Volume 4, 2019 - Issue 2

ISSN(Print): 2377-8091

ISSN(Online): 2377-8083

#### Reading: Orientation on Electron-Transfer Nature for Oxidation of Some Water-Soluble Carbohydrates: Kinetics and Mechanism of Hexacholroiridate (IV) Oxidation of Methyl Cellulose in Aqueous Perchlorate Solutions

Citation

Refat Hassan, Samia Ibrahim; Orientation on Electron-Transfer Nature for Oxidation of Some Water-Soluble Carbohydrates: Kinetics and Mechanism of Hexacholroiridate (IV) Oxidation of Methyl Cellulose in Aqueous Perchlorate Solutions, Trends Journal of Sciences Research, Volume 4, Issue 2, January 22, 2019, Pages 68-79. 10.31586/Chemistry.0402.04

Research Article

Open Access

Peer Reviewed

Article Contents

## Orientation on Electron-Transfer Nature for Oxidation of Some Water-Soluble Carbohydrates: Kinetics and Mechanism of Hexacholroiridate (IV) Oxidation of Methyl Cellulose in Aqueous Perchlorate Solutions

Trends Journal of Sciences Research, Volume 4, Issue 2, 2019, Pages 68–79.

https://doi.org/10.31586/Chemistry.0402.04

Received October 24, 2018; Revised January 12, 2019; Accepted January 22, 2019;
Published February 01, 2019

### Abstract

The kinetics and mechanism of oxidation of methyl cellulose (MC) polysaccharide by hexacholroiridate (IV) in aqueous perchlorate solutions at a constant ionic strength of 0.1 mole dm-3 has been investigated, spectrophotometrically. The experimental results showed first-order dependence in [IrCl6]2- and fractional first-order kinetics with respect to the MC concentration. The reaction rate was found to increase with decreasing the [H+]. A kinetic evidence for formation of 1:1 intermediate complex was revealed. The reaction kinetics seems to be of considerable complexity where one chloride ion from hexacholoiridate (IV) oxidant may act as a bridging ligand between the oxidant and the substrate within the formed intermediate complex. However the added chloride ions and oxidation products were found to have negligible effects on the reaction rate, the added acrylonitrile indicated the intervention of free-radical mechanism during the oxidation process. The kinetic parameters have been evaluated and a tentative reaction mechanism consistent with the kinetic results is discussed.

### 1. Introduction

Methyl cellulose (MC) is a cellulose ether derivative that is water-soluble due to the methylation of hydroxyl moieties which prevent extensive hydrogen bonding. It is hydrophilic macromolecule unless the temperature exceeds that of the lower critical temperature of solution (LCST) of the approximate range 40-70οC 1. Therefore, this material is expected to have advantageous for enzyme immobilization and, hence, methyl cellulose is an important macromolecule which used for preparing the nanofibrous mat through selecting the proper electro-spinning conditions 2.

The kinetics of oxidation of methyl cellulose (MC) polysaccharide by permanganate ion as multi-equivalent oxidant in alkaline solutions 3, 4, 5, showed that the oxidation reaction was proceeding through two distinct stages without free-radical intervention. On the other hand, oxidation of this MC substrate by alkaline hexacyanoferrate (III) ion as one-equivalent oxidant 6 indicated the intervention of free-radical mechanism through a simple one-step reaction-path. Kinetic evidences for formation of 1:1 intermediate complexes through those two catalyzation oxidation in either acid or base media were revealed.

Although, hexachloroiridate(IV) is known as inert oxidant of one-equivalent nature and has been widely used for oxidation of most organic and inorganic substrates, it seems that no attention has been paid to the oxidation of MC macromolecule by this oxidant.

Preliminary experiments on the oxidation of MC by [IrCl6]2- indicated that the oxidation reaction kinetics is complex whereas the influence of added acid showed an acid-inhibition of the rate constants.

In view of the above discrepancies and our interest on the kinetics of oxidation of some macromolecules by various oxidants in acidic media 5, 7, 8, 9, 10, 11, 12, 13, 14 along with the oxidation of some organic and inorganic substrate by this oxidant 15, 16, 17, 18, the present work seems to merit an investigation with the aims at shedding more highlights on the kinetics and mechanistics of oxidation of macromolecules in terms of electron-transfer process as well as to examine the influence of oxidant nature and solvent type on the oxidation kinetics. The results obtained may gain us some univocal information on the chemistry of MC carbohydrate in aqueous acidic media with synthesize of novel keto-biodegradable polymeric precursor derivatives as chelating agents.

### 2. Experimental

##### 2.1. Materials

Sodium hexacholoriridate (IV) of Analar quality (Wako Chemical Reagent) was used without further purification. The preparation and standardization of the stock solutions of [IrCl6]2- were the same as described elsewhere 14, 15, 16, 17, 18, 19, 20.

Methyl cellulose (MC) used in this work is (MP Biomedicals, LLC, France) and was used without further purification. An average molecular weight was about 86000 Dalton which calculated from the viscosity measurements of 2 % solution at 20 C (the intrinsic viscosity was 750 ml/g and the apparent viscosity was 4000 mPas). The average degree of polymerization (DP) was 460. A stock solution was prepared as described earlier 3, 4, 5. This process takes place by stepwise addition of the reagent powder to deionized water whilst vigorously stirring the solution to prevent the formation of coagulate lumps, which swell with difficulty.

All other reagents were of analytical grades. Doubly distilled conductivity water was used in all preparations. The temperature was controlled within ± 0.05 ˚C.

The ionic strength of the reaction mixtures was maintained constants at 0.1 mol dm-3 using NaClO4 as a non-complexing agent.

##### 2.2. Kinetic Measurements

A large excess of MC over that of [IrCl6]2- was employed in order to maintain the pseudo-first-order conditions throughout the kinetic measurements. The procedure of measurements was the same as described elsewhere 3, 4, 5, 16, 17, 18, 19, 20. The zero time was taken when half of [IrCl6]2- solution had been added to MC solution into the reaction cell. The progress of the reaction was followed by measuring the decrease in absorbance of [IrCl6]2- at its absorption maximum, 489 nm, where all other reagents of the reaction mixture do not absorb significantly at this wavelength, as a function of time. The applicability of beer`s law for [IrCl6]2- at 489 nm has been verified giving ε = 4050 ± 20 dm3 mol-1 cm-1 in good agreement with the values reported elsewhere 16, 17, 18, 19, 20, 21. The absorbance measurements were made in a thermostated cell compartment at the desired temperature within ± 0.05˚C on a Shimadzu UV-2101/3101 PC automatic scanning double beam spectrophotometer fitted with a wavelength program controller using cells of pathlength 1.0 cm. The spectral changes during the progress of the oxidation reaction are shown in Figure 1.

### 3. Results

##### 3.1. Stoichiometry and Products Analyses

Since the kinetics of this redox reaction seems to be of complexity nature, determination of the reaction stoichiometry becomes of great importance. Reaction mixtures containing different concentrations of the reactants at [H+] = 1 x 10-4 and I = 0.1 mol dm-3 were equilibrated in dark bottles away from light. The unreacted [IrCl6]2- was estimated periodically until it reaches a constant value. The experimental results revealed that 1.0 mole of MC consumed 4.0 ± 0.1 mol of the oxidant. This result conforms to the following stoichiometric equation.

where C7H12O5 and C7H8O5 denote MC and its corresponding diketo-derivative, respectively. The oxidation product was identified by the spectral data and elemental analysis as described elsewhere 22, 23, 24. The keto-derivatives were characterized by the formation of 2,4-dinitrophenylhydrazone and dioxime derivatives as well as by the IR absorption bands at 1760-1730 cm-1 that characterized to the carbonyl group of α-diketones 25, 26. The enhancement of the absorption band of OH group in the IR spectra of the product indicated the oxidation of OH groups of methyl cellulose to its corresponding ketones.

The formed biopolymer precursor oxidation product could be used as biocatalyst in immobilization systems as well as to encapsulate, protect and deliver bioactive or functional components such as minerals, peptides, proteins, enzymes, drugs, lipids and dietary fibers.

Figure 1.
Spectral changes (200 - 700 nm) during the formation of intermediate complexes in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [MC] = 0.2, [H+] = 0.6x10-4 and I =0.1 moldm-3 at 40 oC (scanning time intervals = 4 min).
##### 3.2. Dependence of the Reaction Rates on [IrCl6]2- and [MC]

Plots of ln (absorbance) vs. time were found to be linear for more than two-half-lives of reaction completion. This result indicated that the reaction is first-order in [IrCl6]2-. The first-order dependency was confirmed not only by the observed linearity of pseudo-first-order plots but also by the independence of the obtained rate constant on different initial concentrations of the oxidant used ranging between 1x10-4 and 6x10-4 moldm-3. Again some experiments were performed in the presence of iridium (III) as an oxidation product. The result indicates that [IrCl6]3- does not inhibit the reaction rate and, hence, the reversible reaction between MC and iridium (IV) is unlikely. The values of the pseudo first-order rate constants, kobs, can be evaluated from the gradients of ln (absorbance)-time plots. These values were calculated by the method of least -squares and are summarized in Table 1.

The order with respect to MC was deduced from the measurements of the reaction rates at several initial concentrations of MC and fixed concentrations of all other reagents. The non-constancy of the second-order rate constants derived from dividing the observed pseudo first-order rate constants by the initial [MC]0 shown in Table 1 indicates that the rate is fractional-first order in [MC]. The magnitude of the order was calculated from the double logarithm of the observed rate constants and the concentrations of MC (kobs = [MC]n ). Again, when the reciprocals of the observed rate constants were plotted against the reciprocals of the substrate, 1/kobs vs. 1/[MC], straight lines with distinct positive intercepts on 1/kobs axis were observed. This behavior was found to be indicative to the Michaelis-Menten kinetics for formation of 1:1 intermediate complexes. A typical plot is shown in Figure 2.

Table 1. The values of the observed first-order rate constants in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [MC] = 0.2, I = 0.1 mol dm-3 at 40 oC.
Figure 2.
A reciprocal Michaelis-Menten plot in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [H+] = 0.6x10-4 and I = 0.1 moldm-3 at 40oC.
##### 3.3. Dependence of the Reaction Rates on [H+]

The effect of [H+] on the rate constants at constant ionic strength of 2 x10-4 mol dm-3 and constants of all other reagents concentration was studied to elucidate a suitable reaction mechanism. An increase in [H+] was found to inhibit the oxidation rates. Plots of kobs vs. [H+]-1 gave curved lines passing through the origin as shown in Figure 3. On the other hand, plots of 1/kobs vs. [H+] were found to be linear with positive intercept on 1/kobs axis as shown in Figure 4. A reverse fractional- first- order in [H+] was revealed from (log kobs- log [H+]-1 plots).

##### 3.4. Dependence of the Reaction Rates on Ionic Strength

The influence of the ionic strength on the reaction rate at constant [H+] as NaClO4 concentration increased up to 0.6 moldm-3 has been examined. The results obtained indicated that the effect of the ionic strength on the reaction rates was negligible under our experimental conditions. The observed pseudo-first-order constants at [MC] = 0.2, [IrCl6]2- =2x10-4 , [H+] = 0.6x10-4 moldm-3 and 40 C were found to be 1.60, 1.58,1.62 and 1.59)x10-4 s-1 at ionic strengths of 0.1, 0.2, 0.4 and 0.6 moldm-3, respectively.

##### 3.5. Polymerization Test

The possibility of formation of free-radicals was examined by adding 10 % (v/v) acrylonitrile to the partially oxidized reaction mixture. After a lapse of 15 min mixing (on warming), a copious precipitate was observed indicating the intervention of the free-radical mechanism in the oxidation reaction.

##### 3.6. Dependence of the Reaction Rates on [KCl] and [IrCl6]3-

In order to investigate either the aquation of [Ir(Cl5) H2O] or [IrCl6]2- is the reactive species of the oxidant in the rate-determining step. Therefore, different concentrations of KCl were added to the reaction mixtures. Again, in order to examining the reaction reversibility, amount of the product [IrCl6]3- were added to the reaction mixture. It was found that the additions of either those two reagents to the reaction mixtures have no influence on the reaction rates.

Figure 3.
Plots of kobs vs [H+]-1 in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [MC] = 0.2, I = 0.1 mol dm-3 at various temperatures and hydrogen ion concentrations.
Figure 4.
Plots of 1/kobs vs. [H+] in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [MC] = 0.2, I = 0.1 mol dm-3 at various temperatures and hydrogen ion concentrations.

### 4. Discussion

A variety of reaction mechanisms was recognized in redox reactions involving hexachloroiridate (IV) as of one-equivalent nature 27, 28, 29, 30, 31. Some of these reactions tend to proceed through formation of intermediate complexes of inner-sphere nature 14, 15, 21, outer-sphere type 32 or by both inner- and outer-sphere 33 mechanisms via free-radicals intervention. Other reactions are proceeding through outer-sphere mechanisms of non-free-radicals intervention 34, 35, 36, 37, 38, 39, 40, 41, 42, whereas the formation binuclear complexes 16, 17 and ion-pairs 20 have been postulated for some other redox systems.

Polysaccharides which containing primary (R-CH2-OH)n, secondary (R-CH-OH)n (where R- is the macromolecule monomer) or those both alcoholic functional groups in their macromolecular chains are well-known to have high tendency for protonation in acidic solutions to form more reactive alkoxnium ions (R-CHOH2+) 5, 7, 8, 9, 10, whereas in alkali they form the reactive alkoxides (R-CH-O- ) by deprotonation 3, 4, 9, 10, 11, 13.

The decrease of the rate constants with increasing the hydrogen ion concentration was surprising. Hence, this behavior can be explained by a release of protons by hydrolysis or ionization processes prior to the rate- determining steps. Since, hexachloroiridate (IV) is known to be extremely inert 21, the release of proton from the oxidant is excluded. Again under our experimental conditions of lower [H+], the protonation of MC seems to be difficult or negligible small and, hence, it is also excluded. Consequently, the observed inverse fractional-order in [H+] in the present oxidation reaction may indicate the existence of a reaction-path involving a release of one proton prior to the rate-determining step.

In view of the aforementioned arguments and under the experimental conditions used of lower acid concentrations, a mechanism in consistent with the observed kinetic results which may be suggested involves the attack of the oxidant to the center of the substrate forming an intermediate complex (C1) with releasing a proton,

followed by decomposition of the formed complex (C1) in the rate-determining step to give rise to a substrate free-radical and hexachloroiridate (III) as initial oxidation products.

where K1 is the formation constant of the complex, the symbol (Ox) represents to the [IrCl6]2- , (Red) denotes the [IrCl6]3- and S is the substrate, respectively. This suggestion is supported by the released protons in the stoichiometric reaction defined by Eq. (1), as well as obeying the oxidation reaction to the Michaelis-Menten kinetics (Figure 2). The change of the rate constants with the change in the hydrogen ion and substrate concentrations can be expressed by the following rate-law equation

where [Ox]T is the total analytical concentration of oxidant. Under pseudo first-order conditions of the presence of a large excess of substrate [S] over that of the [Ox], Eq. (5) can be rewritten in the following form

$\frac{\text{1}}{{\text{k}}_{\text{obs}}}=\frac{\left[{H}^{+}\right]}{{\text{kK}}_{\text{1}}\text{[S]}}+\frac{\text{1}}{\text{k}}$

The rate law expression (6) requires that either 1/kobs -1/[S] at constant [H+] or 1/kobs-1/[H+] at constant [S] plots to be linear with positive intercept on 1/kobs axes. The experimental results were found to satisfy this requirement. However, this relationship did not agree with the kinetic results since the magnitudes of the intercepts obtained from those two plots (which are corresponding to the elementary rate constant (k)) were found to be quite different at the same temperature. Therefore, an alternative mechanism could be suggested. It involves the formation of an intermediate complex (C1) at first. Then, some rearrangement of that complex was occurred with releasing a hydronium ion to give the more reactive complex (C2) as follows,

In a similar manner of the derivation of the above mechanism the following rate-law expression is obtained,

According to Eq. (9 ) plots of either 1/kobs vs. 1/[S] at constant [H+] or 1/kobsvs.1/[H+] at constant [S], respectively, gave straight lines with positive intercepts on 1/kobs axes, from whose slopes and intercepts the values of the various rate constants can be evaluated.

The values of the rate constants of the elementary reaction (k) and the apparent rate constants ( kK1K2 ) at different temperatures were calculated from the temperature dependence of the rate constants at various [H+] and fixed substrate concentration using the least-squares method and are summarized in Table 2. Again, the formation constant (K1) was evaluated from the slopes and intercepts of those two plots and found to be 8.13± 0.5 dm3 mol-1 at 40 ºC.

The negative value of ΔS may confirm the compactness of the intermediates and are characterized by one- electron transfer mechanism of inner-sphere nature. Again, the positive values of ΔG obtained may confirm the non-spontaneity of the complexes formation in the rate-determining steps, as suggested by the cited proposed mechanisms. This means that the activated complexes could be more ordered and more compactness than that of the reactants which stabilized by a large solvation of the electron-transfer step. In addition, the non-dependency of the ionic strength of the rate constants may support this suggestion. But, this suggestion is not conclusive.

If the addition of Cl- ions to the reaction mixture affected the rate-constant of the oxidation reaction, it means that [IrCl6]2- may be aquationed 40 as in Eq.(10),

It was found that the addition of Cl- ions to the reaction mixture not affected the rate constants, therefore, this suggestions is neglected.

Again, addition of [IrCl6]3- product to the reaction mixture was found to has no effect on the reaction rates and, hence, this suggestion was excluded.

Furthermore, Moggi et al 43 showed that [IrCl6]2- is not appreciably photosensive in the visible region (433 or 495 nm) although it is high photosensitive in the UV- region (254 nm). Consequently photochemical induced reactions of [IrCl6]2- would not take place under the conditions used in the present work. Consequently, Eq. (6) can be considered as the sole rate-law expression for oxidation of MC by hexachloroiridiate (IV).

The kinetic parameters of rate constants (k) of the elementary reaction were calculated from the temperature dependence of the rate constants from Eyring and Arrhenius equations. These values were calculated by the method of least-squares and are summarized in Table 3.

Oxidation of alcoholic groups is generally occurring by the rupture of either C-H or C-OH bond through transfer of hydride, proton or hydrogen atom in the rate determining step 44, 45. Transfer of hydride ion is usually accompanied by a simultaneous two-electron transfer in a single step. This transfer occurs in the redox reactions involving multi-equivalent oxidants of labile nature such as permanganate and chromate ions. Whereas, the transfer of hydrogen atom or protons may take place in redox reactions involving either inert oxidants such as hexachloroiridate (IV), hexacyanoferrate (III) or labile Ce(IV) of one-equivalent nature which tend to proceed by one-electron transfer mechanism.

Unfortunately, the rate-law expression here not provided any information on the nature of reaction mechanism whether is of outer- or inner-sphere nature. Therefore, some information may be gained from examining the magnitude of the rate constant and/ or the magnitude of the entropy of activation. It has been reported 46, 47 that the entropies of activation, ΔS, are negative for the oxidation reactions which are proceeding by inner-sphere mechanism; while the ΔS for outer- sphere mechanisms tend to be of more positive values. Hence, two reaction mechanisms for formation of the intermediates may be considered, the first one being the transfer of electrons prior to the proton release. This mechanism corresponds to an outer-sphere type which proceeds by the formation of outer-ion sphere

followed by electron-transfer, then the release of protons as follows,

$\left[\text{S,}{\left[{\text{IrCl}}_{\text{6}}\right]}^{\text{2-}}\right]\stackrel{\text{slow}}{\to }\left[\text{S,}{\left[{\text{IrCl}}_{\text{6}}\right]}^{\text{3-}}\right]$

The second mechanism may involve the release of protons prior to the electron-transfer process which corresponds to the inner-sphere type as follows,

$\left[\text{S,}{\left[{\text{IrCl}}_{\text{6}}\right]}^{\text{2-}}\right]\stackrel{\text{slow}}{\to }\text{S+}{\left[{\text{IrCl}}_{\text{6}}\right]}^{\text{3-}}$

The former mechanism seems likely to be assigned for the present investigation from thermodynamic points of view. Therefore, the inner-sphere mechanism may be considered as a more suitable one for oxidation of MC by hexachloroiridate (IV). This suggestion has been supported by the negative value observed for ΔS (Table 3) along with the small values of the rate-constants of the elementary reaction (Table 2).

In view of the above kinetic interpretations and experimental observations a tentative reaction mechanism in good agreement with the kinetic results may be suggested by Scheme I.

Table 2. The rate constants (k) in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions. [IrCl6]2- = 2x10-4, [MC] = 0.2 and I = 0.1 mol dm-3.
Table 3. Activation parameters of the rate constant (k) in the oxidation of methyl cellulose by hexachloroiridate (IV) in aqueous solutions.
Scheme I.
Mechanism of oxidation of methyl cellulose by hexachloroiridate (IV).

### 5. Conclusions

The kinetics and mechanism of oxidation of methyl cellulose (MC) by hexacholroiridate (IV) in aqueous perchlorate solutions have been investigated spectrophotometrically. Evidence for formation of 1:1 intermediate complexes prior to the rate-determining step was revealed. The reaction kinetics seems to be of considerable complexity where one chloride ion from hexacholoiridate (IV) oxidant may act as a bridging ligand between the oxidant and the substrate into the formed intermediate complex. The kinetic observation indicated that this redox reaction proceeds by one-electron transfer mechanism of inner-sphere nature.

### References

[1]
McAllister JW, Lott JR, Schmidt PW, Sammler RL, Bates FS, Lodge TP (2015) Linear and nonlinear rheological behavior of febrile methylcellulose hydrogels. ACS Macro Lett 4: 538-542.
[2]
Lee JY, Kwak HW, Yun H, Kim YW, Lee KH (2016) Methyl cellulose nanofibrous mat for lipase immobilization via cross-linked enzyme aggregates. Macromolecular Research 24: 218-225.
[3]
Shaker AM (2001) Base catalyzed oxidation of carboxymethyl cellulose polymer by permanganate. I- Kinetics and mechanism of formation of manganate (VI) transient species complexes. J Colloid Interface Sci 233: 197-204.
[4]
Shaker AM (2001) . J Colloid Interface Sci 244: 254-261.
[5]
Hassan RM (2016) “  ”Edited by” Vijay K. Thakur & Manju K. Thakur “Pan Stanford Publishing, Pre Ltd. Publication City/Country Singapore, Singapore 411-454; Hassan RM, Dahy A, Ibrahim SM, Zaafarany IA, Fawzy A(2012) Oxidation of some macromolecules. Kinetics and mechanism of oxidation of methyl cellulose by permanganate ion in acid perchlorate solutions. Ind Eng Chem Res 51: 5424-5432.
[6]
Hassan RM, Ibrahim SM, Zaafarany IA, Fawzy A, Takagi HD (2011) Base-catalyzed oxidation: Kinetics and mechanism of hexacyanoferrate (III)oxidation of methyl cellulose polysaccharide in alkaline solutions. J Mol Cat A 344: 93-98.
[7]
Hassan RM, Fawzy A, Ahmed GA, Zaafarany IA, Asghar BH, Khairou KS (2009) Acid-catalyzed oxidation of some sulfated macromolecules. Kinetics and mechanism of permanganate oxidation of kappa-carrageenan polysaccharides in acid perchlorate solutions. J Mol Cat A 309: 95-102.
[8]
Zaafarany IA, Khairou KS, Hassan RM (2009) Acid-catalysis of chromic acid oxidation of kappa-carrageenan polysaccharide in aqueous perchlorate solutions. J Mol Cat 302: 112-118.
[9]
Hassan RM, Ibrahim SM, Dahy A, Zaafarany IA, Tirkistani F, Takagi HD (2013) Kinetics and mechanism of oxidation of chondroitin-4-sulfate polysaccharide by chromic acid in aqueous perchlorate solutions. Carbohyd Poly 92: 2321-2326.
[10]
Hassan RM (1993) Alginate polyelectrolyte ionotropic gels. XVIII. Oxidation of alginate polysaccharide by potassium permanganate in alkaline solutions. Kinetics of decomposition of the intermediate complex. J Poly Sci 5A, 31: 1147-1151.
[11]
Hassan RM (1993) New coordination polymers. III. Oxidation of poly (vinyl alcohol) by permanganate ion in alkaline solutions. Kinetics and mechanism of formation of an intermediate complex with a spectrophotometric detection of manganate(VI) transient species. Poly Inter 30: 5-9.
[12]
Abdel-Hamid MI, Khairou KS, Hassan RM (2003) Kinetics and mechanism of permanganate oxidation of pectin in acid perchlorate media. Eur Poly 39: 381-387.
[13]
Ahmed GA, Khairou KS, Hassan RM (2003) Kinetics and mechanism of oxidation of chitosan polysaccharide by permanganate ion in aqueous perchlorate solutions. J Chem Res 182-183.
[14]
Hassan RM, Fawzy A, Ahmed GA, Zaafarany IA, Asghar BH, Takagi HD, Ikeda Y (2011) Kinetics and mechanism of permanganate oxidation of iota- and lambda- carrageenan polysaccharides as sulfated carbohydrates in acid perchlorate solutions. Carbohydrate Research 346: 2260-2267.
[15]
Hassan RM, Zaafarany IA, Takagi HD (2013) Oxidation of some water-soluble anionic polyelectrolytes: Oxidation of carboxymethyl cellulose polysaccharide by hexachloroiridate(IV) in aqueous perchlorate solutions. A Kinetic and mechanistic approach to electron transfer process. Indust & Eng Chem Res 52 : 1531-1537; Ahmed G A, Hassan RM (2001) Kinetics and mechanism of reduction of hexachloroiridate(IV) by kojic acid in aqueous perchlorate solutions. Bull Polish Acad Sci 49: 235-244.
[16]
El-Korshy MA, Hassan R M (1998)Kinetics and mechanism of oxidation of ADA by hexachloroiridate (IV) in aqueous perchloric acid. Bull Polish Acad Sci 46: 147-155.
[17]
Hassan RM, Zaafarany IA, Takagi HD, Ikeda Y (2013) Kinetics and mechanism of hexachloroiridate(IV) oxidation of tellurium(IV) in aqueous solutions. New J Chem 37: 2700-2707; Hassan RM (1991) Kinetics of reaction of uranium (IV) and hexachloroiridate(IV) in acid perchlorate solutions. J Chem Soc Dalton Trans 3003-3008.
[18]
Hassan RM (1992) Kinetics and mechanism of hexachloroiridate(IV) oxidation of arsenic(III) in acidic perchlorate solutions. Collect Czech Chem Commun 57: 1451.
[19]
Poulson I A, Garner CS (1962) A thermodynamic and kinetic study of hexachloro and aquopentachloro complexes of iridium(III) in aqueous solutions. J Am Chem Soc 84: 2032-2037.
[20]
Thornley RNF, Sykes SA (1970) Kinetics of the one-equivalent reactions of vanadium(II), vanadium(III), and iron(II) with hexachloroiridate(IV). J Chem Soc A 1036-1037.
[21]
Morris DFC, Ritter TJ (1979) Oxidation of hydrazine by halogeno complexes of iridium (IV) in acidic perchloric solutions. J Chem Soc Dalton Trans 216-219.
[22]
Malik MA, Ilyas M, Khan Z (2009) Kinetics of permanganate oxidation of synthetic macromolecule poly(vinyl alcohol). Ind J Chem 48 A : 189-193.
[23]
Khairou KS, Hassan RM, Shaker MA(2002) Novel synthesis of diketocarboxymethyl cellulose as biopolymer precursors. J Appl Poly Sci 85: 1019-1023.
[24]
Hassan RM, Abd-Alla MA(1992) New coordination polymers. I. Novel synthesis of poly (vinyl alcohol) and characterization as chelating agent. J Mater Sci 2: 609-611.
[25]
Pretch E, Clerc T, Seibl J, Simon W, Tables of Spectral Data for Structure Determination of Organic Compounds, Springer-Verlag, Berlin, Heidelberg, New York, Tokyo, 1983 (Translation).
[26]
Hassan RM(1991) Kinetics and mechanism of oxidation of DL-α-alanine by permanganate ion in acid perchlorate media. Can J Chem 69: 2018-2023.
[27]
Sen Gubta KK, Chatterjee U, Sen PK(1978) Kinetics and mechanism of oxidation of glyoxylate ion by Ir (IV). Ind J Chem 16B: 767-770.
[28]
Sen Gubta KK, Bhattacharjee N(2002) Oxidative behavior and relative reactivities of some unsaturated compounds towards hexachloroiridate(IV) in perchloric acid medium. Int J Chem Kinet 34: 411-417.
[29]
Sen Gubta KK, Uma Chatterjee (1981) Kinetics of oxidation of methanol, ethanol and isopropanol by hexachloroiridate(IV). J Inorg Nucl Chem 43: 2491-2497.
[30]
Ogino H, Bailar Jr. JC (1978) Stereochemistry of complex inorganic compounds. 36. Ammoniation reactions of some optically active 1,2-dihalobis(ethylenediamine) complexes of rhodium(III) and iridium(III) ions. Inorg Chem 17: 1118-1124.
[31]
Po HN, Lo CF, Jones N, Lee RW (1980) The oxidation of 2-thiopyrimidine and 2-thiouracil by Ir(IV) complexes. Inorg Chim Acta 46: 185-189.
[32]
Kottapalli KK, Adari KK, Vani P, Govindan SK (2005) Mechanism of Oxidation of L-Cysteine by Hexachloroiridate(IV) - A Kinetic Study. Trans Metal Chemistry 30 : 773-777; Ayoko GA, Lyun JF, Ekubo T (1992) Kinetics of the reduction of hexachloroiridate (IV) by L-methionine in aqueous-solutions.‏ Ind J Chem 31 A: 975-980.
[33]
Sykes AG, Thornley RNF(1970) Identification of inner- and outer-sphere paths in the reaction of chromium(II) with hexachloroiridate(IV), and the kinetics of the decomposition of the binuclear intermediate. J Chem Soc A 232-238.
[34]
San Gupta K K, Deg S, Sen Gupta S, Banerjee A (1984) Kinetics and mechanism of the oxidation of aromatic aldehydes by hexachloroiridate(IV). J Org Chem 49: 5054-5057.
[35]
Scurlock RD, Gilbert DD, Dekorte IM (1985) Oxidation of ascorbic acid by aquopentabromoiridate(IV): an assessment of the effect of aquation on rates of oxidation by bromo- and chloroiridium(IV) complexes. Inorg Chem 24: 2393-2397.
[36]
Pirk JP (1977) Mechanism of the reversible oxidation of vanadium(IV). Inorg Chem 16: 1381-1383.
[37]
Newton TW (1986) The kinetics of oxidation-reduction reactions - an alternative derivation of Marcus’ cross relation. J Chem Educ 45: 571-576.
[38]
Gordon BJ, Williams LL, Sutin N (1961) The kinetics of oxidation of iron(II) ions and coordination complexes. J Am Chem Soc 83: 2061-2064.
[39]
Wharton RK, Ojo JF, Sykes AG (1975) Mechanism of the oxidation of the molybdenum(V)-ethylenediaminetetra-acetato dimer by hexachloroiridate(IV) and tris(1,10-phenanthroline)iron(III). J Am Chem Soc Dalton Trans 1526-1530.
[40]
Adedinsewo CO, Adgite A (1979) Linear free energy relations in redox reactions. Oxidation of iodide ion by poly(pyridine)-iron(III) complexes and hexachloroiridate(IV) ion. Inorg Chem 18: 3597-3601.
[41]
Hurwitz P, Kustin K (1964) Kinetics of fast electron-transfer reactions. Inorg Chem 3: 823-826.
[42]
Sen-Gupta KK, Chatterjee U (1981) Kinetics of oxidation of methanol, ethanol and isopropanol by hexachloroiridate(IV). J Inorg Nucl Chem 43: 2491-2497.
[43]
Moggi L, Varani G, Manfrin MF, Balazani V (1970) Photochemical-reactions of hexachloroiridate(IV) ion. Inorg Chem Acta 4: 335-341.
[44]
AL-Ajlouni, Bakac A, Esponson JH (1994) Hydride abstraction from 1,2-diols by the pentaaqua(oxo)chromium(IV) ion. Inorg Chem 33: 1011-1014.
[45]
Chimatdar SA, Nandibewoor ST, Sambrani MI, Raju JR (1987) Chromium(III)-catalysed cerium(IV) oxidation of arsenic(III) in aqueous sulphuric acid. J Chem Soc Dalton Trans 573.
[46]
Sutin N (1968) Free energies, barriers and reactivity patterns in electron transfer reactions. Acc Chem Res1: 225-231.
[47]
Hassan RM(1992) A review on oxidation of uranium (IV) by polyvalent metal ions. A linear free-energy correlation. J Coord Chem 27: 255-266.