot of the base hydrolysis of COU in aqueous solution with [compound] = 1 × 10−4 mol dm−3, I = 0.05 M and 298 K.

The suggested mechanism of base-catalyzed hydrolysis of COU.

3.3

3.3 Initial state-transition state analysis

Solvent effects on the rate constants reflect solvation changes on the initial and transition states as these are transferred from water into mixed aqueous methanol or aqueous acetone media. If solvent effects on the initial state can be established from thermodynamic data (solubility data), then solvent effects on the transition state can be calculated from initial state and kinetic data assuming constancy of mechanism and validity of transition state theory.

The transfer chemical potentials for the constituent ions of a salt between two solvents are given by the following equation (Blandamer et al., 1986; Abu-Gharib et al., 2011a,b ).

(4)

δ_{m} μ^{θ} (M_{m} X_{n}) = m δ_{m} μ^{θ} (M^{n +}) + n δ_{m} μ^{θ} (X^{m -})

But there is no absolute method for calculating single ion values. Several procedures have been suggested, with concomitant debate about their validity. The almost identical assumptions we have used in the investigation are

δ_{m} μ^{θ}

(ph₄As⁺, ph₄P⁺) =

δ_{m} μ^{θ}

(ph₄B⁻), which has equal transfer chemical potentials for the two large, hydrophobic and slightly solvated ions [16].

In order to establish solvation effects on the chemical potential of COU⁻ anion, the solubilities of a suitable salt in the water–methanol and water–acetone mixtures are needed. Such a salt should be only sparingly soluble, to minimize activity corrections in calculating transfer chemical potentials. These salts are Ca(COU)₂, Ce(COU)₃ and La(COU)₃. Thus, the transfer chemical potentials of single cations: Ca²⁺, Ce³⁺ and La³⁺ are required.

The transfer chemical potentials of these cations were obtained from reference (). Then the mean of transfer chemical potentials of COU⁻ anion is deduced from the values of the transfer chemical potentials of sparingly soluble salts Ca(COU)₂, La(COU)₃ and Ce(COU)₃ in aqueous-methanol and aqueous-acetone mixtures (cf. Tables 3 and 4). Indeed the values of $δ_{m} μ^{θ}$ (COU⁻) which are calculated from its calcium and lanthanum salts in aqueous-methanol mixtures are in good agreement. Also the values of $δ_{m} μ^{θ}$ (COU⁻) which are calculated from its lanthanum and cerium salts in aqueous-acetone mixtures are the same.

The values of the transfer chemical potentials of OH⁻ anion in methanol are obtained from references (Abu-Gharib et al., 1984) and for OH⁻ in acetone from references Marcus (2007). Also the values of transfer chemical potentials of La³⁺, Ca²⁺, and Ce³⁺ are obtained from reference (Abu-Gharib et al., 2011a,b).

The change in second order rate constant on going from water to medium (aqueous -methanol or aqueous-acetone mixtures) of mole fraction x₂, δ_mΔG^# will be equal to ΔG^#(x₂) minus ΔG^#(x₂ = 0).

(5)

δ Δ_{m} G^{#} = δ_{m} μ^{*} - {δ_{m} μ^{θ} (A) + δ_{m} μ^{θ} (B)}

Where δ_m μ^θ(A) + δ_m μ^θ(B) are the transfer chemical potentials for the reactants A and B and δ_m μ^≠ is the transfer chemical potential of the transition state. In our work δ_mμ^θ of reactant can be obtained from the dependence of the ratio of the solubility of pure reactant in solvent to that in water.

The free energies of activation and transfer chemical potentials are now used in the analysis of reactivity trends with changing solvent composition into initial state and transition state components, i.e., combination of kinetic data, solubility data and transfer chemical potentials yield the effect of solvent on the transition state according to equation (5).

The initial state–transition state analysis of reactivity trends for COU⁻ anion in aqueous–methanol and aqueous–acetone mixtures is set out in Table 5 and depicted in Figs. 5 and 6.

Table 5 Initial state-transition state analysis of solvent effects on reactivity trends for base hydrolysis of COU^- anion in aqueous-methanol and aqueous-acetone mixtures at 298 K.

Solvent	Solvent%	δ_mμ^θ(COU⁻) kJ mol⁻¹	δ_mμ^θ(OH⁻) kJ mol⁻¹	δ_mμ^θ(IS) kJ mol⁻¹	δ_mΔG^# kJ mol⁻¹	δ_mμ^θ(TS) kJ mol⁻¹
MeOH	0
	20	1.51	−0.20	1.31	1.33	2.64
	40	2.81	0.10	2.91	2.37	5.28
	60	3.49	1.60	5.10	4.15	9.25
Acetone	0
	20	6.80	1.64	8.44	0.72	9.16
	40	11.80	2.45	14.25	1.63	15.88
	50	16.65	2.73	19.38	2.90	22.28

Plots of the initial state (IS) and transition state (TS) for the base hydrolysis of COU- anion in different ratios of MeOH at 298 K.

Plots of the initial state (IS) and transition state (TS) for the base hydrolysis of COU- anion in different ratios (v/v) of Ac at 298 K.

Table 5 shows that COU^- anion and the hydroxide ion are destabilized, resulting in the destabilization of the initial state as the proportion of methanol or acetone increases. However the transition state is significantly destabilized on increasing the methanol or acetone content (cf. Figs. 5 and 6). These results affect the decrease of the rate constant of COU⁻ anion i.e. the dominant effect of (TS) on the rate constant of the base hydrolysis of COU⁻ anion as the co-solvent methanol or acetone increases.

3.4

3.4 Determination of the activation parameters

The activation parameters were calculated (cf. Tables 6 and 7) by the least squares of Arrhenius plot (Fig. 7) and Eyring plot (Fig. 8).(see Fig. 9).

Table 6 Second-order rate constant values (10³ k₂, mol⁻¹ dm³ s⁻¹) and the activation parameters of the base hydrolysis of COU in various ratios (v/v) of MeOH at different temperatures at I=(0.05 M).

MeOH%	T (K)
	283	288	293	298	303	308	313	E_a kJ mol⁻¹	ΔH^≠k J mol⁻¹	ΔG^≠ kJ mol⁻¹	ΔS^≠ J mol⁻¹ K⁻¹	A mol⁻¹ dm³ s⁻¹
COU	10³ k₂, mol⁻¹ dm³ s⁻¹
0	24.77	34	49.5	70.0	101.3	....	....	39.25	37.8	67.1	−98.19	5.11 × 10⁸
10	18.35	25.1	37.2	53.14	75.8	104.6	153.6	40.42	38.9	67.35	−95.48	4.15 × 10⁸
20	15.36	20.0	30.7	41.0	62.54	81.0	123	42.18	40.6	68.3	−92.93	3.9 × 10⁸
40	10.43	13.7	20.74	26.90	41.49	55.16	82.95	44.73	43.4	69.7	−88.17	8.45 × 10⁷
50	7.39	9.5	14.43	19.43	28.65	33.61	57.29	46.19	44.7	70.15	−85.39	5.76 × 10⁷
60	4.21	6.45	8.87	13.13	17.76	26.63	35.22	47.45	46.1	70.76	−82.75	1.97 × 10⁷

Table 7 Second-order rate constant values (10³ k₂, mol⁻¹ dm³ s⁻¹) and the activation parameters of the base hydrolysis of COU in various ratios (v/v) of acetone at different temperatures at I = (0.05 M).

	T (K)
Acetone %	283	288	293	298	303	308	313	E_a kJ mol⁻¹	ΔH^≠ kJ mol⁻¹	ΔG^≠ kJ mol⁻¹	ΔS^≠ J mol⁻¹ K⁻¹	A mol⁻¹ dm³ s⁻¹
COU	10³ k₂, mol⁻¹ dm³ s⁻¹
0	24.4	36.9	49.0	72.9	....	....	...	36.81	35.3	65.3	−100.6	2.1 × 10⁹
10	22.2	32.4	44.0	64.9	89.0	....	...	38.12	36.5	65.73	−98.1	1.7 × 10⁹
20	18.23	27.0	36.0	54.6	73.0	110	144	39.78	38.2	66.6	−95.2	1.32 × 10⁹
40	12.90	19.0	25.4	37.8	50.0	76.0	106	41.33	41.1	68.1	−90.6	7.31 × 10⁸
50	10.60	15.4	20.0	31.0	41.0	61.7	80	43.79	42.3	68.5	−87.8	5.53 × 10⁸
60	7.44	11.1	14.7	22.6	30.0	45.0	59	44.55	43.5	68.92	−85.3	3.8 × 10⁸

Plots of lnk2 against 1/T for the base hydrolysis of COU in different ratios of MeOH.

Plots of-ln(k2h/kT) against 1/T for the base hydrolysis of COU in different ratios of MeOH.

Isokinetic plot of chromene-2-one derivatives in aqueous solution. HC = 7-hydroxy-2H-chromen-2-one, HCA = 7-hydroxy-2H-chromen-2-one-4-acetic acid and DMAC = 7-dimethylamino-4-methyl-2H-chromen-2-one were obtained from Ref. (Abu-Gharib, 2011).

It is worth mentioning that the activation parameters are important to determine the rate of reaction and reaction mechanism. The high negative values of entropy of activation (Tables 6 and 7) support the proposal mechanism, that is the investigated reaction takes place via the formation of an intermediate complex (El-Khatib and Nassr, 2007; Awad et al., 2008; Abu-Gharib et al., 2011a,b ), cf. the mechanism. Moreover, these values refer to the rigidity of the complex (B) and its high stability. Thus, the ring opening of the intermediate complex would be the rate controlling step (RCS) and proceeds with an inner-sphere electron transfer step.

It is interesting that, the different thermodynamic functions are consistent in their trends. As entropy of activation (ΔS^≠) increases, i.e., changes to less negative value, the rate constant decreases and activation energy increases. This behavior can be ascribed to enhancing stability of activated complex intermediate. Again the relatively high free energy change would assume that the slow step (RCS) is the ring opening of the established intermediate (B) and many vibrational degrees of freedom have been transformed into translations (Abu-Gharib et al., 2011a,b; Jagan, 2009 ). The large frequency factor would suggest synergistic evidence for the proposal pathway. Furthermore, the positive free energy of transfer from water to methanol or acetone reported in Tables 6 and 7, assumes that the transient species in hand are polar entities [El-Khatib and Nassr, 2007; Abu-Gharib et al., 2011a,b ].

The values of activation enthalpies (ΔH^≠) and activation entropies (ΔS^≠) are important in controlling the reaction rate. The values of ΔH^≠ are plotted versus the values of ΔS^≠ (cf. Fig. 8) and a straight line is obtained. The isokinetic temperature (β) was determined by least square of the slope and found to be 305 K, which is greater than the average experimental temperature (T_exp = 298 K) indicating that the reaction rate is enthalpy controlled (Karunakaran and Chidambaranathan, 2000; Karunakaran and Chidambaranathan, 2001a,b). The straight line obtained, show that the base hydrolysis of COU and other chromen-2-one derivatives followed one mechanism, i.e. the same rate determining step (R. D. S.).

4 Conclusion

Base-catalyzed hydrolysis of coumalic acid follows a rate law with k_obs = k₂[OH⁻]. The decrease in the rate constants of coumalic acid as the proportion of methanol or acetone increases is due to the destabilization of OH^- ion. The values of rate constants (k_obs and k₂) decrease in the following order water > acetone > methanol with increasing the methanol or acetone content. The dominant effect on the rate constant of the base hydrolysis of COU as the co-solvent methanol or acetone increases is the initial state. The high negative values of entropy of activation support the proposal mechanism, i.e. the investigated reaction takes place via the formation of an intermediate complex. Moreover, these values refer to the rigidity and stability of the intermediate complex .Thus, the ring opening of the intermediate complex would be the rate controlling step.

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Jagan, V., 2009. Osmania University, India, 18, 4, 89.
Kaneko T., Baba N., Matsuo M., . Chem. Biol. Interact.. 2003;142:239.
Karunakaran C., Chidambaranathan V., . Mon. für Chemie. 2001;131:1123.
Karunakaran C., Chidambaranathan V., . J. coroatica chemica acta. 2001;74:51.
Kostova I., . Curr. Med. Chem., Anti-Cancer Agents. 2005;5:29.
Lin R., Chiu G., Yu Y., Li P.J.S., Lu Y., Adams M., Fuentes-Pesquera A.R., Ema-nuel S.L., Greenberger L.M., . J. Bioorg. Med. Chem. Lett.. 2007;17:4557.
Marcus Y., . Chem. Rev.. 2007;107(9):3880.
Melagraki G., Afantitis A., Igglessi-Markopoulou O., Detsi A., Koufaki M., Kontogiorgis C., Hadjipavlou-Litina D.J., . Euro. J. Med. Chem.. 2009;44:3020.
Michaelis L., Menten N.L., . Biochem. Z.. 1918;49:333.
Nassr Lobna.A.E., . Int. J. Chem. Kinet.. 2010;42:372.
Reddy N.S., Gumireddy M.R.S.K., Bell S.C., Reddy E.P., . J. Bioorg. Med. Chem. Lett.. 2004;14:4093.
Rosskoph F., Kraus J., Franz J.G., . Pharmazie. 1992;47:139.

Introduction

Experimental

Materials
Kinetic measurements
Solubility measurements

Results and discussion

Kinetic results and reaction mechanism
Michaelis–Menten kinetics
Initial state-transition state analysis
Determination of the activation parameters

Conclusion

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[9] Jagan, V., 2009. Osmania University, India, 18, 4, 89.

[10] Kaneko T., Baba N., Matsuo M., . Chem. Biol. Interact.. 2003;142:239.

[11] Karunakaran C., Chidambaranathan V., . Mon. für Chemie. 2001;131:1123.

[12] Karunakaran C., Chidambaranathan V., . J. coroatica chemica acta. 2001;74:51.

[13] Kostova I., . Curr. Med. Chem., Anti-Cancer Agents. 2005;5:29.

[14] Lin R., Chiu G., Yu Y., Li P.J.S., Lu Y., Adams M., Fuentes-Pesquera A.R., Ema-nuel S.L., Greenberger L.M., . J. Bioorg. Med. Chem. Lett.. 2007;17:4557.

[15] Marcus Y., . Chem. Rev.. 2007;107(9):3880.

[16] Melagraki G., Afantitis A., Igglessi-Markopoulou O., Detsi A., Koufaki M., Kontogiorgis C., Hadjipavlou-Litina D.J., . Euro. J. Med. Chem.. 2009;44:3020.

[17] Michaelis L., Menten N.L., . Biochem. Z.. 1918;49:333.

[18] Nassr Lobna.A.E., . Int. J. Chem. Kinet.. 2010;42:372.

[19] Reddy N.S., Gumireddy M.R.S.K., Bell S.C., Reddy E.P., . J. Bioorg. Med. Chem. Lett.. 2004;14:4093.

[20] Rosskoph F., Kraus J., Franz J.G., . Pharmazie. 1992;47:139.

Kinetics, reactivity, initial-transition state analysis and thermodynamic parameters of base-catalyzed hydrolysis of coumalic acid in solvents with different polarities

Abstract

Keywords

Base-catalyzed hydrolysis

Activation energy barrier

Transfer chemical potential

Initial state

Transition state

Isokinetic temperature

1 Introduction

2 Experimental

2.1 Materials

2.2 Kinetic measurements

2.3 Solubility measurements

3 Results and discussion

3.1 Kinetic results and reaction mechanism

3.2 Michaelis–Menten kinetics

3.3 Initial state-transition state analysis

3.4 Determination of the activation parameters

4 Conclusion

References

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