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KINETIC STUDIES IN ORGANIZED MEDIA

ABSTRACT

THESIS SUBMITTED FOR THE AWARD OF THE DEGREE OF

in

CHEMISTRY

By

S. M. SHAKEEL IQUBAL

DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY

ALIGARH (INDIA)

2006

ABSTRACT

Chemists have long recognized the important role the reaction media

plays in controlling rates, product distribution and stereochemistry. Much effort

has been directed towards the use of organized media to modify reactivity, as

compared to that in isotropic liquids. Among the many ordered or constrained

systems utilized to organize the reactants, the notable ones are micelles,

microemulsions, liquid crystals, inclusion complexes, monolayers, and solid

phases such as adsorbed surfaces and crystals.

Chemical reactivity in colloidal self-assemblies {e.g., micelles,

microemulsion droplets, and vesicles) has got importance owing to similarities

in action with the enzymatic reactions.' Although the analogy between micelle-

catalyzed reactions and enzyme-catalyzed reactions is far from perfect, there

are important similarities.^ The structures of both micelles and enzymes are

similar in that they have hydrophobic cores with polar groups on their surfaces.

Both catalytic micelles and enzymes bind substrates in a non-covalent manner.

The kinetics of micellar catalysis resemble that of enzymatic catalysis in that

the micelle may be saturated by the substrate; and, conversely, the substrate

may be saturated by the micelle.

The mechanism of a chemical reaction cannot be fully described without

the determination of its rate. The kinetic study of a wide range of chemical

processes is seen to be of essential importance, not only in pure research but

increasingly in industrial research, development and, in some instances, in

quality control and analysis as well. Kinetic methods have become an essential

technique in piiotochemistry, enzyme chemistry, study of chemical catalysis,

etc.

In this thesis only one specific type of organized system is considered,

i.e., normal micelles. Micelles are apparently isotropic, although their intimate

nature is microheterogeneous; they fonn spontaneously and they are

thermodynamically stable.

It is well known that permanganate ion has been extensively used and

studied as an oxidizing agent (oxidant). A large number of reports are available

on the kinetics and mechanisms of organic compounds by potassium

permanganate. Soluble forms of colloidal manganese dioxide are formed as

products in many permanganate reactions. They play an important role in

autocatalytic''"'* and oscillating ''"' reactions. They are also involved as

detectable intermediates in some permanganate reactions carried out in acidic

solutions.

A few years ago, Perez-Benito et fl/.'"* reported the preparation of a

stable colloidal Mn02 in aqueous solution. Later, this water soluble form was

used as an oxidizing agent for organic acids. The kinetic and mechanistic

aspects of the interaction of different types of reductants with colloidal

manganese dioxide have been investigated by some workers under different

conditions. Detailed micellar effect on the oxidation of organic acids by water

soluble colloidal manganese dioxide has, however, not been studied so far

except a few reports by Tuncay et al. ' and Kabir-ud-Din et al. ' Therefore, the

present study was designed to extend knowledge on the behavior of colloidal

Mn02 as an oxidant.

The work described in the thesis entitled " Kinetic Studies in

Organized Media" deals with systematic kinetic studies of the reactions of

colloidal manganese dioxide with some organic acids, i.e., glycolic, lactic,

mandelic, malic, tartaric and citric acids, both in aqueous as well as in miceliar

media.

The lay out of the thesis is as follows: (i) Chapter 1 - General

Introduction; (ii) Chapter 2 - Experimental; and (iii) Chapter 3 - Results

and Discussion.

Chapter-1 comprises of an introduction of organized media, surfactants

and miceliar organization, and oxidation by permanganate and colloidal

manganese dioxide. The chapter ends with the statement of the problem which

suggests the importance of this study.

Chapter-2 contains experimental details. The materials used, their

structure and formulas, sources and purities are given along with the method of

preparation and characterization of water- soluble colloidal Mn02. Kinetic

measurement details are also provided in this chapter.

Chapter-3 — Results and Discussion, as the name implies, covers all

the results obtained with their discussion in two parts, namely: (A) Reactions in

Absence of Surfactants; and (B) Reactions in Presence of Surfactants.

The kinetic studies of oxidation of organic acids (glycolic, lactic,

mandelic, malic, tartaric and citric) were performed spectrophotometrically by

following disappearance of the oxidant (colloidal Mn02) at 390 nm in the

absence and presence of surfactant micelles at several concentrations of the

reactants, hydrogen ion, manganese(II), gum arable, temperature, and

surfactants. The log(/l39o) versus time plots show that the reactions of the acids

and colloidal Mn02 proceed in two stages and that the first stage (noncatalytic)

is relatively slower than the second (autocatalytic). The kinetic curves show

inflection and the A:obs values (for both the stages) were estimated from such

curves.

On the basis of order-dependence on different variables, the empirical

rate law is written as

dlMnOj] = {k' + k" [Uy} [Mn02] [reductant]'" (1)

d/

and a mechanism satisfying the observed requirements is proposed for the

noncatalytic reaction path as follows:

(i)pH - independent pathway

^1 K,i^ ^1 (Mn02)n +H0—C—COOH I ^ (Mn02)ir(0H-C-C00H) (2)

R2 R2 (B) (CI)

C I — ^ U - (Mn02)n.i+ Mn(II) + ( p O + COOH (3)

R2 (PI)

COOH + (Mn02)n ^ ^ ^ ^ (Mn02)n.i + CO2 + Mn(III) (4)

R, Rj

HO—(j:—COOH + Mn(III) — ^ HO— C—COO + Mn(II) (5)

R2 R2 (radical)

fast ' radical >• C = 0 + CO2 (6)

R,

(ii) pH - dependent pathway

^ ' . ad2 ? ' (Mn02l;r(0H-(j:-C00H) + H^ ^ = : ^ ir_^Mn02)r(0H-C-C00H) (7)

R2 R2 (CI) (C2)

C2 ^ (Mn02)n-i+ Mn(II)+ COOH + PI (8)

(reactions (4) to (6) then follow).

SCHEME 1

(Ri = R2 = -H glycolic acid; R, = -H, R2 = -CH3 lactic acid; Ri =

-H, R2 = -CftHs mandelic acid; Ri = -H, R2 = -CH2COOH malic

acid; R, = -H, R2 = -CH(OH)COOH tartaric acid; R, = R2 = -CH2COOH

citric acid)

(Mn02)„ stands for colloidal manganese dioxide. Eq. (2) represents

adsorption of the carboxylic acid on the surface of the colloidal Mn02. After

adsorption, species C1 undergoes electron transfer leading to the formation of

product (PI), radical COOH and Mn(II) in a rate determining step. In the

subsequent fast step (Eq. (4)), Mn02 further reacts with COOH and gives

Mn(III) and CO2 as reaction products. The observed dependence of the order

on [reductant] being different (see Eq. (I)) indicates that the reactivity of the

reductant (in the adsorbed state) probably plays a role in the transfer of

electrons to colloidal Mn02.

Freundlich adsorption isothemi has been used to explain the adsorption

of reductants on colloidal Mn02.

The reactions proceed in two stages: the second being faster than the

first. The reactions are thus typically autocatalytic, i.e., catalyzed by one of the

products. In many permanganate - organic substrate reactions Mn(II) has been

the active autocatalyst. As Mn(II) is a product in the present reactions too,

experiments were performed with initially added Mn(II) in order to confirm its

role. The results show sigmoid dependence on [Mn(II)], thereby establishing

Mn(II) as the autocatalyst."''*

In the presence of Mn(II), Scheme 1 is, therefore, modified as Scheme 2:

(Mn02)n

ypath l l path I \ ^

(Mn02)n-i + Mn(II) Mn(IIl) +(Mn02)n-i +

other products

SCHEME 2

Both Mn(II) and reductants (B) compete to react with the colloidal

Mn02 and the values of rate constants are the sum of the paths I and II

(SCHEME 2). Thus, the exact dependence of /robsi/ obsi on [Mn(II)] cannot be

predicted.

Cationic surfactant CTAB caused flocculation of the oppositely charged

colloidal Mn02 and could not be used whereas anionic SDS micellar systems

were found to be ineffective on the reaction rates. However, the reactions were

found to be accelerated by nonionic micelles of Triton X-100 which have been

explained in terms of the mathematical model proposed by Tuncay et al)^

Effect of variables [oxidant], [reductant], [HCIO4], and temperature

were also studied in presence of TX-lOO and it was found that the dependence

of rate constants on all was of the same pattern as in case of aqueous medium.

These observations establish that the same mechanisms are being followed in

both the media.

As far as the role of TX-100 is concerned, hydrogen bonding between

the non-ionic TX-lOO (I) and the reactants plays an important role. The organic

acids contain no hydrophobic character (except mandelic acid) and have OH

CH3 CH3

CH3-C-CH2-C-<^fOCH2CH2);rOH

CH3 CH3

TX-lOO (I)

and COOH groups. Hydrogen bonding may occur between these groups of the

acids (reductants) and the ether-oxygen of the polyoxyethylene chains of

TX-lOO. Due to the presence of a number of donor groups in one TX-lOO

molecule, multiple H-bonding may take place and the number of bound

reductant molecules increase. In this surfactant, the lengths of the hydrophobic

and hydrophilic parts are comparable and have significant amount of water in

the outer shell. The hydrogen bonding between Mn02 sols and hydrophilic part

(polar ethylene oxide) of the TX-lOO cannot be ruled out either. Therefore, the

associated MnO: and reductants with TX-lOO (through hydrogen bonding)

seems responsible of facilitating the reaction. This surfactant thus helps in

bringing the reactants closer, which may orient in a manner suitable for the

redox reactions followed by rearrangement of TX-lOO molecules.

10

References;

(1) J.H. Fendler, "Membrane Mimetic Chemistry", Wiley, New York,

1982.

(2) (a) E.J. Fendler and J.H. Fendler, Adv. Phys. Org. Chem., 8, 271

(1970); (b) J. H. Fendler and E.J. Fendler, ''Catalysis in Micellar

andMacromolecular Systems", Academic Press, New York, 1975.

(3) J. F. Perez-Benito, F. Mata-Perez and E. Brillas, Can. J. Chem.,

65,2329(1987).

(4) J. F. Perez-Benito and C. Arias, Int. J. Chem. Kinet., 23, 717

(1991).

(5) A. Nagy and L. Treindl, Nature, 320, 344 (1986).

(6) M. Orban and 1. R. Epstein, J. Am. Chem., Soc, 111, 8543 (1989).

(7) K. Fujiwara, T. Kashima, H. Tsubota, Y. Toyoshima, M. Aihra

and M. Kiboku, Chem. Lett., 1385 (1990).

(8) M. Orban and I. R. Epstein, J. Am. Chem., Soc, 112, 1812 (1990).

(9) M. Orban, I. Lengyel and I. R. Epstein, J. Am. Chem., Soc, 113,

1978(1991).

(10) C. J. Doona and F. W. Scheider, J. Am. Chem., Soc, 115, 9683

(1993).

(11) A. Nagy, P. G. Sorensen and F. Hynne, Z Phys. Chem. (Munich),

189, 131 (1995).

11

(12) M. Melichercik, M. Mrakarova, A. Nagy, A. Olexova and L.

Treindl, Collect. Czech. Chem. Commun., 60, 781 (1995).

(13) A. Olexova, M. Melicherik and L. Treindle, Chem. Phys. Lett.,

268, 505 (1997).

(14) J. F. Perez-Benito, E. Brillas and R. Pouplana, Inorg. Chem.,

28, 390 (1989).

(15) M. Tuncay, N. Yuce, B. Arlkan and S. Gokturk, Colloids

Surf. A: Physicochem. Eng. Aspects, 149, 279 (1999).

(16) Kabir-ud-Din, W. Fatma and Z. Khan, Colloids Surf. A:

Physicochem. Eng. Aspects, 234, 159 (2004).

(17) J. W. Ladbury and C. F. Cullis, Chem. Rev., 58, 403 (1958).

(18) Z. Khan, Raju, M. Akram and Kabir-ud-Din, Int. J. Chem. Kinet.,

36, 359 (2004).

KINETIC STUDIES IN ORGANIZED MEDIA

THESIS SUBMITTED FOR THE AWARD OF THE DEGREE OF

m

CHEMISTRY

By

S. M. SHAKEEL IQUBAL

DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY

ALIGARH (INDIA)

2006

^;-; A pad f / a > v '^i V ,

m % ^

T7018

HO

PROF. KABIR-UD-DIN CHAIRMAN DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY ALIGARH-202002 (U. P.) INDIA. E-mail: kabir7@rediffmail.com

9 EXT. (0571) -2703515 Int 3353,3351

Certificate

Dated:.3\\9S..\.(i C

This is to certify that the thesis entitled "Kinetic Studies in

Organized Media" is the original work carried out by Mr. S. M. Shakeel

Iqubal under my supervision and is suitable for submission for the award

of Ph.D. degree in Chemistry.

(Prof. Kabir-ud-Din)

1? ACKXoWC(E(DgT:M^ms

When every hope is gone, and helpless flee: help comes from I know not

where, this is the essence of GOD. I am thankful to Allah (SWT), sustainer of

the world, who bestowed upon me the opportunity to accomplish this work,

which would have not seen the light of the day otherwise. The Holy verses and

Prophet's sayings have always infused in me a zest to explore the unknown. I

pray to Allah to advance me in knowledge and that my endeavors guide me

towards the right path.

I am fortunate to have worked under Prof. Kabir-ud-Din, Chairman,

Department of Chemistry, AMU, Aligarh. I am indebted for his invaluable

supervision, encouragement and academic help. His pertinent ability to probe

into the research problems has made this study a pleasant one.

I also place, on record my gratitude to Dr. Zaheer Khan, Reader,

Department of Chemistry, Jamia Millia Islamia, New Delhi, who supported and

encouraged this effort in every way possible at the expense of his own

important preoccupations.

I wish to express my sincere thanks to Prof Imam Hussain and

Prof S. N. Verma, former Heads of the Department of Chemistry, G. D.

College, Begusarai, Bihar.

I am extremely thankful to Dr. Sanjeev Kumar, Dr. M. Z. A. Rafiquee,

Dr. M. Akram, Dr. Andleeb Z. Naqvi, Dr. Ziya Ahmad Khan and Dr. Damyanti

Sharma.

My loving thanks and best wishes for my lab colleagues Dr. Daksha

Shamia, Mrs. Waseefa Fatma, Ms, Nahid Parveen, Mr. Mohd. Sajid Ali, Mrs.

Umme Salma, Mrs. Deepti Sharma, Ms. Nuzhat Gull, Mr. Md. Sayem Alam,

Mr. Tanweer Ahmad, Ms. Neeiam H. Zaidi, Mr. Mohd. Altaf, Mr. Naved

Azam, Mrs. Suraiya and Mr. Mohd. Dabi Ali Al Ahmadi.

I especially thank uncle Akram and Khursheed Bhai who always stood

by me whenever I needed them.

Finally, I don't find suitable words to express my deepest thank and

gratitude to my grandparents, parents, brothers, sisters and brothers-in-law for

their love, untiring assistance, constant support and encouragement throughout

the tenure of this work.

I am also thankful to Mr. Salim-uddin (cartographer) and Mr. Zaheer

Ahmad (Limra Computers) for their work and co-operation.

S. M. Shakeel Iqubal

LISTO'F (F0(BLICATI09{S

(1) Kinetics of the Reduction of Colloidal Mn02 by Citric Acid in the

Absence and Presence of Ionic and Non-ionic Surfactants.

Kabir-ud-Din, S. M. Shakeel Iqubal and Zaheer Khan.

Inorg. React. Mech., 5, 151-166 (2005).

(2) Reduction of Soluble Colloidal Mn02 by DL-Malic Acid in the Absence

and Presence of Non-ionic Triton X-100.

Kabir-ud-Din, S. M. Shakeel Iqubal and Zaheer Khan.

Colloid Polym. Sci., 283, 504-511 (2005).

(3) Effect of Ionic and Non-ionic Surfactants on the Reduction of Water

Soluble Colloidal Mn02 by Glycolic Acid.

Kabir-ud-Din, S. M. Shakeel Iqubal and Zaheer Khan.

Colloid Polym. Sci., 284, 276-283 (2005).

(4) Oxidation of DL-Tartaric Acid by Water Soluble Colloidal Mn02 in the

Absence and Presence of Surfactants.

Kabir-ud-Din, S. M. Shakeel Iqubal and Zaheer Khan.

Indian J. Chem., 44A, 2455-2461 (2005).

(1) Kinetics of Oxidation of DL-Malic Acid by Colloidal Mn02.

Kabir-ud-Din and S. M. Shakeel Iqubal.

91st Indian Science Congress, Chandigarh, Jan. 3-7, 2004.

(2) Effect of Nonionic Triton X-100 on the Oxidation of Malic Acid by

Water Soluble Colloidal MnOj.

Kabir-ud-Din and S. M. Shakeel Iqubal.

15th International Symposium on Surfactants in Solution (SIS

2004), Fortaleza (Brazil), June 6-11, 2004.

CHAPTER 1

CHAPTER 2

CONTENTS

GENERAL INTRODUCTION

A. Organized Media, Surfactants and

Micellar Organization

B. Oxidation by Permanganate and

Colloidal Manganese Dioxide

C. Statement of the Problem

References

EXPERIMENTAL

Materials

Preparation of Water-Soluble Colloidal Mn02

Characterization of Colloidal Mn02

Kinetic Measurements

Detection ofC02

Free Radical Detection

References

CHAPTER 3 RESULTS AND DISCUSSION

A. Reactions in Absence of Surfactants

B. Reactions in Presence of Surfactants

References

Page

1-27

14

19

20

28-40

29

29

33

37

38

39

40

41-139

42

93

137

CHAPTER 1

GENERAL INTRODUCTION

A. Organized Media, Surfactants and Micellar Organization

Chemists have long recognized the important role the reaction media

plays in controlling rates, product distribution and stereochemistry. Much effort

has been directed toward the use of organized media to modify reactivity, as

compared to that in isotropic liquids. A major goal of such studies is to utilize the

order of the medium so as to increase the rate and selectivity of the chemical

process involved in much the same way that enzymes modify the reactivity of

the substrates to which they are bound. Among the many ordered or constrained

systems utilized to organize the reactants, the notable ones are micelles,

microemulsions, liquid crystals, inclusion complexes, monolayers and solid

phases such as adsorbed surfaces and crystals. With these media partial

constraints are imposed upon the reactants, limiting the overall number of

possible transition states which can be formed, subsequently decreasing the

number of products formed. Judicious selection of a given organized system for

a given application requires a sufficient understanding and properties of the

organized media themselves and those of the substrate interactions therein.

Due to their widespread uses in many industrial applications there has

been an increasing interest in the surfactant organized assemblies both from

academic and applied point of view. Self-organized systems composed of

amphiphilic molecules have particular features that make them attractive, not

only as relates to chemical reactivity aspects but also for a large variety of

applications. For instance, surfactants have been used for extraction of metal

ions.' The application of surfactant-based systems as drug delivery vehicles '"' is

a growing research area that may develop further in the coming years. It is quite

interesting that cationic amphiphiles are now widely used as an effective tool in

delivering DNA into cells'*' even mammalian cells.^''° Bilayer-forming synthetic

surfactants have been extensively used as membrane mimetic models, and some

synthetic amphiphiles such as dihexadecyl phosphate or

dioctadecyldimethylammonium salts have found many different uses in strategic

applied areas."

Aqueous association colloids as reaction media offer alternatives to the

use of organic solvents, and there is considerable interest in their use in water as

a reaction medium; they are attractive candidates in "green" chemistry. " In

fact, they are expected to be nontoxic and nonhazardous, they enhance reaction

rates, reactions can usually be carried out under mild conditions, and in favorable

cases surfactants can be separated and reused. Moreover, jeagent organization in

and different microenvironments provided by aqueous micelles open the

possibility of controlling product formation, providing chemo-, regio-, or

stereoselectivity.'^''^ Alignment of reactants at the micellar surface was at the

basis of the observed regioselectivity control in Diels-Alder reactions in which

both reactants were surface active. The Diels- Alder reaction is an important tool

in organic synthesis, and there are a number of studies of this reaction in aqueous

surfactant systems.'^ The use of micelles to induce enantioselectivity is currently

of high interest. High values of enantioselectivity were obtained by Brosch and

Kirmse in the nitrous acid deamination of amines.

The preceding facts show the importance of surfactant organized media in

influencing chemical reactions (rates and equilibria as well). Therefore, a

fundamental understanding of the physical chemistry of such organized

assemblies, their unusual properties and phase behavior is essential.

A surfactant {surface-active ageni^^) is a substance that, when present at

low concentration in a system, has the property of adsorbing onto the surfaces or

interfaces of the system and of altering to a marked degree the surface or

interfacial free energies of those surfaces (or interfaces). On the basis of the

chemical structures of their molecules and the surface active moiety formed

through dissociation in polar aqueous media, surfactants can be classified as

cationic, anionic, nonionic, or zwitterionic.

Cationic: The cation of the compound is the surface active species. The most

prevalent cationic surfactants are based upon quaternary nitrogen.

Alkylammonium halides and tetra-alkyiammonium halides are the most

numerous in this class. Pyridine and related species such as quinoline, iso-

quinoline, pyra2:ine, and their derivatives form the basis for a wide class of aryl-

based quaternary surfactants. Although less numerous, phosphorous can also be

quatemarized with alkyl groups to provide alkylphosphonium surfactants.

Examples:

cetyltrimethylammonium chloride CH3(CH2)i5 >r(CH3)3Cr

dodecylamine hydrochloride CH3(CH2)iiN^H3Cr

dodecylpyridinium chloride {QNC,2H25c r

Anionic: Alkali alkanoates (or soaps) are the most well known anionic

surfactants, the ionized carboxyl group providing the anionic charge.

Sodium dodecylbenzene sulfonate (SDBS) has effectively replaced

soaps. Alkyl sulfates, alkyl ether sulfates, alkyl sulfonates, secondary alkyl

sulfonates, aryl sulfonates, methylester sulfonates, and sulfonates of

alkylsuccinates are other important classes of anionic surfactants. The fatty

acids and these sulfo compounds include the three most important anionic

groups: carboxylate (-CO2"), sulfate (-OSO3"), and sulfonate (-SO3"). On the

basis of basicity and phase data ^ their hydrophilicity ranking is given as:

-C02~ » - SO3" > - 0S03~

Phosphate mono- (-P(0H)02~) and dianions (-POs^') are also important

hydrophilic groups.

Examples:

potassium laurate CH3(CH2)ioCOO~K^

sodium dodecyl sulfate CH3(CH2)i iOS03~ Na

sodium dodecylbenzene sulfonate CH3(CH2) , O C H 2 H ( ^ > - SOI Na

Nonioiiic: The water-soluble moiety of this type can contain hydroxyl groups

or a polyoxyethylene chain. Many nonionic surfactants are structurally

analogous to anionic and cationic surfactants, except that the headgroup is

uncharged. Most prevalent among the headgroups of nonionics are oligomers

of ethylene oxide. Simple saccharides such as glucose and sucrose are also

common as headgroups for nonionic surfactants. The most widely studied class

of alkyl ethylene oxide surfactants, also called alcohol ethoxylates, is

represented as C Ey where x is the length of the alkyl chain and y is the number

of ethylene oxide units in the headgroup. A related class is that of the alkyl

phenol ethoxylates. Triton X-100 (TX-lOO) is perhaps the best known member

of this class.

Block copolymeric nonionic surfactants have become an important class

of dispersants known commonly as polyoxamers and Pluronics, Such block

copolymers are often denoted as AB (diblock) or ABA (triblock), where A

denotes a hydrophilic ("headgroup") block such as poly(ethylene oxide) (EO),

and B denotes a hydrophobic block, such as poly(propylene oxide) (PO) or

polystyrene (PS). Commercially available EO/PO copolymeric surfactants

generally contain a mixture of homologues of various chain lengths. Ethylene

oxide containing surfactants are generally considered to hydrate to the extent of

about three water molecules per ethylene oxide group.

Alkanol amides such as ethanolamides and diethanolamides,

alkylamides, amine ethoxylates, amine oxides (at neutral and alkaline pH), and

poiyamines are the primary nitrogen-based nonionic surfactant types.

Examples:

CH3(CH2),,N(CH3)2 N,N-dimethyldodecyiamine oxide 1

O

polyoxyethylene monohexadecyl ether CH3(CH2)i50(CH2CH20)2iH

polyethylene glycol (/) - octylphenyl ether, TX-lOO

(CH3)3CCH2C(CH2)2-<(O^~0(^"2CH20)9.5H

Zwitterlonic: This type of surfactants can behave as either cationic, anionic or

nonionic species depending on the pH of the solution. True zwitterions such as

a-amino acids, can become ionized by intramolecular proton transfer:

NH2CH(R)C02H o N%CH(R)C02"

Considerations of phase data and basicity^^ suggest that the relative

hydrophilicities of ammonio zwitterionics decrease in the following order:

ammonio-C02~ » ammonio-SOs" > ammonio-OSOs".

The betaines are a very important class of zwitterionic surfactants and

include alkylbetaines, amidoalkylbetaines, and heterocyclic betaines.

Examples:

N-dodccyl-N: N-dimethyl betaine Ci2H25N'"(CH3)2CH2COO"

3-(dimethyldodecylammonio) propane-1-sulfonate

CH3(CH2),,N (CH3)2CH2-CH2-CH2-S03"

2,3-dimethyl-3-dodecy]-],2,4-triazo]ium-5-thio]ate

CH, N=N

N I CH2(CH2)ioCH3

Critical Micelle Concentration (cmc):

In dilute aqueous solution, the ionic surfactants behave like strong

electrolytes, while nonionic ones often resemble simple organic molecules. At

higher amphiphile concentrations, generally greater than 10"'' mol dm'\ a

marked deviation from "ideal behavior" occurs in dilute solution. This

deviation is larger than that exhibited by strong electrolytes. A generalized

diagram for such variations in physical properties as a ftinction of the surfactant

concentration is given in Fig. 1.1. Some of the physical properties found to

exhibit this type of behavior are related to the interfacial tension, the electrical

conductivity, the electromotive force, the transport properties such as the

viscosity, optical and spectroscopic properties of the solution. The well-

defined, but not abrupt, changes in the physical properties are attributable to the

cmc

O

Q.

resono'

Surf octant Concentrat ion

Fig. 1.1: Variation ofpiiysical properties with surfactant concentration.

10

association of the amphiphiles forming aggregates or micelles. The

concentration at which the micelles appear corresponds to the change in the

slope of the plot and is known as critical micelle concentration or cmc. This

change occurs over a narrow concentration range rather than at a precise point

and the magnitude of this range depends on the physical property being

measured. In their excellent compilations Mukerjee and Mysels and Shinoda

et al}* have mentioned techniques for determining cmc.

The primary driving force for micelle formation, in general terms, is the

tendency of the hydrocarbon part of the monomer to associate with itself rather

than to remain in close proximity with water. Micelle formation is thus an

outstanding example of what is increasingly being known as "hydrophobic

bonding". Because hydrophobic bonding has a strong additive character,

substantial free energy changes are involved in micelle formation. Interionic

interactions in ionic association colloids and the substantial role played by the

hydrophilic headgroups give additional complexity to the problem of micelle

formation.

The value of cmc is dependent upon a large number of parameters like

total carbon chain-length, additional polar groups, C = C double bonds, chain

branching of surfactants and various types of additives; polar and nonpolar,

electrolytes, temperature, and pressure.

11

Micellar Structure and Micellar Organization:

The structure of a normal micelle just above cmc can be considered as

roughly spherical." '*' ^ When the hydrophobic portion of the surfactant is a

hydrocarbon chain, the micelle will consist of a liquid-like hydrocarbon core.

The radius of this core is roughly equal to the length of fully extended

hydrocarbon chain (~ 12-30 A). The polar headgroups and bound water are

regularly arranged at the micellar surface which is rough. Menger has

proposed that water can penetrate inside the micelle upto a certain level, ' the

idea gets support from fluorescence and 'H NMR measurements. Partial molar

volume determinations indicate that the alkyl chains in the core are more

expanded than those in the normal liquid state."'"

The nonionic micelles arrest water molecules at the palisade layer by

hydrogen bonding of water with the polyethylene oxide groups.'" Water may

remain trapped in this region.

In ionic micelles the surface potentials are high '"'' and a significant

fraction of the counterions (60-90%)'' are located in a compact region known

as 'Stern layer'^' which extends from the core to within a few angstroms of the

shear surface of the micelle. The core and the Stern layer form the 'kinetic

micelle'. Most of the remaining counterions are, however, located outside the

shear surface in the region called 'Gouy-Chapman electrical double-layer'. The

charge of the kinetic micelle is neutralized by these counterions. Counterions

are bound primarily by the strong electrical field created by the headgroups but

12

also by specific interactions that depend upon headgroup and counterion

type.' ' ''"' A two-site model has been successfully applied to the distribution of

counterions, i.e., they are assumed to be either "bound" to the micellar

pseudophase or "free" in aqueous phase."''*' * The headgroup and counterion

concentrations in the interfacial region of an ionic micelle are on the order of

3-5 mol dm''', which gives the micellar surface some of the properties of

concentrated salt solutions. ' '" Although the solution as a whole is

electrically neutral, both the micellar and aqueous pseudophases carry a net

charge because thermal forces distribute a fraction of the counterions radially

into the aqueous phase.'' ' ^

To summarize, micelles, with dilute surfactant, are typically spherical

but, with increasing [surfactant], and on addition of hydrophobic solutes or low

charge density ions, micelles grow and become ellipsoidal. Twin-tailed

surfactants may form vesicles, addition of cosurfactants and nonpolar solvents

generates microemulsions, and more ordered structures form with more

concentrated surfactant.''"^*''''

Many factors, including temperature, and the concentration of

surfactant, electrolyte, and cosurfactant, determine the shape of surfactant

association structures. Packing considerations constitute a factor which

involves the nature of the head and tail groups of the surfactant. A critical ratio

(Rp) with associated limits for several of the possible aggregation shapes has

been devised by Ninham and his coworkers''"''"

13

where Vj, is the volume of the amphiphile's hydrocarbon tail, Ap is the optimum

cross-sectional area per amphiphile molecule, and / is the length of the fully

extended hydrocarbon tail.

The optimum cross-sectional area is determined experimentally by X-

ray diffraction of bilayer systems, while the volume and length of the

hydrocarbon tail may be calculated following Tanford formulas

F/, = (27.4 +26.9 n)A^ (1.2)

/, = (1.5+ 1.265 n) A (1.3)

where n is the number of methylene groups in the hydrocarbon chain.

The following rules have been derived for predicting the dependence of

43 44

structure on the surfactant packmg parameter: '

R.=-F";, System Architecture

AoK

1 spherical micelles < -

\_ 1 rod-shaped micelles 3 ' I

J . vesicles or bilayers, 3-component o/w 2

and bicontinuous microemulsions

> 1 reverse micelles, w/o microemulsions

14

Thus, the packing parameter allows the choice of a surfactant for a desired type

of organized fluid system from its molecular dimensions (see Fig. 1.2 also).

The Role of Micelles in Chemical Reactions:

Much work on the role of micelles in chemical reactions has been

reported in recent years. ''* ' ' The micelles can play two roles: firstly, they can

lower or raise the free energy of the transition state relative to the ground state,

thereby enhancing or retarding the reaction rate; secondly, through

solubilization and adsorption of counterions in the micelle double layer, they

can increase the effective concentrations of the reactants at the micelle surface,

which is where the reactions probably occur. The first role is closely related to

medium effects in general, and is not well understood. Important factors could

include differences in reactant environment, such as the lower dielectric

constant in the Stern layer and the effects of large electric fields in the micelle

double layer on bond strengths. The second role is better understood. Often, the

observed trends can be explained quantitatively by using concepts such as

partition of organic solutes between micelles and solutions and ion exchange in

the Stern layer.

B. Oxidation bv Permanganate and Colloidal Manganese Dioxide

Permanganate ion has been extensibly used and studied as an oxidizing

agent. The amount of oxidation brought about by one permanganate ion varies

with the pH of the medium. In alkaline, neutral, and weakly acid solution, the

15

V/OhL

1/3 X / •

1/2 X / ^

• >

Aggregate shape

Spherical micelles

/ - N . Cylinders J( that

'^ may be f lex ib le)

Flexible lamella ve-

Lamellar phases

Type of sur factant

Single chain

Ionic or zwitterionic

Single chain Non-ionic or ionic wi th added salt

Double chain

^ V ^ Rzwzrsz micelles

• ^ ^ (in apolar solvents)

Double chain Small area per headgroup

Fig. 1.2: Schematic diagram of possible aggregate sliapes according to tiie packing

factor, R,, = ViMolc, criterion.

16

valence change of the manganese is from VII to IV. In strongly acid solution,

the permanganate ion is further reduced, the valence of manganese changing

from VII to II.

In oxidations by permanganate different stages of the reaction often

involve different oxidative processes, and the kinetics of the overall reaction

may be highly complex. Furthermore, many oxidations by permanganate are

carried out under conditions where hydrated manganese dioxide is precipitated,

and complications may arise from heterogeneous reactions at the surface of the

precipitate.

Manganese dioxide, the brown insoluble material whose degree of

hydration varies, is the nonnal form of manganese(IV). However, under certain

circumstances soluble forms can exist. Manganese dioxide is the normal end

product of permanganate oxidation except in strongly alkaline solution where

manganate is stable enough to be the terminal stage and in certain oxidations in

acid solution. Manganese dioxide is sometimes rather slowly precipitated,

particularly in the presence of phosphate ions, and a neutral solution containing

a yellow brown manganese(IV) species, attributed to H2Mn04 ', has been

found to obey Beer's law. ^

The use of manganese dioxide has as yet been somewhat restricted,

possibly because difficulty has been encountered in obtaining suitable active

material and in predicting the optimum reaction conditions. Unfortunately

hydrated manganese dioxide is difficult to define,"' "* and material of consistent

17

activity is only when the method of preparation has been carefully controlled.

Manganese dioxide prepared by the pyrolysis of the carbonate or oxalate failed

to oxidize allyl alcohols, whereas good yields of aldehyde were obtained if the

dioxide had been first washed with dilute aqueous nitric acid and then dried. In

carrying out kinetic studies yet another restriction is imposed-the manganese

dioxide being present in the form of a fine powder, intermediates formed at the

solid/solution interface cannot optically be detected under these conditions,

although their existence is inferred from the rate measurements.

For several decades, as mentioned above, it has been known that many

permanganate reactions carried out in both aqueous ^ and organic media ^ lead

to the formation of soluble brown-yellow manganese species. Despite its

extensive occurrence, a survey of the permanganate literature reveals that what

appears to be the same species (from both aspect and the UV-vis spectrum) has

received a number of different denominations depending on the nature of the

chemical used to reduce permanganate and the existence or not of buffers in the

medium. Among others, Mn(III)^' and Mn(V) ' ^ reaction intermediates, as

well as several Mn(IV) species such as H2Mn03,' " ^ H2Mn04 ",' and an

undefined Mn(IV)-phosphate complex, ' have been proposed as possible

chemical formulas for the substance detected. It has been suggested that this

species might be a soluble form of colloidal manganese dioxide, ' " but due to

the lack of a clear-cut demonstration some doubts about this interpretation have

arisen.**'

18

Perez-Benito and his co-workers*^ have reported the direct evidence for

this species being a soluble form of colloidal manganese(IV). The soluble

Mn(IV) species was prepared from the reduction of aqueous permanganate by

sodium thiosulfate under neutral conditions. The dark brown solution so

obtained remained perfectly transparent, precipitation not having been seen

even after several months had elapsed. Thus phosphate buffers were not needed

to keep the Mn(IV) species in solution. This species was the same that had

been previously detected in many permanganate reactions.

Precipitation of manganese dioxide was observed when the initially

transparent solutions were stirred in the presence of different electrolytes. It is

found that divalent ions are better coagulating agents than monovalent and for

cations with identical charge, the coagulating efficiency increases with the

ionic radius.

While a change of cation implies a different efficiency of the

corresponding electrolyte as coagulating agent, the same does not apply as far

as the anion is concerned. On the other hand, addition of a protective colloid

results in an increase in the concentration of electrolyte necessary for

precipitation to occur. Gum arable acts as a protective colloid while PbCb acts

as a coagulating agent.

All these results suggest that the soluble Mn(IV) species is present in the

medium in the form of colloidal particles of manganese dioxide with a negative

19

electrostatic charge, which accounts for their stability in solution due to the

adsorption of anions (such as sulfate ions) on their surface.

C. Statement of the Problem

The use of colloidal solutions has been quite successful to study fast

interfacial processes.*'' Consequent upon characterization, colloidal Mn02 -

organic acid reactions have also been studied in some depth focusing on the

determination of kinetic parameters.*'*'*' However, micellar effect on the redox

reactions of water soluble-colloidal Mn02 has not been studied so far except

only on the oxidation of formic and oxalic acids.**'*' Therefore, the present

study was designed to extend knowledge on the behavior of colloidal Mn02 as

an oxidant and to gain further insight on the role of both electrostatic and

hydrophobic interactions on the basis of kinetics of redox reactions carried out

in micellar organized media. As such, the subsequent portion of the thesis deals

with systematic kinetic studies of the reactions of colloidal manganese dioxide

with some organic acids, i.e., glycolic, lactic, mandelic, DL-malic, DL-tartaric

and citric acids, both in aqueous as well as in micellar media. Relevant

experimental details are given in Chapter 2 and Chapter 3 comprises of Results

and Discussion.

20

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26

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27

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CHAPTER 2

EXPERIMENTAL

29

Materials

The chemicals used throughout the study are listed in Table 2.1. All the

surfactants (CTAB, SDS and TX-lOO) and chemicals were used as received.

The oxidant, water-soluble colloidal manganese dioxide, was prepared as

detailed below.

The water used to prepare solutions was double-distilled over alkaline

KMn04 in an all-glass (Pyrex) distillation set-up. The specific conductivity of

the water was in the range (1-2) x 10' S cm''.

Preparation of Water-Soluble Colloidal MnO:

Jones and Noyes' procedure was used to prepare the KMn04 solution

which was standardized by titration with standard sodium oxalate solution.

Water-soluble colloidal Mn02 (5.0 x lO""* mol dm' solution) was prepared by

reduction of potassium permanganate by sodium thiosulfate in neutral aqueous

solution according to the following stoichiometry:^"''

8Mn04"+3S2O3"" + 2H^ — • 8Mn02 + 6S04^" + H2O

A 2 dm"' volumetric flask was filled with water to around 3/4 of its capacity. 20

cm^ of standard solution of Na2S203 (1.88 x 10' mol dm'^) were added. Then,

the required volume of KMn04 solution (5 cm\ 20.0 x 10' mol dm'"') was

added slowly and the reaction mixture was diluted till 2 dm^ with more water.

Each addition was followed of homogenization by gentle shaking. The

30

3 ON

o 0\ OS

OS ON OS

Al

c o U

OS

5 t CQ c

p <

-a c

on

u c

^3

«

c u o

I

CQ

X y +

z w->

+ C3

z 1

o 00

O

X

u

X

o X

(N

X

u u o Y

9

t Os II

O O u

I

u-I

o X

VI

3

1 u

vx

3 u 3

e

C .2 '•3 R *> « I .

E

< E -U

C/3 Q on

o o I

X H

V s c/a

'1 o

p

'c O E £ ,2

"G .E 'C

U

y * ^ N

c«

O •o £ _3

O 00

^-v

1

"ftb-S

D >~>

•£ -= tS ex

d o

,.—*s

o

s:

a •^ ^

0^

yo o

o

o >>

o / " " V CC

oo rn II

N3

o OS

VI

m ON ON

o\ ON

ON ON

in ON ON

31

o u 00 ON

o c 5 to CQ c

u c C3

"O ;/3

,_^ KJ

^ ^—'

U

c y:: -d t/3

/—» 2 ^

O O U

I r^ X X u—u

I

o

o o V ^ u—u

I

X o o u 1

X u-1 o

3:

X o o u 1

X -u

n: o o o

1 t

X 1 1

o a:

X o o u

1

X 1 1

o X

X o o u 1 X u-

X o o u 1

-u -t 1

o X

X o o u 1 X

-u

o

2

>—s ON

rn II ^ CQ

D.

o C3

_o 13 •T3 C c3

5

\ ^ - N

oq ro II —

o.

o

I

Q

m

a. o U

I

Q

II

II !« - o

3 u

<u

m

I 60 C

S 3 CO CO

o

PS

32

ON ON

C o OS

OS OS ON

4> ^-N

73 >S W

^ o

s •

/ ^ * S

•^ T3 C C^

u c G "d c/i

/--N «

c " ^

DS Q U

^ ^ CO

"3 ^ "—'

w

o X

o GO

ce

O u

o

o C u

0)

<L) • * — •

^ 3 M O

S —• E 3

-5 o

00

rs 'o C3

o _o J =

CU

>-.

1

1

1)

^ • ^

3 C/l

;;^ ! ^ ID crt (U c CO 00 c CO s

<L) O

'S o ^ CJ

<

X ) CO

a B 3

o JO (D

33

resulting solution was dark brown and perfectly transparent and remained

stable for several weeks, and aliquots from it were taken when required.

Characterization of Colloidal MnOi

The colloidal solution so prepared was characterized by means of UV-

vis spectrum that showed a large band covering the whole visible region of the

spectrum, with absorbance unifomily decreasing with increasing wavelength,

as well as a single broad band of high intensity at 390 nm (Fig. 2.1). From the

results of Fig. 2.1 and previous results^ it is confirmed that the stabilized

spectrum is that of water-soluble colloidal Mn02. The wavelength 390 nm was

chosen to confirm the fulfillment of the Beer-Lambert law by monitoring the

absorbance of different solutions of colloidal Mn02. The results are represented

graphically in Fig. 2.2 which shows that the law is obeyed for the concentration

range used in the present investigations. It is well known that the precipitation

of manganese dioxide was observed on the addition of different electrolytes."'

Therefore, the effect of adding different salts (LiCl, NaCl, KCl, NH4CI, BaClj,

FeS04, MnCOs, CUCI2, and SrCb) in the transparent solution of water-soluble

colloidal Mn02 was studied. In presence of these electrolytes, precipitation of

manganese dioxide occurs, which suggests that the soluble Mn(IV) species is

present in the form of colloidal particles of manganese dioxide. The negative

logarithm of the minimum [electrolyte] (necessary for precipitation) is plotted

against the cationic radius (Fig. 2.3). These results are in good agreement with

34

0.6 -

0.4 -u c o i . O M

JO <

0.2 -

0.0 340 420 500

Wave length(n m) 580

Fig. 2.1: Absorption spectra of KMn04 (1.6 x 10" mol dm~ ) and of the reaction

product of KMn04 (8.0 x 10" mol dm" ) andNa2S203 (3.0 x 10" mol dm" ).

35

0.0 L.O 8.0

10^CMnO2D(moldm"3)

Fig. 2.2: The Bccr-Lambcrt plot for colloidal Mn02 at 390 nm.

36

0.06 0.10 O.U

Cotionlc rodius { nm )

Fig. 2.3: Plots of log [electrolyte] versus cationic radius.

Reaction conditions: [Mn02] = 8.0 x 10" mol dm"\ temperature = 30 °C.

37

the observations of Perez-Benito et al? and account for the stability of colloidal

particles in solution due to the adsorption of anions on their surface.

Finally, dynamic light scattering (DLS) measurements ' indicate that

the colloidal particles are roughly spherical with a radius of about 57 nm.

Kinetic Measurements

The kinetic experiments were carried out at constant temperatures

controlled within ±0.1 °C in a oil-bath, which was designed and assembled in

the laboratory with commercially available components. A mixture containing

appropriate amounts of all the reactants (except oxidant) was taken in a three-

necked reaction vessel equipped with a double-surface water condenser to

prevent evaporation. The reaction vessel containing the resulting solution was

then immersed in the oil-bath at the required temperature and the solution was

left to stand for 15-20 min to attain equilibrium.

The reaction was then started by adding the requisite volume of

thermally equilibrated oxidant (colloidal Mn02) solution. Purified N2 gas (free

from CO2 and O2) was bubbled through the reaction mixture for stirring as well

as to maintain an inert atmosphere. The zero-time was recorded when half of

the colloidal Mn02 solution has been added. The progress of the reaction was

monitored spectrophotometrically by following disappearance of the colloidal

Mn02 at 390 nm (the X„,^^ of colloidal Mn02) with Bausch & Lomb

Spectronic-20 spectrophotometer. The reaction was usually followed up to not

38

less than 80% completion. Pseudo-first-order conditions were maintained by

keeping the [reductant]T in excess.

Values of the pseudo-first-order rate constants in the absence (kobs) and

presence of surfactants (k^) were estimated from the slopes of the conventional

log(absorbance) versus time plots by least-squares regression analysis of the

data. Multiple kinetic runs showed the data were reproducible within ± 5%.

The dependence of obs or v were obtained as a function of [oxidant],

[H" ], [reductant], [Mn(II)], [surfactant], temperature, and [gum arabic]. The

results are given in Chapter-3.

Detection ofCO:

The experimental set up and the reaction conditions for the estimation of

CO2 were the same as detailed under the kinetic measurements. The evolved

CO2 was swept out by a constant current of purified nitrogen gas and was

absorbed in standard barium hydroxide solution. But, due to the lower

concentration of reductants, the carbon dioxide evolved was not sufficient to be

detected, even though the temperature was raised up to 40 "C. Therefore, higher

reductant concentrations were used to observe the CO2 evolution. The results

are consistent with the report that CO2 is one of the reaction products from the

oxidation of reductants by colloidal Mn02.

39

Free Radical Detection

Presence of free radicals in the reaction mixtures were tested by using

the monomer acrylonitrile. Addition of the monomer solution to the reaction

mixtures containing [Mn02]T = 8.0 x 10' mol dm' , [reductantjx = 16.0 x 10'''

mol dm' and [HC104]T = 10.0 X lO'"* mol dm" at 30 "C lead to appearance of a

precipitate of white polymeric product. The positive response indicated in situ

generation of free radicals in the oxidations (no precipitation occurred in the

absence of reductant or oxidant).

40

References:

1. T. J. Jones and R. M. Noyes, J. Phys. Chem., 87,4686 (1983).

2. J. F. Perez-Benito and D. G. Lee, Can. J. Chem.. 63, 3545 (1985).

3. J. F. Perez-Benito, E. Brillas and R. Pouplana, Inorg. Chem.. 28, 390

(1989).

4. J. F. Perez-Benito and C. Arias, J. Colloid Interface ScL. 152, 70 (1992).

5. ''CRC Handbook of Chemistry and Physics", CRC Press, Inc., West

Palm Beach, Florida, 58 ^ edn., 1977-1978.

6. I. H. Segal, ''Biochemical Calculations", Wiley, New York, 2"'' edn.,

1968.

7. J. F. Perez-Benito, C. Arias and E. Amat, J. Colloid Interface ScL, 111,

288(1996).

8. J. Mata, D. Varade, G. Ghosh and P. Bahadur, Colloids Surf A:

Physicochem. Eng. Aspects, 245, 69 (2004).

9. J. Mata, T. Joshi, D. Varade, G. Ghosh and P. Bahadur, Colloids Surf

A: Physicochem. Eng. Aspects, 247, 1 (2004).

CHAPTER 3

RESULTS AND DISCUSSION

42

A. Reactions in Absence of Surfactants

As has already been mentioned, kinetic studies of oxidation of organic

acids (glycolic, lactic, mandelic, malic, tartaric and citric) were performed

spectrophotometrically by following disappearance of the oxidant (colloidal

Mn02) at 390 nm in the absence and presence of surfactant micelles at several

concentrations of the reactants, hydrogen ion, manganese(II), gum arable and

surfactants.

The log( 39o) versus time plots (Figs. 3.1 - 3.6, insets) show that the

reactions of the acids and colloidal Mn02 proceed in two stages and that the

first stage (noncatalytic) is relatively slower than the second (autocatalytic). '

The kinetic curves (Figs. 3.1 - 3.6) show inflection and the obs values (for both

the stages) were estimated from such curves. It is observed that extent of

noncatalytic path depends upon the reaction conditions (see Figs. 3.1 - 3.6 for

the effect of substrate concentration; similar effects were observed with

increase in [Mn02] as well as temperature). Kinetics of the reactions were,

therefore, investigated at several initial reactant concentrations (as the increase

in temperature resulted in disappearance of the noncatalyltic path, the kinetic

runs were performed at 30 °C — at this temperature the progress of the

reactions was neither fast nor slow) and temperatures both in the absence and

presence of surfactants and the results are summarized in Tables 3.1-3.6.

43

"" 1.4 c o u o

<

o

5 1-0

0.6

—

1

n ^ N °

Q ^

c

I

1' l1.i

o •+1.0 <v* (

\

1 1 D 20 40

Time Imin.)

20 Time (min.)

40

Fig. 3.1: Plot of log(absorbance) versus time. Inset - plot of log(absorbance)

versus time for curve (a). Reaction conditions: [MnO:] = 8.0 x 10' mol dm~\

[glycolic acid] = 24.0(a), 48.0(b), 96.0 x 10" mol dm-^(c), fHC104l = 10.0 x 10"

mol dm"^ temperature = 30 °C.

44

20

Time (min.)

Fig. 3.2: Plots of !og(absorbance) versus time. Inset - plot of log(absorbance)

versus time for curve (a). Reaction conditions: [MnOa] = 8.0 x 10" mol dm'\

[lactic acid] = 2.5(a). 5.0(b). 7.5 x 10" mol dm"-(c), [HCIO4] = lO.O x 10"" mol

dm" , temperature = 30 °C.

45

U c o u o

< Ol o +

10 Time (min)

Fig. 3.3: Plots of log(absorbance) versus time. Inset - plot of log(absorbance)

versus time for curve (a). Reaction conditions: [Mn02] = 8.0 x 10' mol dm"\

[mandelic acid] = 16.0(a), 24.0(b), 40.0 x 10"* mol dm"^(c), [HCIO4] = 10.0 x 10'"

mol dm""', temperature = 30 °C.

46

(U

u c a c o lA

<

O

+

40 Time (min.)

Fig. 3.4: Plots of log(absorbance) versus time. Inset - plot of log(absorbancc)

wrxnM lime for curve (ii), Ri'iirdnn voiulllioiiw fMnO,.] - HO ,\ 10 ^ iiml dm \

[malic acid] = 16.0(a). 24.0(b), 56.0 x 10"* mol dm"^(c), temperature = 30 °C

Inset - plot of log(absorbance) versus time for curve (a).

47

AO Time (min . )

Fig. 3.5: Plots of log(absorbance) versus time. Inset - plot of log(absorbancc)

versus time for curve (a).

Rcaclinn coiuUlhns: [MnO,] = 8.0 x 10" inol cim"\ [tartaric acid] = 16.0(a),

32.0(b), 56.0 X 10-" mol dm-^(c), [HCIO4] = 10.0 x 10"" mol dm-\ temperature =

30 °C.

48

1.A -

o c o u o tA

<

+

12 18

Tlme(min.)

Fig. 3.6: Plots of log(absorbance) versus time. Inset - plot of iog(ubsoibance)

versus time for curve (a). Reaction conditions: [Mn02] = 8.0 x 10 "*' mol dm ^

[citric acid] = 16.0(a), 24.0(b), 40.0(c), 56.0 x lO"* mol dm \d), temperature = 30

°C. Inset - plot of log(absorbancc) versus lime for curve (a).

49

Table 3.1

Dependence of rate constants on the factors influencing the oxidation of

glycolic acid by colloidal Mn02.

Non-variables Variables 10 kohs\l 10 ko\is2l

10^[MnO2]/mol dm~^

-3 [glycolic acid] = 96.0 x 10 mol dm~\ [HCIO4] = 10.0 X 10"'' mol dm"\ Temp. = 30 °C, X„,ax = 390 nm, [Mn(II)] = 0

2.0 3.2 4.0 4.8 6.0 8.0 10.0

9.2 8.9 8.7 8.1 7.7 7.3 6.6

12.8 12.2 11.6 11.0 10.2 9.0 7.1

[gum arabic]/gm dm'

[UnOj] = 8.0 X 10"^ mol dm"\ [glycolic acid] = 96.0 X 10~^ mol dm- \ [HCIO4] = \0.0 X 10~^ mol dm~\ Temp. = 30 °C, X,3, = 390 nm, [Mn(n)] = 0

[MnOj] = 8.0 X 10"^ mol dm"\ [glycolic acid] = 96.0 X 10"^ mol dm"\ Temp. = 30 °C, Xn,,, = 390 nm, [Mn(II)] = 0

1.0 2.0 3.0 4.0 5.0 6.0

4.8 4.4 4.2 3.7 3.2 3.0

not observed not observed not observed not observed not observed not observed

10'[HClO4]/mol dm -3

0.0 5.0 10.0 15.0 20.0 25.0 30.0

5.8 6.7 7.3 8.4 9.0 9.7 10.2

7.3 8.0 9.0 10.0 11.4 11.7 12.8

Contd.

50

Temp. / °C

[MnO.l = 8.0 X 10"^ mol 25 5.5 7.0 \ -3 7 , r -^1 - QA n 30 7.3 9.0 dm ' , [glycohc acid] - 96.0 ^^ ^^ 2 16.2 X 10""' mol dm"\ [HCIOA] = 40 15.9 22.4

- 4 „ „ i j „ - 3 10.0 X 10"' mol dm' ' , X 390 nm, [Mn(II)] = 0

max

Parameters noncatalytic autocatalytic

EJki mol AH*/kJ m o r ' A5*/JK"' mol"' AGVkJ m o r '

[MnOz] = 8.0 X 10"^ mol dm~^, [glycolic acid] = 96.0 X 10"^ mol d m ' \ [HCIO4] = 10.0 X 10"'' mol dm"\ Temp. = 30 °C, ?L,„a, = 390 nm

63 61 -104 92

10^[Mr

5.0 10.0 20.0 30.0 40.0 50.0

67 64 -91 92

i(II)]Vmol dm"^

8.1 8.3 8.5 9.8 10.3 10.7

10.7 12.1 12.8 14.1 15.5 16.2

"added as MnSO^

51

Table 3.2

Dependence of rate constants on the factors influencing the oxidation of

lactic acid by colloidal Mn02.

Non-variables Variables lOUobsl/ lO'^^obs:/

10^[MnO2]/mol dm"

[lactic acid] = 2.5 x 10"^ mol dm"\ [HCIO4] = 10.0 X 10" mol dm~^, Temp. = 30 °C, X,,,, = 390 nm, [Mn(II)] = 0

2.0 3.2 4.0 4.8 6.0 7.2 8.0

5.9 5.1 4.5 4.3 3.6 2.8 2.6

8.4 5.9 7.2 6.9 6.1 5.7 5.4

[gum arabic]/gm dm -3

[MnOz] = 8.0 X 10"^ mol d n r \ [lactic acid] = 2.5 x 10"^ mol dm- \ [HCIO4] = 10.0 X 10-^ mol dm- \ Temp. = 30 °C, X ax = 390 nm, [Mn(II)] = 0

1.0 2.0 3.0 4.0 5.0 6.0

3.8 3.4 3.3 3.2 3.1 2.7

not observed not observed not observed not observed not observed not observed

10''[HClO4]/mol dm -3

[MnOz] = 8.0 X 10"^ mol dm"\ [lactic acid] = 2.5 x 10"^ mol dm"\ Temp. = 30 °C. ?L,„ax = 390 nm, [Mn(II)] = 0

0 5 10 15 20 25 30

1.9 2.3 2.6 3.6 4.1 4.5 5.1

3.5 3.8 5.4 5.8 6.4 7.3 8.2

Contd.

52

[MnO:] = 8.0 x 10"^ mol dm""', [lactic acid] = 2.5 x 10"^ mol dm- \ [HCIO4] = 10.0 X 10" 390 nm, [Mn(II)] =

Parameters

mol dm "'. Xr

0

EJk] mol A//*/kJ mol"' A^-VJK"' mor AG*/kJ mol"'

Temp. / °C

25 30 35 40

1.8 2.6 4.7 7.9

noncatalytic autocatalytic

74 72 76 95

54 51

- 137 93

4.4 5.4 7.8 11.8

10^[Mn(II)]Vmol dm"

[MnOz] = 8.0 X 10"^ mol dm~\ [lactic acid] = 2.5 x 10-2 j^Qj jj^-3^ [HCIO4] =

10.0 X lO"'* mol dm"\ Temp.

= 30 °C, A.,„ax = 390 nm,

5.0 10.0 20.0 30.0 40.0 50.0

2.3 3.0 3.8 4.3 4.8 5.0

5.5 5.8 6.4 7.1 7.5 7.7

"added as MnSOa

53

Table 3.3

Dependence of rate constants on the factors influencing the oxidation of

mandelic acid by colloidal Mn02.

Non-variables Variables lOUobsi/ 10 obs2/

10^[MnO2]/mol dm"^

[mandelic acid] = 16.0 x 10"" mol dm-^ [HCIO4] = 10.0 X lO""* mol dm"\ Temp. = 30 °C, Xmax = 390 nm, [Mn(II)] = 0

[MnO.] = 8.0 X 10"^ mol dm'"', [mandelic acid] = 16.0 X 10"'' mol dm"\ [HCIO4] = 10.0 X lO""* mol dm~\ Temp. = 30 ^C, X,„ax = 390 nm, [Mn(II)] = 0

2.0 3.2 4.0 4.8 6.0 7.2 8.0

11.0 10.0 10.0 10.0 9.6 9.6 9.0

26.9 26.0 25.0 25.0 23.0 21.0 18.3

[gum arabic]/gm dm -3

1.0 2.0 3.0 4.0 5.0 6.0

5.8 5.4 5.3 5.1 4.8 4.6

not observed not observed not observed not observed not observed not observed

10''[HClO4]/mol dm-

[Mn02] = 8.0 X 10"^ mol dm""', [mandelic acid] = 16.0 X lO""* mol dm"^ Temp. = 30 °C, Xn,ax = 390 nm, [Mn(Il)] = 0

0 5 10 15 20 25 30

6.1 7.1 9.0 9.6

. 10.1 10.8 12.2

12.8 15.5 18.3 20.8 22.8 25.1 26.9

Contd.

54

Temp. / °C

[MnO,] = 8.0 X 10"^ mol 25 5.8 13.4 . -3 ; . ,• -^1 - 1/; n 30 9.0 18.3

dm [mandehc acid] - 1 6 . 0 ^^ ^^.8 29.6 X 10 mol dm , [HCIO4] = 40 152 38.4 10.0 X lO""* mol dm"^ Ji ax = 390 nm, [Mn(II)] = 0

Parameters noncatalytic autocatalytic

EJkJ moP' A//*/kJ mol"' A5*/JK"' m o r ' AG*/kJ mol"'

10^[Mn(II)]Vmol dm~^

55 53 130 92

58 56

- 112 90

[MnO dm"\

,1

X 10"^ 10.0 X = 30 °

"added

2] = 8.0 X [mandelic mol dm"''.

10"^ mol acid] = 16.0

J

. [HCIO4] = : 10"'' mol dm""*, Temp. r y =

as MnSOd

390 nm

5.0 10.0 20.0 30.0 40.0 50.0

9.6 10.7 10.7 12.6 13.4 14.1

20.5 22.7 23.0 24.0 26.4 27.4

55

Table 3.4

Dependence of rate constants on the factors influencing tiie oxidation of

malic acid by colloidal Mn02.

Non-variables Variables 1 OUobs i / 10 K^J

[malic acid] = 16.0 x 10"^ mol dm"^ [HCIO4] = 10.0 x lO"** mol dm"\ Temp. = 30 °C, ?.n,ax = 390 nm, [Mn(II)] = 0

[Mn02] = 8.0 X 10"' mol dm"^, [malic acid] = 16.0 x lO"'' mol dm- \ [HCIO4] = 10.0 X 10-" mol dm- \ Temp. = 30 °C, ?L,„ax = 390 nm, [Mn(II)] = 0

[MnOz] = 8.0 X 10"' mol dm"'', [malic acid] = 16.0 x 10"'' mol dm- \ Temp. = 30 °C, ?.,„ax = 390 nm, [Mn(II)] = 0

10'[MnO2]/mol dm -3

2.0 3.2 4.0 4.8 6.0 7.2 8.0 8.8 10.0

4.6 4.5 4.3 3.2 3.1 2.9 2.9 2.7 1.9

7.2 6.8 6.1 6.0 5.7 4.7 4.6 4.4 4.3

[gum arabic]/gm dm"

1.0 2.0 3.0 4.0 5.0 6.0

3.3 3.0 2.9 2.6 2.4 2.3

not observed not observed not observed not observed not observed not observed

10"[HClO4]/mol dm"

0 5 10 15 20 25 30

2.9 3.6 3.8 4.0 4.2 4.5 4.9

4.6 5.5 6.3 7.1 8.5 9.3 10.1

Contd.

56

Temp. / °C

[MnOs] = 8.0 X 10"^ mol 25 2.1 2.8

dm- \ [malic acid] = 16.0 x ^^ II I't A 1 j-> j - o y . o

lO"* mol dm"\ [HCIO4] = 40 95 22 8 10.0 X 10"^ mol dm~\ X^^ax = 390 nm, [Mn(II)] = 0

Parameters noncatalytic autocatalytic

EJkJ mol" A/T/kJ m o r ' A5*/JK-' m o r ' AGVkJ m o r '

[MnOz] - 8.0 X 10"' mol dm~^, [malic acid] = 16.0 x 10"^ mol dm- \ [HCIO4] = 10.0 X 10"^ mol dm~\ Temp.

= 30 °C, ?.,„ax = 390 nm

"added as MnS04

85 83 -39 95

10^[Mn(II)]'

5.0 10.0 20.0 30.0 40.0 50.0

118 116 -60 97

Vmol dm"^

3.8 4.0 4.5 4.9 5.4 6.0

6.5 6.9 7.4 8.4 9.4 10.1

57

Table 3.5

Dependence of rate constants on the factors influencing the oxidation of

tartaric acid by colloidal MnOa-

Non-variables Variables lO ' ^obs l /

„ - i

10^[MnO2]/mol dm'

[tartaric acid] = 16.0 x 10"'* mol dm'% [HCIO4] = 10.0 x 10"" mol dm~\ Temp. = 30 °C, X„„„ = 390 nm, [Mn(II)] = 0

2.0 3.2 4.0 4.8 6.0 7.2 8.0 10.0

5.0 4.4 4.2 4.1 3.8 3.3 2.7 1.6

11.9 11.3 10.9 10.8 10.7 10.5 9.6 9.4

[gum arabic]/gm dm"

[Mn02] = 8.0 X 10"^ mol dm""', [tartaric acid] = 16.0 x 10"'* mol dm- \ [HCIO4] = 10.0 X 10"^ mol dm"\Temp. = 30 °C, X„,^ = 390 nm, [Mn(II)] = 0

1.0 2.0 3.0 4.0 5.0 6.0

3.0 2.8 2.7 2.6 2.4 2.3

not observed not observed not observed not observed not observed not observed

10''[HClO4]/mol dm-

[MnOz] = 8.0 X 10"^ mol dm"\ [tartaric acid] = 16.0 x lO""* mol d m ' \ Temp. = 30 °C, ?L,„ax = 390 nm, [Mn(II)] = 0

0 5 10 15 20 25 30

2.3 2.6 2.7 3.4 3.5 3.6 3.7

9.5 9.6 6.6 9.7 10.4 11.3 12.4

Contd.

58

[MnOjJ = 8.0 X 10'^ mol dm"\ [tartaric acid] = 16.0 x 10"^ mol dm- \ [HCIO4] = 10.0 X lO""* mo

?max = 390 nm. [Mn(II)] = 0

Parameters

EJki moP' A//*/kJ mor ' A5*/JK"' mol"' AG*/kJ mol"'

1 dm" • i

Temp. / °C

25 30 35 40

noncatalytic

81 79

-52 94

2.3 2.7 5.5 9.6

autocatalytic

80 77

-48 92

6.1 9.6 13.5 28.1

10^[Mn(II)]Vmol dm -3

[MnOs] = 8.0 X 10"^ mol dm~^, [tartaric acid] = 16.0 x 10"^ mol dm- \ [HCIO4] = 10.0 X 10"'' mol dm"\ Temp. = 30 °C, ?.max = 390 nm

5.0 10.0 20.0 30.0 40.0 50.0

2.8 3.4 4.1 4.8 4.9 5.0

9.9 10.1 11.3 12.2 13.4 14.8

'added as MnS04.

59

Table 3.6

Dependence of rate constants on the factors influencing the oxidation of

citric acid by colloidal Mn02.

Non-variables Variables 10 ^obsi/ 10 obsi/

s-' s-'

[citric acid] = 16.0 x 10"^ mol dm"\ Temp. = 30 °C, ?-max = 390 nm, [Mn(II)] = 0

[MnOs] = 8.0 X 10"^ mol dm"\ [citric acid] = 16.0 x 10"^ mol d n r \ Temp. = 30 °C, X,r„^ = 390 nm, [Mn(II)] = 0

[MnOa] = 8.0 x I0~^ mol dm"\ [citric acid] = 16.0 x lO""* mol dm"\ Temp. = 30 °C, A,n,ax = 390 nm, [Mn(II)] = 0

10^[MnO2]/mol dm"^

2.0 3.2 4.0 4.8 6.0 8.0 10.0

12.8 11.3 10.7 10.1 10.3 9.6 8.4

35.2 34.5 34.1 33.1 33.6 32.9 30.1

[gum arabic]/gm dm'

1.0 7.0 2.0 6.7 3.0 6.4 4.0 5.4 5.0 5.3 6.0 4.8 7.0 5.5 8.0 3.7

10^[HClO4]/mol dm"^

0 9.6 5 11.8 10 12.8 15 13.2 20 14.8 25 15.8 30 16.7

not observed not observed not observed not observed not observed not observed not observed not observed

32.9 34.5 34.8 36.2 38.4 40.9 44.3

Contd.

60

Temp. / °C

20 25 30 35 40

4.0 6.1 9.6 13.7 20.9

13.9 22.0 32.9 57.9 85.3

[Mn02] = 8.0 X 10"^ mol dm~^, [citric acid] = 16.0 x 10"" mol dm"\ ?.niax = 390 nm, [Mn(II)] = 0

Parameters noncatalytic autocatalytic

E^/k] mol"' A/T/kJ mol"' A5*/JK"' mol"' AG*/kJ mol"'

[MnOz] = 8.0 X 10"^ mol dm"'', [citric acid] = 16.0 x lO""* mol dm"\ Temp. = 30 °C, X.,„ax = 390 nm

65 73 63 71

-96 -59 92 89

10^[Mn(II)]VmoI dm"^

5.0 10.8 10.0 11.0 20.0 15.6 30.0 19.2 40.0 25.6 50.0 29.5

32.7 32.7 37.7 37.3 42.8 50.8

"added as MnSOa.

61

To assess the effect and to find out the order with respect to [Mn02], the

/Cobsi and k^bsi values (pertaining, respectively, to noncatalytic and autocatalytic

pathways) were determined at different initial MnOi concentrations at constant

[reductant] and temperature (= 30 °C). It was observed that as the initial Mn02

increases the values of both /:obsi and obs2 decrease (Tables 3.1 - 3.6). These

observations are against the norms of kinetics (the pseudo-first-order rate

constants should be independent of the initial concentration of the reactant in

defect). Such type of behavior may be due to flocculation of the colloidal

particles. In order to suppress any possible flocculafion, the experiments were

repeated in presence of gum arable (a good water-soluble polysaccharide

known to stabilize colloidal Mn02).''''* Comparison between the obs obtained in

the absence and presence of gum arable (Tables 3.1 - 3.6) suggests that no

improvement is observed. Such type of behavior has earlier been observed in

many oxidation reactions of organic oxidants by permanganate.'*"' On the other

hand, in presence of gum arable, there seems a competition between the

reductant and gum arable for the adsorption on the surface of the colloidal

Mn02 (the -OH groups of reductant and gum arable are responsible for the

adsorption—see later).

To investigate the effect of acid concentration on rate, the reactions were

studied as a function of [HCIO4] (0.0 to 30.0 x 10" mol dm" ) at fixed [Mn02]

and [reductant] at 30 °C. It was observed that the reaction rate increased

markedly with increasing [HCIO4] (Tables 3.1-3.6). The plots of tobsi and

62

obs2 versus [HCIO4] were straight lines (Figs. 3.7 - 3.12) with positive

intercepts on the y-axis, indicating acid- independent and acid- dependent

reaction paths being involved in the reduction of colloidal Mn02 by the organic

acids. Further, double-logarithmic plots of obsi and [HCIO4] yielded straight

lines with slopes of 0.25 (r = 0.9930 — glycolic), 0.48 (r = 0.9792 — lactic),

0.39 (r = 0.9844 — mandelic), 0.16 (r = 0.9596 — malic), 0.23 (r = 0.9534 —

tartaric), and 0.19 (r = 0.9604 — citric), indicating the order with respect to

[HCIO4] to be fractional for the noncatalytic pathway (see insets of Figs. 3.7 -

3.12 where kohs\ versus [H" ]" are shown to yield good straight lines).

At fixed [Mn02] (= 8.0 x 10" mol dm" ), the dependence of obsi sJ d ko\)s2 on

[reductant] were also determined at different concentrations. The results are

depicted graphically in Figs. 3.13-3.18. The values of rate constants ( obsi and

'tobs2) are given in Tables 3.7-3.12. Double-logarithmic fits between /robsi and

[reductant] yielded apparent order to be 0.67 (r = 0.9978 — glycolic), 0.96 (r =

0.9740 — lactic), 0.90 (r = 0.9940 — mandelic), 1.0 (r = 0.9895 — malic),

0.89 (r = 0.9964 — tartaric), and 0.49 (r = 0.9790 — citric). The plots of obsi

and /:obs2 versus [citric acid] show positive intercepts on the y-axis (Fig. 3.18).

It should, however, be zero at [citric acid] = 0.0 mol dm~ . Complexity in the

structure of citric acid may be the cause — citric acid is a special acid which

has tertiary -OH group in the skeleton and oxidation of tertiary -OH group is

quite diftlcult under normal conditions.

63

13.0

1. ^ 8.0

o

i n

0 ^

I 1

k o b s 2

^•"^""^ lao

W o b S l r ^

- -^^ r-

o

o

5.0

P

P

/

1 1 aiO ai5 0.2 0 0.25

CH*D0-25(mol0-25dm-0.75)

1 10 20 30

I O ^ C H C I G ^ O (moldm~3)

Fig. 3.7: Effect of HCIO4 on the pseudo-first-order rate constants.

Reaction conditions: [MnOz] = 8,0 x 10" mol dm"\ [glycolic acid] = 96.0 x 10"''

mol dm"\ temperature = 30 °C. Inset plot showing the effect of hydrogen ion

concentration on the pseudo-first- order rate constant.

64

10 20

Fig. 3.8: Effect of HCIO4 on the pseudo-first-order rate constants.

Reaction conditions: [Mn02] = 8.0 x 10" mol dm"\ [lactic acid] = 2.5 x 10" mol

dm"\ temperature = 30 °C. Inset plot showing the effect of hydrogen ion

concentration on the pscudo-first- order rate constant.

65

2 6 . 0 -

J3 O

n.o-

10 20 10'CHClO4l(moldm-3 )

Fiy. 3.9: i:riccl of I IClO.i on Ihc pscudo-nrsl order rate conslanls.

Reaclion conditions: [MnOj] = 8.0 x 10"* mol dm"^ [mandclic acid] = 16.0 x 10'"

niol dnr \ temperature = 30 °C. Inset - plots showing the effect of hydrogen ion

concentration on the pseudo-first order rate constant.

66

10.0-

J o

10^CHClO^D(moldm-3)

Fig. 3.10: Effect of HCIO4 on the pseudo-first-order rate constants.

Reaction conditions: [MnOj] = 8.0 x 10"' mol dm-\ [malic acid] = 16.0 x 10^

mol dm-\ temperature = 30 °C. Inset - plots showing the effect of hydrogen ion

concentration on the pseudo-first-order rate constant.

67

10 20 30 10^CHCl04D(moldm-3)

Fig. 3.11: Effect of HC10.| on the pseudo-first order rate constants. ri ,-3 Reaction conditions: [MnO:] = 8.0 x 10 mol dm"% [tartaric acid] = 16,0 x 10

mol dm"\ temperature = 30 °C. Inset - plots showing the effect of hydrogen ion

concentration on the pseudo-first order rnle conslnnt.

68

I/)

o

m O

2.0 h

0.0 I 10 20 30

1 0 ^ C H C I 0 4 D ( m o l d m - 3 )

Fig. 3.12: Effect of HCIO4 on the pseudo-first-order rate constants.

Reaction conditions: [colloidal MnOj] = 8,0 x 10" mol dm" , [citric acid] = 16.0

X 10"* mol dm" , temperature = 30 "C. Inset - plots showing the effect of

hydrogen ion concentration on the pseudo-first order rate constant.

69

tn

1.2 -

J2 °-® JO

o

0.^ -

—

^

9/.

1

oZ^

1

kobs2 y p

/kobsl

1 1 0.0 0 40 80 120 160

10^CGlycolicacld3(moldm~3) Fig. 3.13: Effect of glycolic acid on the pseudo-first-order rate constants.

Reaction conditions: [MnOz] = 8.0 x 10"* mol dm-\ [HCIO4] = 10.0 x lO^ mol

dm'^, temperature = 30 °C.

70

in

O

m O

^.l.

0.7

-

Ox-

1

^ ^ J ^ - ' - ' ' ' ^ kobs l

1 1 25 50

10^C Lactic o c i d l l m o l d m " ^ )

75

Fig. 3.14: Effect of lactic acid on the pseudo-first-order rate constants.

Reaction conditions: [MnOa] = 8.0 x 10" mol dm"\ [HCIO4] = 10.0 x 10"" mol

dm"\ temperature = 30 °C.

71

2 0 40

10^ QMande l i cQc id l Imo ldm" " )

Fig. 3.15: Effect of mandelic acid on the pseudo-first order rate constants. Reaction

conditions: [MnOj] = 8.0 x 10' mol dm"\ [HCIO4] = 10.0 x lO"' mol dm'\

temperature = 30 °C.

72

25 50

10^ C Malic acid3 (moldm"^)

Fig. 3.16: Effect of mandelic acid on tiie pseudo-firrt order rate constants.

Reaction condition: [MnOzl = 8.0 x 10" mol dm"^ temperature = 30 °C.

73

tn

(/) Si

o CO

O

0 25 50

10^ CTartaric acid3 (moldmT^)

Fig. 3.17: Effect of DL-tartaric acid on the pseudo-first order rate constants.

Reaction conditions: [MnOa] = 8.0 x 10" mol dm''\ [HCIO4] = 10.0 x lO""* mol

dm"^ temperature = 30 °C.

74

20 40 lO^CCi t r ic acid 3 ( m o l d m - 3 )

60

Fig. 3.18: Effect of DL-cilric acid on the pseudo-first order rate constants.

Reaction conditions: [Mn02] * 8.0 x 10" mol dm" , temperature = 30 °C,

75

Table 3.7

Dependence of rate constants on glycolic acid for its oxidation by colloidal

MnOz".

1 O^[glycolic acid]/ mol dm" 1 O^obsi/ s"' 10" k^^J s~'

24.0 3 l 4^0

48.0

72.0

96.0

120.0

144.0

''Conditions: [UnOj] = 8.0 x 10"^ mol dm-\ [HCIO4] = 10.0 x

1 0 " V o I d m " \ 3 0 ° C

4.9

6.4

7.3

9.0

10.8

6.0

7.7

9.0

11.0

12.4

76

Table 3.8

Dependence of rate constants on lactic acid for its oxidation by colloidal

MnOz".

10^[lactic acid]/ mol dm" 1 oXbsi/ s"' 10'' K\>J s"'

T 2 5 iF 18

2.5

3.75

5.0

6.25

7.50 10.2 14.9

'"Conditions: [MnOj] = 8.0 x 10"^ mol dm"\ [HCIO4] = 10.0 x

10" ' 'moldm"\30°C

2.6

4.7

6.8

8.5

5.4

7.8

9.5

11.1

Table 3.9

Dependence of rate constants on mandelic acid for its oxidation by

colloidal MnOi*.

10'*[mandelic acid]/ mol dm ^ 1 O^bsi/s' 10 ^nho/s

16.0

20.0

24.0

32.0

36.0

40.0

11.7

14.8

16.6

20.7

24.0

24.9

22.6

32.0

40.8

53.3

57.6

65.2

''Conditions: [Mn02] = 8.0 x 10"^ mol dm"\ [HCIO4] = 10.0 x

10"' 'moldm~\30°C

78

Table 3.10

Dependence of rate constants on malic acid for its oxidation by colloidal

MnOa".

10''[malicacid]/moldm"^ lOXbsi/s"' Wk'^JT'^

16.0 19 4 6

24.0

32.0

40.0

48.0

56.0

64.0

''Conditions: [Mn02] = 8.0 x 10"^ mol dm"^ 30 °C

5.3

6.7

7.8

11.0

11.8

13.3

1.1

11.2

11.5

16.0

17.6

19.6

79

Table 3.11

Dependence of rate constants on tartaric acid for its oxidation by colloidal

MnOz".

10'*[tartaric acid]/ mol dm' 1 oXbsi/ s"' 10* kotsil s"'

16.0 l 8 iTs

24.0

32.0

36.0

40.0

44.0

""Conditions: [UnOi] = 8.0 x 10" mol dm"^ [HCIO4] = 10.0 x

10"*moldm"\30°C

6.9

9.6

10.9

12.2

13.4

15.6

21.6

24.3

25.2

30.0

80

Table 3.12

Dependence of rate constants on citric acid for its oxidation by colloidal

MnO,".

10'*[citricacid]/moldm"^ lOXbsi/s"' 10''^obs2/s"'

16.0 9 6 33

24.0

32.0

40.0

48.0

56.0

''Conditions: [MnOj] = 8.0 x 10"^ mol d m ' \ 30 °C

11.3

12.6

13.9

16.2

17.7

3.4

3.9

4.1

4.3

4.6

81

In view of the above points, the empirical rate law is written as

cl[Mn02] - = {k! + Jd' [H ]"} [MnOz] [reductant]"' (3.1)

dr

Activation parameters {E^, AH* AS^, and AG*) are believed to provide

useful information regarding environment in which chemical reactions take

place. Therefore, a series of kinetic runs were carried out at different

temperatures at fixed [reductant], and [Mn02]. The obs (Tables 3.1 - 3.6) were

found to fit the Arrhenius and Eyring Eqs. (3.2) and (3.3):

obs = exp(-£a/RT) (3.2)

/' obs = (k^T/h) exp(-A/r/RT) exp(A^/R) (3.3)

where the symbols have their usual meanings. The non-linear least squares

method was used to obtain the values of parameters which are recorded in

Tables 3.1-3.6.

The following mechanism satisfying the above requirements is proposed

for the noncatalytic reaction path:

82

(a) pH - independent pathway

^1 ^adl ^1

(Mn02)n + H O — C — C O O H ^ " (Mn02) i r (0H-C-C00H) (3.4)

R2 R2 (CI)

CI U - (Mn02)n.i+ Mn(II) + (p=0 + COOH (3.5)

R2 (PI)

COOH + (Mn02)n —^-*-(Mn02)n. , + COj + Mn(III) (3.6)

HO—(p—COOH + Mn(III) J ^ HO— f ^ C O O + Mn(II) (3-7)

R2 R2 (radical)

fast I radical *• 0 = 0 + CO2 (3-8)

R.

(ii) pH - dependent pathway

^ ' ^ ^ad2 ^ '

(Mn02-);7(OH-(j:-COOH)+H' ^ ; = ^ H^-fMnOj^COH-C-COOH) (3.9)

R2 R2 (^1) (C2)

C2 • (Mn02)n.i+Mn(II)+ COOH+ PI (3-10)

(reactions (3.6) to (3.8) then follow)

SCHEME 1

83

(R, = R2 = -H glycolic acid; Ri = -H, Rj = -CH3 lactic acid; Rl =

-H, R2 = -CfiHs mandelic acid; R, = -H, R2 = -CH2COOH malic

acid; R, = -H, R2 = -CH(OH)COOH tartaric acid; R, = R2 = -CH2COOH

citric acid)

In Scheme 1, (Mn02)n stands for colloidal manganese dioxide. Eq. (3.4)

represents adsorption of the carboxylic acid on the surface of the colloidal

Mn02. After adsorption, species CI undergoes electron transfer process leading

to the formation of product (PI), radical COOH and Mn(II) in a rate

determining step. In the subsequent fast step, as depicted in Eq. (3.6), Mn02

further reacts with COOH and gives Mn(III) and CO2 as reaction products. The

observed dependence of the order on [reductant] being different (see Eq. (3.1))

indicates that the reactivity of the individual carboxylic acid (in the adsorbed

state) probably plays a role in the transfer of electrons to colloidal Mn02.

As evident from Figs. 3.1 - 3.6, the reactions (oxidation) of organic

acids with colloidal Mn02 proceed in two stages: the second being faster than

the first. The reactions are thus typically autocatalytic, i.e., catalyzed by one of

the products. Many permanganate - organic substrate reactions have been

shown to be autocatalytic'' "' whereby Mn04~ undergoes reduction (by the

organic substrate) to the Mn(II) stage, which then gets oxidised by Mn04" into

Mn(III).

MnOr + 4 Mn(II) + 8 H^ • 5 Mn(III) + 4 H2O (3.11)

84

Mn(II) is, thus, considered to be the active autocatalyst. As Mn(II) is a product

in the present reactions too, experiments were performed with initially added

Mn(II) in order to confirm its role. The rate constants, obtained as a function of

[Mn(II)], were found to increase with increasing [Mn(II)] (Tables 3.1 - 3.6).

The results are shown graphically in Figs. 3.19 - 3.24 which show sigmoid

dependence on [Mn(n)], thereby establishing Mn(II) as the autocatalyst. '"^

In the presence of Mn(II), Scheme 1 is, therefore, modified as

Scheme 2:

pathi (Mn02)n + Mn(II) »- (MnOj^-i + Mn(III) (312)

pathll (MnOj),! + reductant >• (Mn02)n-i + Mn(II) + other products (3.13)

SCHEME 2

Both Mn(II) and reductant compete to react with the colloidal Mn02 and

the values of rate constants (Table 3.1 - 3.6) are the sum of the paths I and II

(Scheme 2). Thus, the exact dependence of obsi/ obsi on [Mn(II)] cannot be

predicted. Additionally, another possibility exists involving intermediates/

products for the observed autocatalysis.

85

1 0 ^ C M n ( I I ) 3 ( m o l d m ' 3 )

Fig. 3.19: Effect of Mn(II) on the pseudo-first-order rate constants.

Reaction conditions: [MnOi] = 8.0 x 10" mol dm"\ [glycolic acid] = 96.0 x 10"

mol dm"\ [HCIO4] = 10.0 X 10"" mol dm~\ temperature = 30 °C.

86

O

o

lO^CMn ( IDD (moldm-^J

Fig. 3.20: Effect of Mn(II) on the pseudo-first-order rate constants.

Reaction conditions: [UnOz] = 8.0 x 10" mol dm'\ [lactic acid] = 2.5 x 10" mol

dm'\ [HCIO4] = 10.0 X 10"" mol dm"\ temperature = 30 °C.

87

O

lO^CMn l i n D t m o L d m - ^ )

Fig. 3.21: Effect of Mn(II) on the pseudo-first order rate constants.

Reaction conditions: [MnO^] = 8.0 x 10" mol dm" , [mandelic acid] = 16.0 x 10

mol dmr\ [HCIO4] = 10.0 x lO"'' mol dm'\ temperature = 30 °C.

-4

88

0 25 50

l O ^ C M n d D D ( m o l d m - 3 )

Fig. 3.22: Effect of Mn(II) on the pseudo-first order rate constants.

Reaction conditions: [MnOj] = 8.0 x 10"* mol dm" , [malic acid] = 16.0 x 10

mol dm" , temperature = 30 °C.

89

^^ V) X)

o

o

10SCMn(II)D(moldm"3)

Fig. 3.23: Effect of Mn(II) on the pseudo-first order rate constants.

Reaction conditions: [MnOi] = 8.0 x 10" mol dm""*, [tartaric acid] = 16.0 x 10

nu)l dm \ 11 IC10.,J 10.0 .\ 10 ' mol dm ^ Icmpcralurc - 30 "C.

-4

90

50

1 V)

(/I

** o

^_-—S--"^

1

kobs2/"^

>^*<obs1

1 25 50

lO^CMnl wO (moldm-^

Fig. 3.24: Effect of Mn(II) on the pseudo-first order rate constants.

Reaction conditions: [Mn02] = 8.0 x 10'* mol dm" , [citric acid] = 16.0 x 10" mol

dm'^ temperature = 30 °C.

91

To check the formation of Mn(III) species, studies were made by

monitoring absorbance at 470 nm (Mn(III) is the only absorbing species at this

wavelength"^). But, no build up of Mn(III) was observed during the course of

the reactions. As the reactivities of organic acids toward colloidal Mn02 are

higher in comparison to Mn(II), '' the participation of the colloidal Mn02-

Mn(II) reaction (Eq. (3.12)) in the autocatalytic pathway can either be ruled out

or might be of minor significance.

As far as the sigmoid dependence on [Mn(II)] is concerned (Figs. 3.19 -

3.24) for the autocatalytic reaction pathway, the concentration of Mn(III)

increased first and then started to decrease due to the Mn(III)-reductant

reaction (Eq. (3.7)), whereas that of Mn(II) increased, passed through

maximum, decreased, and finally showed a slow increase. The final conversion

of Mn(III) to Mn(II) seemed to occur when Mn(IV) was no longer present in

the system. '''

The total reaction rate, derived by the sum of the contributions

corresponding to pathways (i) and (ii), is:

v = {lc',+ k"nm\]} [MnOj] [(reductant)^] (3.11)

{k'l and ^''H are the rate constants corresponding to pH-independent and pH-

dependent pathways, whereas (H )s and (reductant)s refer to adsorbed species

on the colloid surface).

92

Adsorption of organic acids on the surface of colloidal particles in quasi-

equilibrium reactions prior to the redox rate-determining steps is widely

accepted."" ' " Therefore, adsorption isotherm can be used to explain the

observed results. Writing Freundlich adsorption isotherm for the species

present in the bulk solution and adsorbed on colloidal Mn02, we get, according

to equilibrium steps (3.4) and (3.9)

[(H")s] = a'[HY,and (3.12)

[(reductant)s] = a" [reductant]''" (3.13)

where a and b are the adsorption parameters for the respective species. Eq.

(3.11) can now be written as

v = {k\ a" + k"n a' a" [H^]'''}[Mn02] [reductant]''" (3.14)

Eq. (3.14) agrees very well with the experimental results (see Eq. (3.1)) with k'

= k\ a" and k" - k"i\ a' a". Moreover, the values obtained for the Freundlich

exponents are b' =0.25, 0.48, 0.39, 0.16, 0.23, 0.19 and b" = 0.67, 0.96, 0.90,

1.0, 0.89, 0.49 for glycolic, lactic, mandelic, malic, tartaric, and citric acids

respectively, in agreement with one of the requirements of the Freundlich

isotherm — it is well known that for applicability of that isotherm, the

exponent values must lie within 0 -1 interval. '

93

(B) Reactions in Presence of Surfactants

To see the role of surfactants (anionic, cationic or non-ionic), attempts

were made to carry out the oxidation reactions under different experimental

conditions, i.e., reaction mixture + SDS, reaction mixture + CTAB, and

reaction mixture + TX-100. The values of rate constants ( 4/, s'') are

summarized in Tables 3.13 - 3.24. Preliminary observations showed that a

reaction mixture containing [Mn02] (= 8.0 x 10" mol dm" ), cationic CTAB

and reductant became turbid due to flocculation. CTAB micelles have positive

charge on their head group whereas colloidal Mn02 has negative charge.''' ' ' ^

The anionic SDS surfactant neither catalyzed nor inhibited the oxidation

reactions (Tables 3.13 - 3.18). The observation is not unexpected; anionic

micelles, simply based on electrostatic considerations, will repel the oxidant

(colloidal Mn02) keeping the bulk reaction unaffected.

The effect of varying [TX-100] upon the rate of oxidation of the organic

acids by colloidal Mn02 were studied at constant [reductant], [Mn02] (= 8.0 x

10" mol dm" ), and temperature (= 30 °C). The observed data (Tables 3.13 -

3.18) are shown graphically in Figs. 3.25 - 3.30 as rate constant - [surfactant]

profiles.

To see the effects of [oxidant], [reductant], [HCIO4] and temperature,

and to confirm whether or not the aqueous medium mechanism is operative in

micellar TX-100, a series of kinefic runs were carried out at constant [TX-100]

94

Table 3.13

Dependence of rate constants on [TX-lOO] for the oxidation of glycolic acid

by colloidal Mn02''.

10^[TX-100]/moldm-^

0.0

5.0

10.0

15.0

20.0

30.0

40.0

50.0

lOX,/ s"'

7.3 (7.3, 7

7.7

8.2

8.3

8.8

8.7

8.7

8.8

.3)''

lO'^V s"'

9.0(9.1,9.0)"

12.8

13.6

15.1

16.3

16.3

16.2

16.3

''Conditions: [glycolic acid] = 96.0 x 10"^ mol dm~ , [Mn02] =

8.0 X 10" mol dm~^ [HCIO4] = 10.0 x 10"* mol dm"\ 30 "C.

''The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 x 10"^ and 20.0 x W^ mol dm"l

95

Table 3.14

Dependence of rate constants on [TX-lOO] for the oxidation of lactic acid by

colloidal MnOz".

10^[TX-100]/moldm-^ 1 0 \ , / s - ' I0X2/S"'

_ _

5.0

10.0

15.0

20.0

30.0

40.0

50.0 4.9 7.7

''Conditions: [lactic acid] = 2.5 x 10"^ mol d n r \ [Mn02] = 8.0 x

10" mol dm-\ [HCIO4] = 10.0 x lO"'' mol dm"\ 30 °C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10'^ and 20.0 x 10"^ mol dm"l

2.6 (2.6, 2.5)"

3.0

3.9

4.3

4.9

5.1

4.9

5.4 (5.4, 5.4)"

5.6

6.4

7.0

7.7

7.9

7.7

96

Table 3.15

Dependence of rate constants on [TX-lOO] for the oxidation of mandelic

acid by colloidal MnOa .

10^[TX-100]/moldm"^

0.0

5.0

10.0

15.0

20.0

30.0

40.0

50.0

lOXi/i

9.0(9.1,

10.2

11.5

12.0

13.0

13.0

13.0

13.0

r'

, 9.0)"

10 Vs" '

18.3(18.3,18.2)"

20.5

22.0

22.6

24.0

24.0

24.0

24.0

^Conditions: [mandelic acid] = 16.0 x 10"'* mol dm"''[Mn02] =

8.0 X 10" mol dm-^ [HCIO4] = 10.0 x lO'^'mol dm"^ 30 °C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10'^ and 20.0 x 10"^ mol dm"^

97

Table 3.16

Dependence of rate constants on [TX-lOO] for the oxidation of malic acid by

colloidal Mn02^.

10^[TX-100]/moldm-^ 10Xi/s~' IOX2/S"'

T o 2.9 (2.9,2.9)" 4.6 (4.5,4.6)"

5.0 3.8 4.9

10.0 4.2 5.1

15.0 4.3 5.2

20.0 4.6 5.4

30.0 3.5 4.5

40.0 3.3 4.4

50.0 3.2 4.3

""Conditions: [malic acid] = 16.0 x 10'" mol dm"\ [Mn02] = 8.0

xlO"^moldm"\30°C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10"^ and 20.0 x 10"^ mol dm"l

98

Table 3.17

Dependence of rate constants on [TX-lOO] for the oxidation of tartaric acid

by colloidal Mn02*.

10^[TX-100]/moldm"^ 10%,,/s"' lO^Vs '"

T o 2.7 (2.6,2.6)' 9.6 (9.6, 9.6)"

1.0 3.2 9.6

5.0

10.0

15.0

20.0

30.0

40.0

50.0

4.6

5.2

5.7

5.9

5.7

5.7

5.9

10.6

11.0

11.5

12.1

11.8

11.8

12.1

""Conditions: [tartaric acid] = 16.0 x 10 " mol dm \ [Mn02] =

8.0 X 10'^ mol dm-\ [HCIO4] = 10.0 x 10-Vol dm-\ 30 °C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10"^ and 20.0 x 10"^ mol dm"\

99

Table 3.18

Dependence of rate constants on [TX-lOO] for the oxidation of citric acid by

colloidal Mn02^.

10^[TX-100]/moldm"^ lOXi/s"' IOX2/S''

0.0 9.6(9.6,9.7)'' 33.0(33.1,33.0)''

I.O 12.0 43.0

5.0

10.0

15.0

20.0

30.0

40.0 16.3 42.1

50.0 14.6 37.6

''Conditions: [citric acid] = 16.0 x 10"" mol dm"\ [Mn02] = 8.0 x

10"^moldm"\30°C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10"^ and 20.0 x 10"^ mol dm'^

16.4

17.7

22.6

19.2

17.4

48.0

51.9

54.2

44.8

42.5

99

Table 3.18

Dependence of rate constants on [TX-lOO] for the oxidation of citric acid by

colloidal MnOa .

10^[TX-100]/moldm-^ 1 0 \ i / s " ' lOXi/s"'

0.0

1.0

5.0

10.0

15.0

20.0

30.0

40.0 16.3 42.1

50.0 14.6 37.6

''Conditions: [citric acid] = 16.0 x 10"" mol dm"^ [Mn02] = 8.0 x

10"^moldm"\30 °C.

The values quoted in parentheses were obtained, respectively,

at [SDS] = 4.0 X 10"^ and 20.0 x 10~ mol dn^^

9.6 (9.6, 9.7)''

12.0

16.4

17.7

22.6

19.2

17.4

33.0(33.1,33.0)''

43.0

48.0

51.9

54.2

44.8

42.5

100

Table 3.19

Dependence of rate constants on the factors influencing the oxidation

of glycolic acid by colloidal Mn02 in presence of TX-100.

Non-variables Variables s - ' S - '

[TX-100] = 15.0 X lO'^mol dm ' \ [MnO2] = 8.0x 10'^ mol dm"\ [HCIO4] = 10.0 X lO"'' mol dm~\ Temp. = 30 °C,

max - 390 nm

10^[glycolic acid]/mol dm '

24.0 48.0 72.0 96.0 120.0 144.0

4.1 5.1 7.0 8.3 10.2 11.6

6.6 9.2 11.5 15.1 16.7 19.7

10^[MnO2]/moldm"^

[TX-100] = 15.0 X 10"^ mol dm"\ [glycolic acid] = 96.0 x 10" mol dm"\ [HCIO4] = 10.0 X 10"* moidm~\ Temp. = 30 °C,

max ~ 390 nm

2.0 3.2 4.0 4.8 6.0 7.2 8.0

12.8 11.5 11.3 11.1 10.5 9.6 8.3

16.8 16.2 16.0 15.9 15.8 15.3 15.1

10''[HClO4]/mol dm-

-3 [TX-100] = 15.0 X 10"^ mol dm [glycolic acid] = 96.0 x lO"-' mol dm' l [MnOz] = 8.0 x 10"^ mol dm"" Temp. = 30 °C, max - 390 nm

0.0 5.0 10.0 15.0 20.0 30.0

7.1 8.2 8.3 9.0 10.5 10.7

11.3 12.8 15.1 16.3 17.9 20.3

Contd.

101

Temp. / °C

[TX-lOO] = 15.0 xlO'^ mol 25 6.2 10.1 d m - , [MnO,] = 8.0 X 10- 'mol ^0 O ^ 15.1

dm , [glycolic acid] = 96.0 x ^Q J9 2 34.8 10"'' mol dm"^ [HCIO4] = 10.0 x 10~^moldm"\ ?max = 390 nm

Parameters

-1 EJV.] mol A/r/kJ mol" AS-VJK-' mol"'

AG*/kJ mol"'

noncatalytic

58 58

- 112 92

autocatalytic

66 63

- 90 91

102

Table 3.20

Dependence of rate constants on the factors influencing the oxidation

of lactic acid by colloidal Mn02 in presence of TX-100.

Non-variables

[TX-100] = 15.0 X 10" mol dm'^ [MnO2] = 8.0xl0"^moldm-^ [HCIO4] = 10.0 x 10"'* mol dm-^ Temp. = 30 °C, ?iniax = 390 nm

Variables loX,/ S-'

10^[lactic acid]/ mol dm"''

1.25 3.9 2.50 4.3 3.75 5.7 5.00 8.3 6.25 9.8 7.50 11.6

S-'

6.4 7.0 9.6 11.1 14.4 18.3

10^[MnO2]/moldm"

[TX-100] = 15.0 x 10" mol dm"\ 2.0 [lactic acid] = 2.5x10"^ mol dm"^ ^"^ [HCIO4] = 10.0x10"^ mol dm-\ 4'g

Temp. = 30 °C, >-niax = 390 nm 6.0 8.0

8.1 7.3 7.0 6.0 5.1 4.3

12.8 12.2 11.7 10.0 9.4 7.0

10''[HClO4]/moldm -3

[TX-100] = 15.0 x 10" mol dm-\ 0-0 [lactic acid] = 2.5x10'^ mol dm'^ 10.0 [MnOz] = 8.0x10"^ mol dm-\ 15.0 Temp. = 30 °C, X^^^ = 390 nm 20.0

25.0 30.0

4.0 4.3 5.0 5.3 5.5 6.0

6.8 7.0 7.5 7.7 7.9 8.5

Contd.

103

Temp. / °C

[MnO2] = 8.0xl0"^moldm"^ [lactic acid] = 2.5 X 10" mol dm""', [TX-100]=15.0 [HClO4] = 10.0>

?k.max = 390 nm

Parameters

EJk} moP' A//*/kJ m o r ' Ar/JK"' mol"' AG*/kJ moP'

xlO~^moldm"^ : 10"^moldm"\

25 30 35 40

noncatalytic

50

48 -152

94

3.4 6.1 4.3 7.0 6.0 9.6 8.3 12.8

autocatalytic

43

41 -170

92

104

Table 3.21

Dependence of rate constants on the factors influencing the oxidation

of mandelic acid by colloidal MnOi in presence of TX-100.

Non-variables Variables 10X1/ 10U^2/ S-' s-'

10 [mandelic acid]/mol dm'

[TX-100] = 15.0 X lO'^mol dm"\ [MnOz] = 8.0 x 10"^ mol dm"\ [HCIO4] = 10.0 X 10" mol dm"^ Temp. = 30 °C, ?tmax = 390 nm

16.0 20.0 24.0 32.0 36.0 40.0

11.7 14.8 16.6 20.7 24.0 24.9

22.6 32.0 40.8 53.3 57.6 65.2

10^[MnO2]/moldm-^

-3 [TX-100] = 15.0 X lO'^mol dm""', [mandelic acid] = 16.0 x 10"'' mol dm~\ [HCIO4] = 10.0 x 10"''moldm-\Temp. = 30°C, A-max = 390 nm

2.0 3.2 4.0 6.0 7.2 8.0

13.5 12.8 12.5 12.0 11.7 11.7

28.1 26.9 25.6 24.0 23.0 22.6

10^[HClO4]/moldm -3

[TX-100] = 15.0 x 10"Vol dm""', [mandelic acid] = 16.0 x 10"'' mol dm"\ [MnOz] = 8.0 x 10"^ mol dm"\ Temp. = 30 °C,

max = 390 nm

0.0 10.0 15.0 20.0

25.0

30.0

9.6 11.7 12.8 13.2

13.9 14.2

21.9 22.6 26.4 28.8

33.1 34.9

Contd.

105

[MnOz] = 8.0 X 10"^ mol dm"^ [mandelic acid] = 16.0 x lO""* mol dm"\ [TX-lOO] = 15.0 x

10"^ mol dm-\ [HCIO4] = 10.0 x 10" mol dm"^ ?.„,ax = 390 nm

Parameters

EJV.} mol"' A/r/kJ m o r ' A^^/JK"' mor ' AGVkJ mol"'

Temp. /

25 30 35 40

noncat£

41 38

- 175 91

/ OQ

ilytic

8.4 14.8 11.7 22.6 13.7 35.3 18.3 44.1

autocatalytic

61 58

- 103 90

106

Table 3.22

Dependence of rate constants on the factors influencing the oxidation

of malic acid by colloidal Mn02 in presence of TX-100.

Non-variables Variables I0^k,.i/ 10%,2/ S-'

10'*[malic acid]/ mol dm ^

[TX-100] = 15.0 X 10~ mol dm-^ [MnO2] = 8.0x 10"^moldm-\ Temp. = 30 °C, Xn,ax = 390 nm

[TX-100] = 15.0 X 10~^moldm~^ [malic acid] = 16.0 x 10"'' mol dm"^ Temp. = 30 °C, max - 390 nm

16.0 24.0 32.0 40.0 48.0 52.0

4.3 7.3 9.1 10.8 11.8 13.3

10^[MnO2]/moldm-^

2.0 3.2 4.0 4.8 6.0 7.2 8.0 8.8 10.8

5.6 5.4 4.5 4.8 4.6 4.4 4.3 4.1 3.8

5.2 10.1 11.6 13.4 14.5 16.3

8.0 6.7 6.1 5.8 5.6 5.4 5.2 5.1 4.9

10''[HClO4]/moldm"^

[TX-100] = 15.0 X 10"^ mol dm-\ 0-0 4.3 5.2 [MnOj] = 8.0x 10-^ mol dm-^ 5.0 4.4 5.6 [malic acid] = 16.0 X 10"^ mol 10.0 4.8 6.4 dm-\Temp. = 30 °C, X„,,, = 390 nm ^^'^ 5.0 7.7

20.0 5.5 8.9 30.0 6.1 11.1

Contd.

107

[MnOj] = 8.0 X 10"^ mol dm~\ [malic acid] = 16.0 x lO''* mol dm"', [TX-lOO] = 15.0 x 10 mol dm"'', ?imax = 390 nm

Parameters

EJkJ mol"' A//*/kJ mol"' A5*/JK'' mol"' AGVkJ mol"'

-3

Temp. / °C

25 30 35 40

noncatalytic

80

77 - 55

94

2.3 3.6 4.3 5.2 6.2 8.7 11.1 26.8

autocatalytic

107

104 37 93

108

Table 3.23

Dependence of rate constants on the factors influencing the oxidation

of tartaric acid by colloidal Mn02 in presence of TX-100.

Non-variables Variables lOU, „-i

x^U lOU, H/2'

10'*[tartaric acid]/mol dm"''

[MnO2] = 8.0x 10"^moldm"^ [TX-100] = 15.0 X 10"^ mol d m ' \ [HCIO4] = 10.0 X lO-** mol dm~\ Temp. = 30 °C, X„,^^ = 390 nm

16.0 24.0 32.0 36.0 40.0 44.0

5.7 6.9 9.6 10.8 12.2 13.4

11.5 15.6 21.6 24.3 25.2 30.0

[TX-100] = 15.0 X 10"^ mol dm"^ [tartaric acid] = 16.0 x 10"'' mol dm-\ [HCIO4] = 10.0 X 10" mol dm"\ Temp. = 30 °C, X ax = 390 nm

10'[MnO2]/moldm"^

2.0 3.2 4.0 4.8 6.0 8.0

7.1 6.4 6.1 6.0 5.9 5.7

13.4 12.8 12.2 12.0 11.6 11.5

10^[HClO4]/moldm'

[TX-100] = 15.0 X 10"^ mol dm"^ [tartaric acid] = 16.0 x lO"'* mol dm~\ [MnOj] = 8.0 x 10"^ mol dm-\ Temp. = 30 °C, X^,^ = 390 nm

0,0 10.0 15.0 20.0 25.0

30.0

5.1 5.7 6.1 6.4 6.7

7.2

10.5 11.5 11.7 12.0 12.8

13.4

Contd.

109

Temp. / °C

[MnOj] = 8.0 X 10"^ mol d n r \ 25 4.6 7.8

[tartaric acid] = 16.0 X ID""* mo! ^^ ^•'l j ^ ' ^ dm"\ [TX-lOO] = 15.0x10"^ 4Q J 2 g 30*2 mol d m ' \ [HCIO4] = 10.0 x lO""* mol dm"'', A,,„ax = 390 nm

Parameters

EJkJ mop ' A//*/kJ m o r ' A S V J K " ' m o r '

AG*/kJ m o r '

noncatalytic

58

55 - 25

93

autocatalytic

71

69 - 75

91

110

Table 3.24

Dependence of rate constants on the factors influencing the oxidation

of citric acid by colloidal Mn02 in presence of TX-100.

Non-variables Variables lO'^ k^il lOUv.2/ S-' S-'

10 [citric acid]/ mol dm'

[MnO2] = 8.0x 10~^moldm"^ 12.0 [TX-100] = 15.0 X 10"^ mol dm-^ i^ '?

24.0 Temp, = 30 °C, Xm&x = 390 nm 23 n

32.0 40.0

20.2 22.6 24.5 24.8 25.6 26.9

51.2 54.2 55.6 56.6 57.6 60.0

10^[MnO2]/moldm -3

[citric acid] = 16.0 x 10 ' mol dm"\ [TX-100] = 15.0 x 10"^ mol dm"\ Temp. = 30 °C,

max = 390 nm

2.0 3.2 4.0 6.0 8.0 10.0

26.6 25.2 24.8 23.0 22.6 21.2

59.5 57.6 57.6 55.4 54.2 52.3

10^[HClO4]/moldm -3

[TX-100] = 15.0 x 10-^ mol dm-^ [citric acid] = 16.0 x lO"'* mol dm' [MnO2] = 8.0x 10-^moldm-^ Temp. = 30 °C, X^^^ = 390 nm

0.0 5.0 10.0 15.0 20.0 30.0

22.6 23.4 24.4 26.4 27.1 29.5

54.2 55.4 56.1 59.7 62.4 65.8

Contd.

Ill

[MnOj] = 8.0 X 10"^ mol dm"\ [citric acid] = 16.0 x 10"'' mol d m ' \ [TX-lOO] = 15.0 xlO"^ mol dm"'', X,,„ax - 390 nm

Parameters

EJkJ mol ' AfT/kJ mol"' A ^ V J K - ' mol"'

AG*/kJ mol"'

Temp. / °C

25 30 35 40

noncatalytic

60

57 - 107

89

18.3 36.2 22.6 54.2 36.4 69.4 52.1 95.9

autocatalytic

52

49 -126

88

112

n O

1.S

1.0

0.5

o /

— /

1

A

1

o v^

A

1

/^ ( J -

1

ky2

Xvl

- 0

*

1

10 20 30 40 103CTX-100D(moldm-3)

50

Fig. 3.25: Effect of TX-IOO on the pseudo-first-order rate constants.

Reaction conditions: [MnOa] = 8.0 x 10"' mol dm'^, [glycolic acid] = 96.0 x 10"

mol dm"\ [HCIO4] = 10.0 X lO"'* mol dm~^ temperature = 30 °C.

113

10 20 30

lO^CTX-IOOn l m o l d m - 3 )

to

Fig. 3.26: Effect of TX-lOO on the pseudo-first-ordcr rate constants.

Reaction conditions: [MnOi] = 8.0 x 10" mol dnr \ [lactic acid] « 2.5 x 10" niul

dm'\ fHClO.,] = 10.0 X 10"" mo! dm"\ temperature = 30 °C.

114

10 20 30

lO'CTX-lDODlmoldm-')

40 50

Fig. 3.27: Effect of TX-lOO on the pseudo-first order rate constants. Reaction

conditions: [MnOi] = 8.0 x 10' mol dm'\ [mandelic acid] = 16.0 x 10''* mo! dm"\

[HClO l = 10.0 X ID" mol dm"\ temperature = 30 °C.

115

0.50 -

(/)

> 0.25

O

10 20 30

1 0 3 C T X - 1 0 0 3 ( m o l d m - 3 )

AO

Fig. 3.28: Effect of TX-lOO on the pseudo-first order rate constants.

Reaction conditions: [Mn02] = 8.0 x 10"* mol dm"\ [malic acid] = 16.0 x 10""

mol dm"'', temperature = 30 °C.

I {A

3-

O

0.0 10 20 30 40

10^ C TX-IOOD (moldm-3)

50

Fig. 3.29: I'lTccl ol'TX-100 on the pscuclo-first order rate constants.

Reaction iondilions: fMiiOj| • 8.0 x 10 ^ nuil dm •\ |Uirl;iric acidj - 16.0 ,v 10 '

mol dm'-'. fl-IClO.,] = 10.0 x 10" mol dm'^ temperature = 30 °C.

117

10 20 30

I O ^ C T X - I O O D ( m o l d m - 3 )

40 50

Fig, 3.30: Effect of TX-lOO on the pseudo-first order rate constants.

Reaction conditions: [Mn02] = 8.0 x 10" mol dm" , [citric acid] = 16.0 x 10"'' inol dm" , temperature = 30 °C.

118

(= 15.0 X 10" mol dm ^). The k^ - values, obtained as a function of different

variables, are summarized in Tables 3.19 - 3.24. We can see that the same

pattern is being followed as regards the variations in [Mn02], [reductant] and

[HCIO4]. The effect of varying temperature on the rate was also the same,

namely, the rate constants increased with increase in temperature. These

observations establish that the same mechanism is being followed in both the

media. The activation parameters {E^, Aff, Alf and A(f), evaluated for the

micellar media (Tables 3.19 - 3.24), show that the micelles act as catalyst

providing a new reaction path with lower Ea (aqueous > TX-lOO). A complete

account of all the factors that influence Aff, ZLS* and A(f is not possible

because the rate constant ( obs or kyy) does not represent a single elementary step, as

it is a complex function of adsorption and other constants (see Scheme 1).

The observed effect of TX-lOO on k^ (Tables 3.13 - 3.18 and Figs. 3.25

- 3.30) are catalytic up to certain [TX-lOO] and then remained constant (except

in case of malic and citric acids). In order to explain the micellar catalytic

effect on the oxidations, an attempt has been made to employ the mathematical

model proposed by Tuncay etal. '* (Eqs. (3.15) and (3.16))

log s>i = X, log [TX-lOO] - y, (3.15)

log k^2 = X2 log [TX-100] - y2 (3.16)

119

where k^y] and k^2 are pseudo-first-order rate constants, respectively, for the non-

catalytic and autocatalytic pathways (x and y being empirical constants). The above

equations predict that plots of log k^ versus log [TX-lOO] should be linear with

slopes (xi and X2) and intercepts (yi and y2). When the data were fitted, clear linear

plots resulted, Figs. 3.31 - 3.36, indicating that the model is adequate to explain the

present reactions. The values of empirical constants, obtained from Figs. 3.31-3.36

plots, are given in Table 3.25.

As the catalytic - effect curves are of Michaelis-Menten type^^ of enzyme

catalysis, the use of Lineweaver-Burk^^ method of rearranged equations (3.17)

and (3.18) were found correct.

1 b, = a, + (3.17)

(^v,i-^obsi) [TX-lOO]

= 32 + (3.18) (^v,2-^obs2) [TX-lOO]

The fulfilment of the above relafions can be seen in Figs. 3.37 - 3.42. The

values of parameters ai(a2) and bi(b2), evaluated from the intercepts and slopes

of Figs. 3.37 - 3.42, are given in Table 3.26.

120

-logCTX-1003

Fig. 3.31: Log-log plots between ,,,and fTX-lOO].

Reaction conditions: [MnO:] = 8.0 x 10" mol dm""', [glycolic acid] = 96.0 x 10"

mol dm'", [IICIO4] = 10.0 X lO""* mol dm"\ temperature = 30 °C. The data belong

to the part up to which the effect of TX-IOO was catalytic (cf. Fig. 3.25).

121

o I

l o g C T X - l O O D

Fig. 3.32: I.og-Iog plots between A',|,and [TX-lOO].

Reaction conditions: [Mn02] = 8.0 x 10 mo! dm , [lactic acid] = 2.5 x 10' mol

dm'\ [HCIO4] = 10.0 X lO'"* mol dm'\ temperature = 30 ''C. The data belong to the

part up to which the effect of TX-lOO was catalytic (cf. Fig. 3.26).

122

C7)

o

- l o g C T X - 1 0 0 3

Fig. 3.33: Log-log plots between Arv and [TX-lOO].

Reaction conditions: [MnO:] = 8.0 x 10" mol dm"\ [mandelic acid] = 16.0 x 10"

mol dm"\ [HC10.(] = 10.0 x 10" mol dm~\ temperature = 30 "C. The data belong

to the part upto which the effect of TX-lOO was catalytic (cf Fig. 3.27).

123

o

- log C T X - 1 0 0 3

Fig. 3.34: Log-log plots between /:,,, and [TX-lOO].

Reaction conditions: [malic acid] = 16.0 x 10"* mol dm"\ [Mn02] = 8.0 x 10"'' mol

dm"\ temperature = 30 °C. The data belong to the part upto which the effect of TX-

100 was catalytic, (cf. Fig. 3.28).

124

o

log CTX-100D

Fiii. 3.35: f.oj'-lo)- plols hclWLvn />,,;iiul \TX-\()0].

Reaction conditions: [MnOi] = 8.0 x 10" mol dm"'', [tartaric acid] = 16.0 x lO""*

mol din'\ [HCIO4] = 10.0 x lO"* mo! dm"\ temperature = 30 "C. The data belong

to the part upto, which the effect of TX-lOO was catalytic (cf. Fig. 3.29).

125

2.2

>

cr» 2.6 O

T • " " ^ • ^

.^__J<y2

'-^^^^kyi

3.0 J_

- logCTX-lOOD

Fig. 3.36: Log-log plots between k^^, and [TX-lOO].

Reaction conditions: [Mn02] = 8.0 x 10"* mol dm""', [cilric ticldj ^- 16.0 x 10 ' inol

dm"\ temperature = 30 ^C. The data belong to the part up to which the cITccl of

TX-lOO was catalytic (cf. Fig. 3.30).

126

Table 3.25

Values of the empirical constants of Eqs. (3.15) and (3.16) for the oxidation of

organic acids by colloidal MnOi (= 8.0 x 10" mol dm" ) in presence of TX-lOO

at 30 °C.

Acid xi yi X2 yi

Glycolic acid , , 0.0822 2.9192 0.1666 2.5166

(96.0 X 10" mol dm" )

Lactic acid 0.3348 2.7455 0.2210 2.7455

(2.5xl0-^moldm~^)

Mandelicacid 0.1600 2.6221 0.1138 2.4300

(16.0 x 10- moldm" )

DL-Malicacid 0.1197 3.1385 0.0653 3.1597

(16.0xl0"'moldm"^)

DL-Tartaric acid 0.2035 2.8737 0.0714 2.8075

(16.0 X I0~^moldm~ )

Citric acid 0.2146 2.2831 0.0779 2.1299

(16.0xl0-^moldm-^)

127

0.1 0.2

10~3/c TX-1003(morldm3)

Fig.3.37: \I{K,, - ohs) against 1/[TX-100] plots.

Reaction cmulilinns: fMnOi] = 8.0 x 10"' mol dm"\ [glycolic acid] = 96.0 x \{T^

mol Cim'^, |I1C10.|] = 10.0 x 10"" mol dm"\ temperature = 30 °C. The data belong

to the p:irl up to which the effect of TX-lOO wjis catalytic (ef. Fig. 3.2.S).

128

10" 3 / C T X - I O O D (mol ~''dm^)

I'l^. .\JH: I/(A,,, - A „„.J iigiiiiiM l/| I X-|{)()) |)K)l.s.

Rciiclion co/ic/i/iu/is: [Mn02j - 8.0 x 10"' mol dnr \ [lactic acidj = 2.5 x 10 ' mol

dnr\ [HCIO.,] - 10.0 X lO:"* mo! dnr \ lempcralure = 30 °C. The data belong to the

part up to which the effect of TX-lOO was catalytic (of. Fig. 3.26).

129

10" V C T X - I O O D l m o r ' ' d m 3 )

-4

Fig. 3.39: \/{k^ - /tobs) against 1/[TX-100] plots.

Reliction conditions: [MnOij = 8.0 x 10' mol dm"'', [mandelic acid] = 16.0 x 10

mo! dm'\ [HCIO4] = 10.0 x 10"" mol dm~\ temperature = 30 °C. The data belong

10 ihc pari upio which the effect of TX-lOO was catalytic (cf. Fig. 3.27).

130

0.10 0.20

1 0 " 3 / l l T X - 1 0 0 l l ( m o r * ' d m 3 )

Fig. 3.40: l/(/:^ - kobs) against 1/[TX-100] plots.

Reaction conditions: [malic acid] = 16.0 x 10" mol dm"\ [Mn02] = 8.0 x 10"

mol dm~\ temperatiire = 30 °C. The data belong to the part up to which the effect

of TX-lOO was catalytic (cf. Fig. 3.28).

131

10.0 -

0 0.1 0.2 l 6 3 / C T X - 1 0 0 3 ( m o r ^ d m 3 )

Fig. 3.41: \/{k^ - k^bs) against 1/[TX-I00] plots.

Reaction condliiom: [MnOj] = 8.0 x 10"* mol dm"\ [tartaric acid] = 16.0 x lO""

mol dm-\ [HCIO4] = 10.0 x 10" mol dm-\ temperature = 30 °C. TJie data belong

to the part up to, which the effect of TX-lOO was catalytic (cf Fig. 3.29).

132

0 0.5

10-3/ C T X - I O O D (mol"^dm3 )

Fig. 3.42: l/(Jt - ;tobs) against 1/[TX-100] plots.

Reaction conditions: [Mn02] = 8.0 x 10" mol dm"\ [citric acid] = 16.0 x 10"" mol

dm'\ [MnOz] = 8.0 x 10"* mol dm"\ temperature = 30 °C. The data belong to the

part up to which the effect of TX-lOO was catalytic (cf. Fig. 3.30).

133

Table 3.26

Values of the empirical constants of Eqs. (3.17) and (3.18) for the oxidation of

organic acids by colloidal MnOi (= 8.0 x 10" mol dm" ) in presence of TX-lOO

at 30 °C.

Acid 10"' ai/ bi/ lO'^az/ bz/

mol dm' s s mol dm' s

Glycolic acid

(96.0 X 10- mol dm~ ) 8.30 119.0 11.30 7.0

Lactic acid 20.34

(2.5 X 10" mol dm" )

107.7 74.89 199.7

Mandelic acid 45.60

(16.0xl0"^moldm~^)

38.9 9.30 18.1

DL-Malic acid 0.33

(16.0xl0"*moldm"^)

4452.1 1.34 6450.8

DL-Tartaric acid 23.15

(16.0xl0"'moldm-^)

14.1 25.21 37.9

Citric acid 7.26 -3> , (16.0xl0"^moldm~0

3.4 4.89 0.5

134

Probable Role ofTX-100

Adsorption of non-ionic surfactants and gum arabic (a protective colloid) on

the surface of the colloidal particles is a well known phenomenon. As a result, gum

arabic stabilizes the colloidal Mn02^ whereas non-ionic surfactants enhance the

dispersion stability. ^ Therefore, both should have the same influence on the reaction

rate. As pointed out earlier, the effects of the two are opposite to one another, i.e., TX-

100 had a catalytic but gum arabic an inhibitory effect. These results indicate that

adsorption is not the only factor responsible to explain the catalytic effect observed

with TX - 100 on the reduction of Mn02 by organic acids. In addition, the role of

other factors, e.g., hydrogen bonding, properties of interfacial water (known to be less

polar but more structured than bulk water), differences in stabilization of the initial

and transition states by surfactant molecules, reaction rates in the bulk and micellar

pseudo phase, etc., cannot be ruled out completely.

As far as the role of TX-lOO is concerned, hydrogen bonding between

the non-ionic TX-lOO (I) and the reactants may play an important role. The

organic acids contain no hydrophobic character (except mandelic acid) and

CH3 CH3

C H 3 - C - C H 2 - C - < ^ f O C H 2 C H 2 ) ^ O H

CH3 CH3

TX-lOO (I)

have OH and COOH groups. Hydrogen bonding may occur between these groups

of the acids (reductants) and the ether-oxygen of the polyoxyethylene chains of TX-

135

100. Due to the presence of a number of donor groups in one TX-lOO molecule,

multiple H-bonding may take place and the number of bound reductant molecules

increase. In this surfactant, the lengths of the hydrophobic and hydrophilic parts are

comparable and have significant amount of water in the outer shell. The hydrogen

bonding between Mn02 sols and hydrophilic part (polar ethylene oxide) of the TX-

100 can not be ruled out either (let us call it adsorption!). Therefore, the associated

Mn02 and reductant (glycolic, lactic, mandelic, malic, tartaric, citric acids) with TX-

100 (through hydrogen bonding) seems responsible of facilitating the reaction; this

might be the role of TX-lOO towards the observed catalysis. This surfactant thus

helps in bringing the reactants closer, which may orient in a manner suitable for the

redox reactions followed by rearrangement of TX-lOO molecules.

Finally, to confirm Tuncay's propositions "* of hydroxyl ions bonded to

colloidal Mn02 being the active points for substrate adsorption on the colloidal

surface, the reactivity of different reductants towards colloidal Mn02 has been

compared (Table 3.27). Based on their reactivity, the reductants can be ordered as:

citric > mandelic > tartaric > malic > lactic > glycolic. As four hydroj 'l groups

(three from -COOH groups and one from -OH) are contained in each citric acid

molecule, reaction rate is expectedly higher in comparison to others (Table 3.27).

Mandelic acid rate being greater as compared to tartaric and malic acids seems to be

due to hydrophobicity of-CeHs group (phenyl group) present in mandelic acid, i.e.,

presence of a hydrophobic moiety in the acid influences the rate as well.

136

Table 3.27

Comparison of second-order-rate constants (A:") for the oxidation of organic

acids by colloidal MnOa at 25 °C.

Acid Medium l O V / m o r ' dm^s

Glycolic acid^

Lactic acid''

Mandelic acid'

DL-Malic acid**

DL-Tartaric acid*

Citric acid

10.0 X 10~*moldm-'HClO4

10.0 X 10"'moldm"^HClO4

10.0 X 10~''moldm"^HClO4

aqueous neutral

10.0 X 10~'*moldm~^HClO4

aqueous neutral

0.6

0.8

36.2

13.1

14.4

38.2

''[Mn02] = 8.0 X 10"^ mol dm"^ [glycolic acid] = 96.0 x 10"^ mol dm'^

''[MnOz] = 8.0 x 10"^ mol d m ' \ [lactic acid] = 2.5 x 10"^ mol dm"\

'[MnOz] = 8.0 X 10"^ mol dm"^ [mandelic acid] = 16.0 x lO"'* mol dm"^

"[MnOz] = 8.0 X 10"' mol dm"\ [DL- malic acid] = 16.0 x 10"^ mol dm"^

'[MnOj] = 8.0 x 10"^ mol dm~\ [DL-tartaric acid] = 16.0 x 10"'* mol dm"^

'•[MnOj] = 8.0 X 10'^ mol dm'^ [citric acid] = 16.0 x lO"'* mol dm"l

137

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