Advances in Biochemistry
Volume 4, Issue 1, February 2016, Pages: 1-8

Oxidation of Formamidines by Cerium(IV) in Aqueous Perchlorate Solutions: A Kinetics and Mechanistic Approach

Ahmed Fawzy1, 2, *, Ishaq Zaafarany1, Ismail Althagafi1, Moataz Morad1, Jabir Alfahmi1

1Chemistry Department, Faculty of Applied Science, Umm Al-Qura University, Makkah, Saudi Arabia

2Chemistry Department, Faculty of Science, Assiut University, Assiut, Egypt

Email address:

(A. Fawzy)
(I. Zaafarany)
(I. Althagafi)
(M. Morad)
(J. Alfahemi)

*Corresponding author

To cite this article:

Ahmed Fawzy, Ishaq Zaafarany, Ismail Althagafi, Moataz Morad, Jabir Alfahmi. Oxidation of Formamidines by Cerium(IV) in Aqueous Perchlorate Solutions: A Kinetics and Mechanistic Approach. Advances in Biochemistry. Vol. 4, No. 1, 2016, pp. 1-8. doi: 10.11648/j.ab.20160401.11

Received: March 6, 2016; Accepted: March 11, 2016; Published: March 24, 2016


Abstract: The kinetics of oxidation of one of the formamidine derivatives, namely, methylaminopyrazole formamidine (MAPF) by cerium(IV) has been investigated in aqueous perchlorate solutions at a constant ionic strength of 1.2 mol dm-3 and at 25°C. The progress of the reaction was followed spectrophotometrically. The reaction exhibited a first order kinetics with respect to [Ce(IV)], whereas the order with respect to [MAPF] was found to be less than unity. The reaction rate decreased with the increase in acid concentration with a negative less than unit order. The rate of reaction was not affected by increasing either ionic strength or dielectric constant of the medium. Furthermore, addition of cerium(III) product did not affect the reaction rate. The plausible oxidation mechanism involves formation of an intermediate complex between MAPF substrate and the kinetically active species of cerium(IV) in a pre-equilibrium step which decomposes in the slow step to give free radical derived from the substrate and Ce(III). The free radical is attacked by another Ce(IV) species to yield the final oxidation products which were identified by spectral and chemical analyses as methylaminopyrazole, dimethylamine and carbon dioxide. The activation parameters have been evaluated and discussed. The rate law associated with the reaction mechanism was derived.

Keywords: Oxidation, Methylaminopyrazole Formamidine, Cerium(IV), Kinetics, Mechanism


1. Introduction

Cerium(IV) is a powerful one electron oxidizing agent in acid media [1-18]. Oxidation by cerium(IV) in sulfuric acid solutions have conclusively been established [1-9]. Nevertheless, little reports on cerium(IV) oxidation in perchloric acid solutions [10-17], probably due to presence of dimers and polymers of cerium(IV) in such solutions [17]. However, the reactions of cerium(IV) in perchloric acid medium proceed much faster than those in sulfuric acid medium. Cerium(IV) oxidations of various substrates are found to follow different mechanisms, depending upon the acid medium used. Such oxidation are generally interesting to understand the different pathways of the reactions and different active forms of cerium(IV).

Formamidines have achieved considerable importance in the last decades due to their very broad spectrum of biological activity [19, 20]. The N,N-dialkyl derivatives of formamidines are highly effective acaricides and the most rewarding of these studies resulted in discovery of the acaricide insecticide chlordimeform [21, 22]. The oxidative cleavage of formamidines is quite important, since the N,N-dialkylformamidine group is one of the most versatile protecting groups, especially in biosynthetic applications. Also, methylaminopyrazole was employed as beta-sheet template to investigate its interaction with ferrocenoyl-dipeptides [23].

A literature survey revealed that there is no work has been reported about the kinetics and mechanism of oxidation of methylaminopyrazole formamidine (MAPF) by cerium(IV). This observation prompted us to investigate the title reaction. We aim to establish the optimum conditions affecting oxidation of MAPF by cerium(IV) in perchlorate solutions and to propose a plausible oxidation mechanism.

2. Experimental

2.1. Materials

All chemicals employed in the present work were of reagent grade and their solutions were prepared by dissolving the requisite amounts of the samples in doubly distilled water. The stock solution of MAPF was prepared as reported elsewhere [24]. Cerium(IV) solution was freshly prepared by dissolving ceric ammonium sulfate in a 1.0 mol dm-3 sulfuric acid, diluted with double distilled water and kept for overnight. Concentration of cerium(IV) was ascertained by titrating against standard ferrous ammonium sulfate solution using ferroin as an internal indicator. The solution of cerium(IV) was stored in a dark glass bottle and was used after 24 h, since the hydrolysis is negligible small or ruled out after 12 h of preparation [25]. Cerium(III) solution was prepared by dissolving cerium(III) acetate in water.

2.2. Kinetic Measurements

Kinetic measurements were followed under pseudo-first order conditions where MAPF was present in a large excess over that of cerium(IV). The progress of the reaction was followed by monitoring the decrease in absorbance of cerium(IV) as a function of timeat λ = 316 nm, its absorption maximum, whereas the other constituents of the reaction mixture do not absorb significantly at this wavelength. The applicability of Beer’s law for cerium(IV) at 315 nm has been verified giving ɛ = 4371 ± 28 dm3 mol1 cm1 in agreement with the earlier reports [11, 12]. The absorbance measurements were made in a thermostatted quartz cell of a pathlength 1.0 cm on a Shimadzu UV-VIS-NIR-3600 double-beam spectrophotometer. The effect of dissolved oxygen on the reaction rate was checked by monitoring the reaction in a nitrogen atmosphere. No significant difference between the results obtained under nitrogen and in the presence of air was observed.

First order plots of ln(absorbance) versus time were found to be good straight lines and the observed first order rate constant values (kobs) were calculated as the gradients of such plots. The rate constants were reproducible to within 4% and were the average of at least two independent kinetic runs. The reaction orders with respect to the different reactants were determined from the slopes of log kobs versus log(concentration) plots by varying the concentrations (C) of substrate and acid, in turn, while keeping other conditions constant.

3. Results

3.1. Stoichiometry and Product Analysis

Reaction mixtures containing various amounts of Ce(IV) and MAPF at constant [H+], ionic strength and temperature were allowed to react for 24 h in closed vessels for completion of reactions. The unreacted [Ce(IV)] was assayed spectrophotometrically at 316 nm as well as by titrating against standard ferrous ammonium sulfate solution. The results indicated that two moles of Ce(IV) were consumed by one mole of MAPF yielding the oxidation products as shown in the following equation,

The above stoichiometric equation is consistent with the results of product analysis as described elsewhere [24, 26]. The products, methylaminopyrazole (MAP) and dimethylamine (HNMe2) were identified by liquid chromatography and spot tests [27], respectively, and carbon dioxide by lime water.

3.2. Spectral Changes

The spectral changes during the oxidation of MAPF by Ce(IV) in aqueous perchlorate solutions are shown in Figure 1. The scanned spectra indicate gradual disappearance of Ce(IV) band at its absorption maximum with time as a result of its reduction to Ce(III). It can be observed that there is a change in the shape of Ce(IV) spectrum with a hypsochromic shift of its absorption maximum to a shorter wavelength.

[MAPF] = 5.0 x 10-3, [Ce(IV)] = 2.0 x 10-4, [H+] = 0.6 and I = 1.2 mol dm-3 at 25°C.

Figure 1. Spectral changes during oxidation of MAPF by Ce(IV) in perchloate solutions.

3.3. Dependence of the Reaction Rate on [Permanganate]

The oxidant, cerium(IV) was varied in the range of (1.0-8.0) x 10-4 mol dm3 keeping other variables constant. It has been observed that the increase in the oxidant concentration did not alter the rate constant value (Table 1). Also, the plots of ln(absorbance) versus time were found to be good straight lines. These results indicate that the order of reaction with respect to the oxidant concentration is confirmed to be one.

3.4. Dependence of the Reaction Rate on [MAPF]

The observed rate constant was determined at different initial concentrations of MAPF while other variables were kept constant. It was found that increasing [MAPF] increased the reaction rate as listed in Table 1. A plot of the kobs versus [MAPF] at constant pH is linear with a positive intercept as shown in Figure 2 confirming fractional-first order dependence with respect to [MAPF].

3.5. Dependence of the Reaction Rate on [H+]

Kinetic runs were carried out by varying the hydrogen ion concentration (0.2 – 1.2 mol dm–3) and keeping the concentrations of all other reactants constant. It was observed that the rate of the reaction decreased with increasing [H+] (Table 1). A plot of log kobs versus log [H+] was linear with a negative slope of -0.58 (Figure 3) confirming the negative fractional-first order dependence with respect to [H+].

Table 1. Effect of variation of [Ce(IV)], [MAPF], [H+] and ionic strength on the observed first order rate constant (kobs) in the oxidation of MAPF by Ce(IV) in perchlorate solutions at 25°C.

104 [Ce(IV)], mol dm-3 103 [MAPF], mol dm-3 [H+], mol dm-3 I, mol dm-3 105kobs, s-1
1.0 5.0 0.6 1.2 20.7
2.0 5.0 0.6 1.2 20.2
3.0 5.0 0.6 1.2 21.3
4.0 5.0 0.6 1.2 19.8
8.0 5.0 0.6 1.2 19.5
8.0 5.0 0.6 1.2 21.9
2.0 1.0 0.6 1.2 9.4
2.0 3.0 0.6 1.2 15.1
2.0 5.0 0.6 1.2 20.2
2.0 7.0 0.6 1.2 24.9
2.0 9.0 0.6 1.2 28.8
2.0 11.0 0.6 1.2 32.8
2.0 5.0 0.2 1.2 32.2
2.0 5.0 0.4 1.2 26.4
2.0 5.0 0.6 1.2 20.2
2.0 5.0 0.8 1.2 17.0
2.0 5.0 1.0 1.2 14.5
2.0 5.0 1.2 1.2 13.0
2.0 5.0 0.6 1.2 20.2
2.0 5.0 0.6 1.4 19.8
2.0 5.0 0.6 1.6 22.3
2.0 5.0 0.6 1.8 21.4
2.0 5.0 0.6 2.1 20.9
2.0 5.0 0.6 2.5 19.2

Experimental error ± 3%.

3.6. Effect of Ionic Strength and Dielectric Constant of the Reaction Medium

The effect of ionic strength on the rate of the reaction was studied by varying the ionic strength in the range 1.2 – 2.5 mol dm–3 using sodium perchlorate as an inert electrolyte with keeping the concentrations of all other reactants constant. The results indicated that variation of ionic strength was found to have negligible effect on the reaction rate.

The dielectric constant effect was studied by varying the acetic acid - water content in the reaction mixture with all other conditions being kept constant. The rate constant increased negligibly with decreasing the dielectric constants of the reaction medium.

[Ce(IV)] = 2.0 x 10-4, [H+] = 0.6 and I = 1.2 mol dm-3 at 25°C.

Figure 2. Plot of kobs versus [MAPF] in the oxidation of MAPF by Ce(IV) in perchlorate solutions.

[MAPF] = 5.0 x 10-3, [Ce(IV)] = 2.0 x 10-4 and I = 1.2 mol dm-3 at 25°C.

Figure 3. Plot of log kobs versus log [H+] in the oxidation of MAPF by Ce(IV) in perchlorate solutions.

3.7. Effect of Initially Added Product

The effect of added cerium(III) product was studied in the concentration range 1.0 x 10-4 to 1.0 x 10-3 mol dm-3 at constant concentrations of the oxidant, reductant and acid. It was found that Ce(III) did not have any significant effect on the rate of reaction.

3.8. Effect of Temperature

The rate of reaction was performed at four temperatures namely in the range of 288 - 318 K, at constant concentrations of the reactants and other conditions being constant. The results indicate that the observed first order rate constant was increased with rise in temperature. The activation parameters of the second order rate constant (k2) are calculated using Arrheniusand Eyringplots and are listed in Table 2.

Table 2. Activation parameters of the second order rate constant k2 in the oxidation of MAPF by Ce(IV) in perchlorate solutions. [MAPF] = 5.0 x 10-3, [Ce(IV)] = 2.0 x 10-4, [H+] = 0.6 and I = 1.2 mol dm-3.

Experimental error ±4%

3.9. Polymerization Test

The reaction mixture to which a known quantity of acrylonitrile scavenger has been added initially was kept in an inert atmosphere for about 8 h. Upon diluting the reaction mixture with methanol, a white precipitate has been formed, suggesting generation of a free radical in the reaction. When the experiment was repeated in the absence of MAPF under similar conditions, the test was negative. This indicates that the reaction was routed through a free radical path.

4. Discussion

It was reported [28-30] that the active species of cerium(IV) in perchlorate solutions was found to be either free monomeric species Ce4+, its hydrolyzed forms, Ce(OH)3+ and Ce(OH)22+ or partially in the form of dimeric species (Ce–O–Ce)6+ and (HOCe–O–CeOH)4+. The spectrophotometric studies [31] showed that the hydrated form, Ce4+, is the predominant species at [H+] 1.0 mol dm3 up to the concentration of 1.5 x 103 mol dm3 of cerium(IV), whereas the hydrolyzed forms and dimers are the more predominant at [H+] < 0.8 mol dm3. Therefore, under our experimental conditions of low [H+] and deceasing reaction rate with increasing [H+], Ce(OH)3+ may be regarded as the kinetically active form of cerium(IV) according to the following equilibria,

The present reaction between MAPF and Ce(IV) in aqueous perchlorate solutions has a stoichiometry of 1:2, i.e., one mole of MAPF requires two moles of Ce(IV). The reaction exhibited a first order dependence with respect to [Ce(IV)], less than unit order with respect to [MAPF] and a negative fractional order in [H+]. The rate is not considerably affected by Ce(III) suggesting that the probability of any fast equilibrium with the product preceding the rate-determining step was ruled out. The rate-determining step should be irreversible as is generally the case for one electron oxidants [32, 33] and the oxidation takes place through generation of free radical as obtained experimentally. The rate of reaction was not considerably affected by variation of either ionic strength or dielectric constant of the reaction medium suggesting that the reaction occur between a neutral molecule and a charged ion [34, 35].

Scheme 1. Mechanism of oxidation of MAPF by Ce(IV) in perchlorate solutions.

The less than unit order with respect to [MAPF] may be as a result of complex formation between the kinetically active Ce(IV) species and MAPF (C1) prior to the rate-determining step. Complex formation was proved kinetically by the non-zero intercepts of the plot of 1/kobs versus 1/[MAPF] (Figure 4) in favor of possible formation of an intermediate complex between the oxidant and substrate, similar to the well-known Michaelis–Menten [36] mechanism for enzyme–substrate reactions. An additional prove for complex formation is the change in the shape of Ce(IV) spectrum as well as the hypsochromic shift of its absorption maximum shown in Figure 1.The formed complex was slowly decomposed in the rate-determining step to give rise to the initial oxidation products as the substrate radical (MAPF.) and Ce(III). The substrate radical reacts with another Ce(IV) species in a subsequent fast step to yield an intermediate product. In a further fast step, the intermediate product is hydrolyzed to give the final oxidation products as given in Scheme 1. Since mechanistic Scheme (1) is in accordance with the generally well-accepted principle of non-complementary oxidations taking place in a sequence of one-electron steps, the reaction between the substrate and oxidant would afford radical intermediate as was obtained experimentally.

The suggested mechanism leads to the following rate law expression (see Appendix A),

(1)

The rate law (1) is consistent with all the observed orders with respect to different species.

Under pseudo-first order condition, the rate-law can be expressed by Eq. (2),

Rate = =kobs[Ce(IV)]            (2)

Comparing Eqs. (1) and (2), the following relationship is obtained,

(3)

Rearranging Eq. (3) leads to the following equations,

(4)

(5)

According to Eqs. (4) and (5), plots of 1/kobs versus 1/[MAPF] at constant [H+], and 1/kobs versus [H+] at constant [MAPF] should be linear with positive intercepts and are found to be so as shown in Figures 4 and 5, respectively. From the slopes and intercepts of these plots, the values of k1, KOH, and K could be evaluated and were found to be 4.71 x 10-4 s-1, 8.33 mol dm-3 and 4.21 x 103 dm3 mol-1, respectively, at 25°C.

[Ce(IV)] = 2.0 x 10-4, [H+] = 0.5 and I = 1.2 mol dm-3 at 25°C.

Figure 4. Plot of 1/kobs versus 1/[MAPF] in the oxidation of MAPF by Ce(IV) in perchloate solutions.

[MAPF] = 6.0 x 10-3, [Ce(IV)] = 2.0 x 10-4 and I = 1.2 mol dm-3 at 25°C.

Figure 5. Plot of 1/kobs versus [H+] in the oxidation of MAPF by Ce(IV) in perchlorate solutions.

The obtained activation parameters listed in Table 2 can be discussed as follows. The negative value of ∆S#suggest interaction of reacting ions to form an intermediate complex of inner-sphere nature [37, 38]. This value is within the range of radical reactions and has been ascribed to the nature of the electron pairing and unpairing process and to the loss of degree of freedom formerly available to the reactants upon formation of a rigid transition state [39, 40]. The values of DH and DSare both favorable for electron transfer processes. On the other hand, the positive values of both DH and DGindicate endothermic formation of the complex and its non-spontaneity, respectively.

Appendix A.

Derivation of the Rate Law Expression

According to the suggested mechanistic Scheme 1,

Rate =  = k1[C1]              (A1)

KOH= , [Ce(OH)3+] =     (A2)

,

[C1] = K[Ce(OH)3+][MAPF]=   (A3)

Substituting Eq. (A3) into Eq. (A1) leads to,

Rate =          (A4)

The total concentration of Py-F is given by:

[MAPF]T = [MAPF]F + [C1]             (A5)

where ‘T’ and ‘F’ stand for total and free concentrations.

Substituting Eq. (A3) into Eq. (A5) and rearrangement gives,

[MAPF]T = [MAPF]F +      (A6)

[MAPF]T = [MAPF]F(1+ )       (A7)

Therefore,

         (A8)

In view of low [Ce4+], the second denominator term KOHK[Ce4+]/[H+] in the above equation is neglected. Therefore,

[MAPF]F = [MAPF]T                 (A9)

Also,

[Ce(IV)]T = [Ce4+]F+ [Ce(OH)3+] + [C1]      (A10)

Substituting Eqs. (A2) and (A3) into Eq. (A10),

[Ce(IV)]T = [Ce4+]F+ +  (A11)

[Ce(IV)]T = [Ce4+]F (1 + + )    (A12)

         (A13)

Substituting Eqs. (A9) and (A13) into Eq. (A4) (and omitting ‘T’ and ‘F’ subscripts) leads to,

 (A14)

Under pseudo-first order condition, the rate-law can be expressed by Eq. (A15),

Rate = = kobs[Ce(IV)]        (A15)

Comparing Eqs. (A14) and (A15), the following relationship is obtained,

        (A16)

and with rearrangement, the following equations are obtained,

         (A17)

   (A18)


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