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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 ,
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T7018
HO
PROF. KABIR-UD-DIN CHAIRMAN DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY ALIGARH-202002 (U. P.) INDIA. E-mail: [email protected]
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
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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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