1 CHAPTER 1 INTRODUCTION 1.1 CHEMICAL KINETICS Chemical kinetics has been developed in a logical and coherent fashion over the years. However, an understanding of the way we achieved our present ideas on chemical kinetics is a very good basis to understand the subject. The great success of Newtonian mechanics in the area of mechanics and astronomy, which involved the idea of explaining phenomena by simple forces acting between particles. This led scientists in the nineteenth century to introduce such mechanical explanation to all areas involving natural phenomenon. In chemistry, for example, these concepts are applied to interpret chemical affinity or chemical mechanics. Hence understanding of chemical kinetics within this context is important. The concept of chemical kinetics evolved relatively late in terms of the overall studies of reactions and reactivity. The study of chemical kinetics can be traced back to Wilhemy (1850), who carried out the first study of inversion of cane sugar in the presence of acid. He formulated in terms of a first order mathematical expression to interpret the progress of the reaction. Unfortunately, this work went unrecognized until Ostwald (1884) drew attention 34 years later. It may seem strange today that such an idea of studying the variation of chemical affinity with time had not occurred earlier. Farber (1961) had tried to explain this and shown that there were some earlier attempts to study the time evolution of reaction even before Wilhemy (1850).
Transcript
1
CHAPTER 1
INTRODUCTION
1.1 CHEMICAL KINETICS
Chemical kinetics has been developed in a logical and coherent
fashion over the years. However, an understanding of the way we achieved
our present ideas on chemical kinetics is a very good basis to understand the
subject. The great success of Newtonian mechanics in the area of mechanics
and astronomy, which involved the idea of explaining phenomena by simple
forces acting between particles. This led scientists in the nineteenth century to
introduce such mechanical explanation to all areas involving natural
phenomenon. In chemistry, for example, these concepts are applied to
interpret chemical affinity or chemical mechanics. Hence understanding of
chemical kinetics within this context is important.
The concept of chemical kinetics evolved relatively late in terms of
the overall studies of reactions and reactivity. The study of chemical kinetics
can be traced back to Wilhemy (1850), who carried out the first study of
inversion of cane sugar in the presence of acid. He formulated in terms of a
first order mathematical expression to interpret the progress of the reaction.
Unfortunately, this work went unrecognized until Ostwald (1884) drew
attention 34 years later. It may seem strange today that such an idea of
studying the variation of chemical affinity with time had not occurred earlier.
Farber (1961) had tried to explain this and shown that there were some earlier
attempts to study the time evolution of reaction even before Wilhemy (1850).
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The practical importance of such studies did not exist at the end of
eighteenth century, and it was only with the advent of chemical industry at the
beginning of the nineteenth century that chemists need to consider this
problem. Eventually, this became important for the development of industrial
research at the end of that century. An excellent discussion of this problem is
given by King (1982 and 1984), King and Laidler (1984) in their studies on
the history of chemical kinetics where they analyzed the impact of various
theoretical, experimental and conceptual works of Berthelot and Giles (1862)
and Harcourt and Essen (1866). These researchers are truly considered to be
the founder of this new branch of chemistry called chemical kinetics.
The formation of activated complex during the reaction was
proposed and reported by Arrhenius (1889) in which chemical kinetics was
slow to develop because complex phenomenon involved when molecules
collide to produce chemical reactions. He interpreted the effect of temperature
on reaction rates of a chemical reaction. For rates measured under standard
concentration conditions, Arrhenius expressed this effect by equation (1.1)
k = Ae (Ea/RT) (1.1)
where k is the rate constant under standard conditions and A and Ea are
constants, which are practically independent of temperature. A is called the
frequency factor and Ea is the activation energy. The Arrhenius law took a
long time to become accepted, and from this many other expressions were
derived and proposed to explain the dependence of rate on temperature.
Many of the conceptual and experimental difficulties would
disappear with the work of Vant’s Hoff, who introduced the concept of order
of reaction and proposed the mechanism of a chemical reaction based on the
basis of chemical kinetics. In the first phase of development, the theories of
chemical kinetics tried to resolve the problem of calculation of
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pre-exponential factor and activation energy in the Arrhenius equation. Lewis
(1918) developed mathematical expressions that allowed the formulation of
collision theory for pre-exponential factor. Henry Eyring (Laidler and King
1983) came forward to develop the transition state theory based on
thermodynamics and statistical mechanics.
The energy differences between reactants and products in solution
are normally measured in terms of their equilibrium constants. The
equilibrium constant of a reaction is related to the energy as given in
equation (1.2)
G0 = RT ln Keq (1.2)
where G0 is the standard free energy, Keq is the equilibrium constant, R is
the gas constant and T is the temperature in absolute. However, the science of
kinetics does not end here. The next task is to look at the chemical steps
involved in a chemical reaction and develop a mechanism. Chemical kinetics
is not just an aspect of physical chemistry but it is a unifying topic covering
the whole of chemistry, many aspects of biochemistry and pharmaceutical
industries. In this context it covers the measurement of rates of reaction and
analysis of experimental data to give a systematic collection of information
which summarises all quantitative kinetic information of any given reaction.
This, in turn, enables comparison of reactions. The sort of information used
here is summarised in terms of the
factors influencing rates of a reaction,
dependence of the rate of reaction on concentration called the
order of the reaction,
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rate expression, which is an equation which summarizes the
dependence of the rate on the concentration of substances
which affect the rate of reaction,
expression involves the rate constants which is a constant of
proportionality linking the rate with the various concentration
terms,
rate constant collects in one quantity all the information
needed to calculate the rate under specific conditions and
effect of temperature on the rate of reaction. Increase in
temperature generally increases the rate of a reaction. It gives
information to a deeper understanding of a reaction to occur.
1.2 CATALYSIS
The term catalysis was introduced by Berzelius (1835) in order to
explain various decomposition and transformation reactions. He assumed that
catalysts possess special powers that can influence the affinity of chemical
substance. In continuation, Ostwald offered a valid definition for catalysis in a
book of Gates (1992). A catalyst is a substance which alters the speed of a
chemical reaction without itself undergoing any chemical change and the
phenomenon is known as catalysis. Example
MnO2 2KClO3 2KCl + 3O2 (1.3)
The suitable catalyst for an industrial process mainly depends on the activity,
selectivity and stability (deactivation behavior).
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1.2.1 Types of Catalytic Reactions
Catalytic reactions are classified into two broad types, viz.,
1. Homogeneous catalysis
2. Heterogeneous catalysis
1.2.1.1 Homogeneous catalysis
In this reaction, the reactants and catalyst are in the same phase.Homogeneous catalysts have a higher degree of dispersion thanheterogeneous catalysts since in theory each individual atom can becatalytically active. Due to their high degree of dispersion, homogeneouscatalysts exhibit higher activity than heterogeneous catalysts. The highmobility of the molecules in the reaction mixture gives more collisions withsubstrate molecule. The reactants can approach the catalytically active centerfrom any direction and a reaction at an active center does not block theneighboring centre. This allows the use of low catalyst concentration andmilder reaction conditions. Homogeneous catalyzed reactions are controlledmainly by kinetics and less by material transport because diffusion of thereactants to the catalyst. Hence the mechanism of homogeneous catalysis isrelatively clear and mechanistic investigations can readily be carried outunder reaction conditions by means of spectroscopic methods.
Metal ions, being Lewis acids, can be expected to assist a reactionby drawing the electrons of a bond to themselves and thereby weakening thebond to be broken. Though a proton can do this, metal ions which carry morethan a single positive charge are much more efficient than H+. Thus metalions can function as super acids. For example, polyvalent metal ion such asCu2+, Fe2+, Fe3+, etc., catalyze decarboxylation of -keto acids like oxolacticacid. In this reaction, transition metal ions are found to be better catalyst thannon-transition metal ion like Al3+ because of their ability to coordinate withreactant molecules.
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Transition metal ions and their complexes catalyze a wide variety
of reactions like hydrogenation, oxidation and polymerization. Apart from
their ability to coordinate with the substrate molecules, the availability of
vacant d-orbitals enables the formation of sigma bonds with ligands which are
Lewis bases. The three t2g orbitals remain non-bonding. However, these non-
bonding electrons on the metal ion can also be donated to suitable -acceptor
ligands. This ability to form both sigma and pi-bonds with suitable substrates
constitutes one of the important factors in catalysis by these ions. Their ability
to assume a variety of oxidation states and different coordination numbers
enable them to act as catalysts in redox reactions. The original objective of
binding metal complexes to insoluble supports is to get over the problem of
separation of homogeneous catalyst from the reaction mixture. A number of
materials have been used as supports for complexes viz., silica, -alumina,
molecular sieves (zeolites) and clays.
1.2.1.2 Heterogeneous Catalysis
The reactants and catalyst (usually solid) are in the different phase
in heterogeneous catalysis. Heterogeneous catalysts provide a surface for the
chemical reaction to take place. In order for the reaction to occur, one or more
of the reactants must diffuse to the catalyst surface and adsorb onto it. After
reaction, the products must desorb from the surface and diffuse away from the
solid surface. Frequently, this transport of reactants and products from one
phase to another plays a dominant role in limiting the reaction rate.
Understanding these transport phenomenon and surface chemistry such as
dispersion is an important area of heterogeneous catalyst research. If diffusion
rates are not taken into account, the reaction rate for various reactions on
surface depends solely on the rate constant and reactant concentration. The
solid catalysts that are used in major industries for various purposes include
food industry (hydrogenation and isomerisation of olefins), fine chemicals
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(catalytic hydrogenation), petroleum industry (catalytic cracking and
reforming process) and heavy inorganic chemicals (ammonia synthesis and
nitric acid). Some of the solid catalysts are clays, zeolites, mesoporous
molecular sieves and heteroploy acids.
1.3 OXIDATION
The combination of processes involved in the flow of electrons
from a reducing agent (reductant) to an oxidizing agent (oxidant) is oxidation.
The total number of electrons lost by one substance is the same as the total
number of electrons gained by another substance. Oxidation and reduction
always occur simultaneously and are really opposite sides of the same
reaction, which is often called redox reaction. Oxidation involves the loss of
electrons from the reducing agent. Since electrons carry negative charge,
oxidation results in an increase of positive valence. Therefore oxidizing agent
is the substance that brings the oxidation.
Oxidation state of an atom in a chemical compound is counted by
set of rules: (1) the oxidation state of a free element (uncombined element) is
zero, (2) for a simple (monoatomic) ion, the oxidation state is equal to the net
charge on the ion and (3) hydrogen has an oxidation state of +1 and oxygen
has an oxidation state of 2 when they are present in compounds. Exceptions
to this are that hydrogen has an oxidation state of 1 in hydrides of active
metals, e.g. LiH and oxygen has an oxidation state of 1 in peroxides, e.g.
H2O2. The algebraic sum of the oxidation states of all atoms in a neutral
molecule must be zero, while in ions the algebraic sum of the oxidation states
of the constituent atoms must be equal to the charge on the ion. For example,
the oxidation states of sulfur in H2S, SO2 and H2SO4 are 2, +4 and +6
respectively. Hence higher the oxidation state of a given atom, greater is its
degree of oxidation and lower is the oxidation state, greater is its degree of
reduction.
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1.3.1 Importance of Oxidation
Oxidation reactions play an important role in organic synthesis and
there is currently demand for selective and efficient oxidation methods
(Backwall 2004). Pressure from society has placed restrictions on industrial
oxidation technology because of the need for sustainable and ecofriendly
processes. Today there is an increasing demand to use oxidants such as
molecular oxygen and H2O2 which are environmental friendly and do not give
rise to any waste products (Sheldon et al 2002).
The direct oxidation of organic substrates by either O2 or H2O2 is
rare as the energy barrier for electron transfer from the organic substrate to the
oxidant is usually high. This high energy barrier is the way nature protecting
organic compounds from destructive oxidation. Nature has also found
methods to make controlled aerobic oxidation under highly mild condition.
The unfavourable kinetics associated with direct aerobic oxidation is in the
respiratory chain, which is involved in many biological oxidations (Duester
1996, Gille and Nohl 2000, Berkessel 2006).
Oxidation reactions are of fundamental importance in nature and
are key transformations in organic synthesis (Julio and Backwall 2008). The
developments of new processes employ transition metals as substrate-
selective catalysts and stoichiometric friendly oxidants. Direct oxidation of
the catalyst by molecular oxygen or hydrogen peroxide is often kinetically
unfavored. Selective oxidation reactions based on palladium catalyst led to
excitement in 1950s since the direct reoxidation of palladium by molecular
oxygen is difficult in many cases (Stahl 2005). Therefore Pd catalyst is now
used in the fuel cells and biological enzymes for selective oxidation. Hence
metal catalyzed oxidation has been providing a better alternative route and
they provide new opportunities for industrially relevant reactions. Improved
oxidation methods also provide inventive methods to improve the substrate
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conversion and product selectivity by increasing the solubility of oxidizing
agent in the reaction mixture. In modern organic synthesis, oxidation of
organic substrates has been carried out in aprotic solvents under mild and
neutral conditions to get maximum product.
1.3.2 Oxidation of Amino acids
Oxidation of amino acids is one of the most prevalent forms of
chemical reactions and is susceptible to modification by a wide array of
oxidants. Uncharacteristic oxidation reactions are of particular concern in
biotechnology and medicine. Therefore, it is important to understand the
amino acid metabolism (Bender 1975) with a model system or model reaction
towards oxidants.
Amino acids have both amino and carboxylic acid functional
groups and therefore they are both acids and bases. In certain compound
specific pH known as isoelectric point, the number of protonated amino
groups with a positive charge and deprotonated carboxylate groups with a
negative charge are equal, resulting a net neutral charge. These ions are
known as zwitter ions. Amino acids are zwitter ions in solid phase and in
polar solutions such as water depending on the pH but not in the gas phase.
Zwitter ions have minimal solubility at their isoelectric point. Simple
amino acids present in municipal wastewater can cause serious eutrophication
in water bodies. It is essential to remove them from wastewater. In textile
industry and tanneries the waste sent out to nearby area contain some simple
amino acids. It rises up the nutrients on that particular place. This may not be
suitable for living plants. It affects their growth and population. Therefore,
before discharge, they may be treated completely and converted into
ecofriendly product.
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The importance of biologically active pyridoxal phosphate
(Vitamin B6) in the metabolism of amino acids has been well established. It is
the cofactor for enzymes catalyzing a number of reactions of amino acids
such as transfer of amino group to -ketoacids (Braunstein 1973, Metzler
1979, Walsh 1979). The initial reaction of an amino acid with pyridoxal
phosphate is the formation of an intermediate complex, a Schiff’s base, by
condensation between -amino group of the substrate and 4'-formyl group of
pyridoxal phosphate (Braunstein and Shemyakin 1953, Metzler et al 1954,
Metzler 1957, Braunstein et al 1960). The electron withdrawing effect of
heterocyclic nitrogen results withdrawal of electrons from all the three bonds
around the -carbon. The structural requirements for non-enzymic and
coenzyme activity of pyridoxal phosphate analogues are the heterocyclic
nitrogen for electron withdrawal, 4'-formyl group and phenolic hydroxyl
group (Groman et al 1972). Hence in the non-enzymatic model reactions,
heterocyclic nitrogen can be placed by an oxidant which will act as an
electron withdrawer.
It was observed that thermal decomposition of -amino acids in the
presence of aldehydes and ketones is relatively fast (Chatelus 1964). It was
also noted that in the oxidation of -amino acids, the product aldehyde
enhanced the rate of the reaction. Therefore a chemical model pyridoxal-
catalyzed amino acid metabolism can also be constructed using formaldehyde
and oxidant. The efficacy of the model can be measured from the oxidation of
amino acids. However, this study brings out the importance of Schiff’s bases
(aldimines) in the oxidative decarboxylation-deamination of amino acids.
1.3.3 Oxidation of Organic Substrates with Oxidizing Agents and
Metal Ions
Kinetic study of dioxygen complexes has been reported by Martell
(1983). Reaction scheme is described in which oxygen complexes of cobalt,
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iron and copper may be involved in the oxidation of organic substrates by
electron transfer (oxidase models) and oxygen insertion (oxygenase models).
Many non-metal redox reactions take place by atom transfer or ion transfer
processes. However, much less attention has been paid to electron transfer
reactions. A detailed study of the electron transfer reactions was investigated
by Kumar and Margerum (1987). Based on the study, the following rate
expression was proposed for the oxidation of bromide by HOCl and OCl
THA T
--d[OCl ] - - = k [HA] [OCl ] [Br ]dt
(1.4)
where [OCl ]T = [OCl ] + [HOCl] and HA is the general acid.
The kinetics of oxidation of thiosemicarbazide, thiocarbohydrazide
and hydrazone by hydrogen peroxide in both perchloric and sulphuric acid
media were investigated by Thimme Gowda and Vasi Reddy (1990). The
rates revealed first order kinetics each with respect to [oxidant] and [substrate]
in all the cases.
In aqueous acid solution, the complex [Ag(H2L)]3+ (H2L= ethylene
bis (bigunaide)) quantitatively oxidized ethanol and isopropyl alcohol to the
corresponding carbonyl products at moderate rate (Banerjee et al 1990). The
rate law is expressed as in equation (1.5).
k-d[complex] 2 = k + [alcohol] [complex]+1dt [H ] (1.5)
During the reaction, deprotonation of the alcohols assisted by axial
coordination to [Ag(H2L)]3+ is suggested to be the source for inverse acid
dependence.
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The kinetics of oxidation of propionic acid by Ce(IV) in aqueous
perchloric acid was studied by Alvarez-Macho and Mata-Perez (1992). The
analysis revealed that the rate of oxidation is proportional to propionic acid
concentration and the over all order is two. The activated complex consists of
different species such as propionic acid, Ce(IV), H2O and H+. Amels et al
(1997) reported the kinetics of oxidation of dimethyl sulfide (DMS) by
hydroperoxides such as hydrogen peroxide, peroxo formic acid, peroxo acetic
acid, peroxo nitrous acid and peroxomonosulfuric acid anion.
The kinetics of oxidation of mandelic acid by permanganate in
aqueous alkaline medium was studied spectrophotometrically (Rafeek et al
1998). The reaction showed first order kinetics over [permanganate ion] and
fractional order dependence over [mandelic acid] and [alkali]. There was no
significant change on the reaction rate upon the addition of aldehyde to the
product. An increase in the ionic strength and decrease in dielectric constant
of the medium increased the rate. The oxidation process in alkaline medium
under the conditions employed in the present investigation proceeded first by
formation of an alkali permanganate complex, which combined with mandelic
acid to form another complex. The latter decomposed slowly followed by fast
reaction between free radical of mandelic acid and another molecule of
permanganate to give the final products. There is a good agreement between
the observed and calculated rate constants under different experimental
conditions.
The kinetics of oxidation of glyoxylic acid by
peroxomonophosphoric acid (PMPA) in acidic medium was reported by Vijai
et al (2008). The stoichiometry corresponds to the reaction of one mole of
PMPA and one mole of glyoxylic acid. The reaction was found to be second
order.
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CHO
COOH + H3PO5 HCOOOH + CO2 + H3PO4 (1.6)
Ag(I) catalyzed cold aqueous persulphate oxidation of primary
amines are converted into aldehydes or ketones. When applied to secondary
amines of the types (RCH2)2 NH and (R1R2CH)2 NH gave poor yields in
which -amino acids are converted into aldehyde. But the yields are not better
than with other inorganic oxidants (Bacon et al 1966).
(RCH2)2NH + S2O82 RCHO + RCH2.NH2
+ + H+ + 2 SO42 (1.7)
The reaction between malic acid and potassium peroxydisulphate in
the absence of a catalyst was found to be very slow. However, the rate was
enhanced by the introduction of Ag(I) ion as a catalyst (Agrawal et al 1970).
The total order of the reaction was nearly unity. The observed reaction and
rate equation are expressed in equations (1.8) and (1.9) respectively.
Ag+ + S2O82 Ag2+ + 2 SO4
2 (1.8)
22 0.3 +2 8
2 8-d[S O ] = k [S O ] [malic acid] [Ag ]
dt (1.9)
The kinetics of Ag(I) catalyzed oxidation of benzoin and four
substituted benzoins by peroxydisulphate in acid-water mixture was
investigated by Khandual (1990). The rate law was explained on the basis of
free radical mechanism using steady state principle
2+ 22 8
2 8
-d[S O ] = k [Ag ][S O ]dt
(1.10)
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Maleic anhydride, a precursor for polyester resins, is made by the
oxidation of n-butane over vanadium phosphate catalyst (Coulston et al
1997). The rate of malic anhydride formation is proportional to the rate of
decay of V5+ species in the catalyst. Thus V5+ species are kinetically
significant for the production of malic anhydride and not just for the
production of by-products. The results also suggested that V5+ species may
play a role in the initial hydrogen abstraction from n-butane, the
rate-determining step in the reaction sequence. V4+ sites appeared to be
responsible for the by-product formation.
The oxidation of 6-aminocaproic acid by Ag(III) complex
(dihydroxydiperiodatoargentate(III)) was studied in alkaline medium (Huo
et al 2007). The reaction was first order with respect to Ag(III) complex and
the order with respect to 6-aminocaproic acid was found to be from one to
two. A plausible mechanism involving a pre-equilibrium adduct formation
between the complex and reductant was proposed from the kinetic study.
1.4 LITERATURE REVIEW ON THE OXIDATION OF AMINO
ACIDS WITH OXIDIZING AGENTS
Toennies and Homiller (1942) reported the action of hydrogen
peroxide dissolved in formic acid oxidation of tryptophan, methionine and
cystine. The rates of Caro’s acid oxidation of dimethylaniline and
p-substituted dimethylaniline were measured in aqueous solution at various
pH and temperatures. The rates were found to be proportional to the product
of concentration of both the reactants. The effects of pH and
p-substituents on the rate were estimated. The rate constant in acid medium
increased with increasing electron withdrawing power of the substituents of
dimethylaniline (Ogata and Tabushu 1958).
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Amino acid oxidation was carried out with a variety of oxidizing
agents to produce a variety of products. Silver oxide oxidation produces
carboxylic acid with one carbon less (Clarke et al 1968). The kinetics of
oxidation of glycine and valine by chloramine-T in hydrochloric acid medium
was studied. The rate of disappearance of chloramine-T showed first order
dependence on both chloramine-T and amino acid and an inverse first order
with respect to [H+] (Gowda and Mahadevappa 1979). The kinetics of
oxidation of glycine, alanine, phenyl alanine, serine, threonine, aspartic acid
and glutamic acid by permanganate in aqueous medium were investigated and
proposed suitable mechanism for the reactions (Surender Rao et al 1979). The
rate law was found to be
0-d[Mn(VII)] = k + [amino acid] [Mn(VII)]
dt (1.11)
Laloo and Mahanthi (1990) studied the kinetics of oxidation of
glutamic acid and aspartic acid by alkaline hexacyanoferrate (III). The rate
depended on the concentration of substrate and oxidant but independent of the
concentration of alkali. During the reaction, there is loss of hydrogen atom
from the C-H bond in a slow step, giving a radical species which was
characterized by ESR spectroscopy. The reaction proceeds with formation of
-imino acid in a fast step, which undergoes hydrolysis to give the
corresponding -keto acid.
Kinetics of oxidation of twelve different -amino acids by
N-chlorosuccinimide (NCS) in aqueous alkaline media were studied and
compared with those of N-bromosuccinimide (NBS) oxidation
(Ramachandran et al 1990). Perusal of the results showed that NCS/NBS
reacted with -amino acid anion to produce -amino acylhypohalite which
then decomposed in the rate-determining step to give the products. The
intermediate -amino acylhypohalite was identified by UV-Visible absorption
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spectra. But glycine behaved differently from other amino acids with both the
oxidants.
Aliphatic amino acids were smoothly oxidized by
peroxomonophosphoric acid (PMPA) in the pH range 6-10. Kinetic studies
revealed second order dependence on the amino acid, which was not observed
earlier (Panigrahi and Paichha 1991). However the dependence of rate on
[PMPA] was unity. The reaction path included oxidation step due to H2PO5
and HPO5 . Second order dependence on the amino acid was attributed to the
participation of two zwitter ionic molecules of the amino acid to form a
nucleophile.
Electrochemical oxidation of amino acids in the presence of
triphenylphosphine gave -amino aldehydes (Maeda et al 1992). Amino acids
with longer side chains, hydrogen abstraction could lead to the formation of
hydroxyl derivatives via oxy-radical-mediated reaction pathway which could
form alkoxyl and peroxyl radicals (Luxford et al 2002). Studies on the
modification of valine to hydroxide form of valine and leucine to hydroxy
leucine (Fu et al 1995 and Fu and Dean 1997) supported this study.
Kinetics of amino acid oxidation by sodium salt of 2-p-
phenylsulfonicacid-2-phenyl-1-piperylhydrazyl and 2,2-di-p-phenylsulfonicacid -
2-phenyl-1-piperylhydrazyl at isoelectric point of amino acids were studied in
the temperature range 298-318 K. The analysis of the results provided
information about the mode of action during the reaction. The mechanistic
pathway of amino acids oxidation occurred probably by an intermediate of
amino acid radical type which led to keto acids (Ionita et al 2000).
Oxidative decarboxylation of L-ornithine by permanganate in
sulfuric acid medium was found to be first order with respect to both oxidant
and substrate concentration (Usha and Choubey 2006). The kinetics of
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oxidation of L-phenyl alanine by diperiodatoargentate (III) (DPA) in aqueous
alkaline medium was studied spectrophotometrically (Lamani et al 2009). The
reaction showed first order in [DPA], less than unit order dependence each in
[L-phenyl alanine] and [alkali] and retarding effect of [IO4 ] under the
reaction conditions. The active species of DPA was monoperiodato
argentate (III) (MPA). The reaction is shown to proceed via MPA-L-phenyl
alanine complex, which decomposed in the rate-determining step to give
intermediates followed by a fast step to give the products. The products were
identified by spot and spectroscopic studies.
1.4.1 Oxidation of Amino acids by Oxidizing Agents in the Presence
of Metal Ions
Silver catalyzed persulphate oxidation of amino acids generates
aldehyde with one carbon less (Bacon et al 1966). The reactions are found to
be acid catalyzed, and the kinetic data indicate the participation of water
molecules in the rate-determining step as a proton abstracting agent from the
substrate as per Bunnett's hypothesis. As Ag(I) was found to catalyze these
reactions, the oxidation of glycine and glutamic acid was studied. The rate
law was found to be
k k'' [amino acid] [Ag(I)]- dln[Mn(VII)] c =dt 1+ k [amino acid] = k [Ag(I)]
(1.12)
Studies on the oxidative decarboxylation of amino acids by
peroxydisulphate in the presence of Ag(I) revealed that aldehyde is formed
through radical intermediates (Minisci et al 1983, Zelechonok and Silverman
1992). It served as the model system to understand the catalysis by
monoamine oxidase (Gates and Silverman 1990, Silverman 1992). These
results stimulated interest to study the kinetics and mechanism of oxidative
decarboxylation of amino acids by peroxydisulphate and also in the presence
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of various metal ions (Chandra and Srivasatava 1971, 1972, 1973, Kumar and
Saxena 1970 and Ram Reddy et al 1978a, 1978b, 1979). The mechanism
formulated in these studies involved the formation of SO4-..
S2O82 2 SO4
. (1.13)
The rate determining step is catalyzed by metal ions. The above
mechanism explained the zero order dependence of rate with respect to amino
acid concentration. The following rate equation is deduced for Cu(II)
catalyzed oxidation of glycine and alanine (Ram Reddy et al 1978b).
22 3 0.52 8
2 8-d[S O ] = k [S O ] [amino acid] [Cu(II)]
dt (1.14)
The kinetics of oxidation of dl-alanine by hydrogen peroxide in the
presence of Fe(II) ions was studied by Ashraf et al (1979). The reaction was
first order with respect to [alanine], [Fe(II)] and zero order in [H2O2]. A free
radical mechanism as shown below was postulated for this reaction in which
carbon dioxide, acetaldehyde and ammonia were the reaction products.
Fe2+ + H2O2 Fe (OH)2 + HO. (1.15)
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R CH COOH+ OH. H2O
H2O2
OHO
+ NH4+
RCHO + CO2
R C COOH +.
R C COOH
NH3+
NH3+NH3+
NH3+NH3+
R C COOH
Scheme 1.1 Plausible mechanism for the oxidation of dl-alanine by H2O2
in the presence of Fe(II) ion
The kinetics of oxidation of glycine, alanine, phenyl alanine, serine,
threonine and aspartic acid by permanganate in aqueous medium were
investigated by Surender Rao et al (1979). The rate was found to be
-d[Mn(VII)] = k [amino acid][Mn(VII)]dt
(1.16)
The reactions were found to be acid catalyzed and the kinetic data indicated
participation of water molecules in the rate determining step as proton
abstracting agent from the substrate as per Bunnett’s hypothesis. As Ag(I)
was found to catalyze these reactions, the rate law was found to be
K k'' [amino acid] [Ag(I)]-d[Mn(VII)] c=dt 1+ K [Ag(I)]+ K [amino acid]
(1.17)
Kinetics of oxidation of amino acids viz., glycine, alanine, valine,
leucine, phenyl glycine and phenyl alanine by diperiodatoargentate (III) was
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investigated in alkaline medium by Jayaprakash Rao et al (1985). The
oxidation products were found to be ammonia and the corresponding
keto acids. The presence of electron-withdrawing groups at the -carbon
increased the rate of oxidation while the rate decreased with increasing alkyl
chain length. A two electron transfer mechanism was proposed to explain the
results.
The rate of Mn(II) catalyzed O2 production in the presence of
alanine or leucine was about 4-fold the rate observed in the absence of amino
acids and accounts for about half of the H2O2 consumed. The other half of the
H2O2 was consumed in the oxidation of amino acids. In contrast, O2
production was increased nearly 18-fold in the presence of -methylalanine
and accounts for about 90% of the H2O2 consumed. Oxidation of amino acid
in this complex most likely proceeded by a free radical mechanism involving
hydrogen abstraction from the -carbon. Mn(II) could able to facilitate
Fenton-type reactions (Berlett et al 1990).
2 H2O2 + RCHNH3+COO RCOO + CO2 + NH4
+ + 2H2O (1.18)
Metal catalyzed oxidation (MCO) or non-enzymatic oxidation gave
rise to highly reactive intermediates such as hydroxyl radicals. The hydroxyl
radicals cause simple amino acids to lose a hydrogen atom at the -position
and form a carbon-centered radical, which in the presence of O2 generates
hydroperoxides (Fu et al 1995). These hydroperoxides are unstable in the
presence of redox active transition metal ions like copper and decompose
rapidly to -keto acids, aldehydes and/or carboxylic acids corresponding to
the oxidized amino acid after losing the amine group as NH3 and/or carboxyl
group as CO2 (Stadtman and Berlett 1991).
The oxidation of amino acids by Fentons reagent (H2O2 + Fe(II))
led to the formation of NH4+, -keto acids, CO2, oximes, aldehydes or
21
carboxylic acids containing one carbon atom less. Oxidation is almost
completely dependent on the presence of bicarbonate ion and is stimulated by
iron chelators at levels which are sub-stoichiometric with respect to [Fe(II)]
but is inhibited at higher concentrations. The stimulatory effect of chelators is
not merely due to solubilization of catalytically inactive polymeric forms of
Fe(OH)3 but also due to the conversion of Fe(II) to complexes. The results
suggested that an iron chelate and another form of iron are required for
maximum rate of amino acid oxidation. The metal ion-catalyzed oxidation of
amino acids is likely to be a caged process since the oxidation is not inhibited
by hydroxyl radical scavengers. The relative rates of oxidation of various
amino acids by Fenton system as well as the distribution of products formed
especially products of aromatic amino acids are significantly different from
those reported for amino acid oxidation by ionizing radiation. Several iron-
binding proteins, peptides and hemoglobin degradation products can replace
Fe(II) or Fe(III) in the bicarbonate dependent oxidation of amino acids
(Stadtman and Berlett 1991).
Several metals or metalloid ions exist in multiple oxidation states
and can undergo electron transfer reactions that are important in biological
and environmental systems. There are endogenous metal ions such as Fe(II),
Cu(II) and Co(II) that participate in the oxidation-reduction reactions with
species like molecular dioxygen, superoxide and hydrogen peroxide. These
reactions may be modulated by endogenous reducing agents such as
glutathione, ascorbate and tocopherol. The reactions can be described in terms
of thermodynamics using the standard electrode potentials. The favorable
reaction will depend on the concentration of the reactants and may depend on
the pH and/or the presence of organic ligands that form complexes with metal
or metalloid. As(V) can react with glutathione in buffered aqueous solution to
produce As(III) and oxidized glutathione. This reaction may be important in
the methylation reactions of arsenic. Arsenic species can decrease the red
22
blood cell levels of reduced glutathione, but the products of oxidation and the
mechanism of oxidation are more complex than those found in water alone.
Cr(VI) was found to interact with DNA after reacting first with a reducing
agent such as glutathione to form lower oxidation states of chromium. These
examples illustrate the importance of oxidation-reduction reactions for toxic
metals and metalloids (Carter 1995).
Tungsten catalyzed oxidative decarboxylation of N-alkyl amino
acids with hydrogen peroxide under phase transfer condition gave the
corresponding nitrones with good yield (Murahashi et al 1994). The oxidation
of organic compounds including amino acids by tungsten catalyzed H2O2 was
studied by Miochowski and Said (1977). The kinetic and mechanism of
permanganic oxidation of L-glutamine in sulfuric acid was carried out both in
the presence and absence of Ag(I) (Iloukhani and Bahrami 1999). The overall
rate expression for the oxidation may be written as
- d[Mn(VII)] = k [L- glutamine] [Mn(VII)]°dt (1.19)
The rate law is expressed as in equation (1.20) in the presence of Ag(I),
K k'' [L- glutamine] [Ag(I)]- d[Mn(VII)] c °= [Mn(VII)]dt 1+ K [Ag(I)] + K [L-glutamine]
(1.20)
The kinetics of Os(VIII) catalyzed oxidation of DL-methionine by
hexacyanoferrate (III) in aqueous alkaline medium was studied spectrophoto
metrically (Jose et al 2006). The analysis revealed that there is a decrease in
the dielectric constant of the medium with increase in the rate of the reaction.
The addition of products has no effect on the rate of reaction. Os(VIII) binds
to OH species in the step prior to equilibrium step to form hydroxide species
which reacts with [Fe(CN)6]3 in a slow step to form an intermediate species.
This reacts with a molecule of DL-methionine in a fast step to give sulfur
23
radical cation of methionine and yields sulfoxide product by reacting with
another molecule of [Fe(CN)6]3 .
The kinetics of Cu(II) autocatalyzed oxidation of threonine by well
recognized analytical reagent diperiodatocuprate (III) (DPC) in aqueous
alkaline medium was studied spectrophotometrically (Jose and Tuwar 2007).
The reaction between DPC and threonine in alkaline medium exhibited 2:1
stoichiometry (DPC: threonine). The reaction was first order each in [DPC]
and [threonine] and less than unity in [alkali]. Periodate has retarding effect
on the rate of reaction. Ionic strength has negligible effect on the reaction.
Increase in the dielectric constant of the medium with a decrease in the rate of
the reaction was observed. The main products were identified by spot test and
IR spectra.
Kinetics and mechanism of oxidation of leucine and alanine by
Ag(III) complex were studied spectrophotometrically in alkaline medium at a
constant ionic strength (Song et al 2008). The reaction was first order with
respect to Ag(III) complex and amino acids (leucine and alanine). The
second-order rate constant (k) decreased with increase in [OH ] and [IO4].
The oxidation of amino acid, L-tryptophan (L-trp) by diperiodatoargentate (III)
(DPA) in alkaline medium was studied spectrophotometrically (Tatagar et al
2009). The reaction between L-trp and DPA in alkaline medium exhibited 1:2
stoichiometry. The involvement of free radicals was observed in the reaction
based on the observed order and experimental evidences. The products were
identified by spot test and characterized by spectral studies. The reaction
constants revealed different steps of mechanism.
Kinetics of oxidation of L-cystine by hexacyanoferrate (III) was
studied in alkaline medium at 300 ºC (Nowduri et al 2009). The reaction was
followed spectrophotometrically at max = 420 nm. The reaction was found to
be first order dependence each on [hexacyanoferrate (III)] and [cystine]. It
24
was found that the rate of the reaction increased with increase in [OH-]. The
oxidation product of the reaction was found to be cysteic acid. The oxidation
of glycine and alanine by bis(dihydrogen-tellurto)argentite (III) ion was
studied by stopped-flow spectrophotometer (Huo et al 2009). The reaction
was first order in Ag(III) complex and less than unit order in glycine and
alanine. A plausible mechanism was proposed from the observed kinetic
results.
1.5 PEROXOCOMPOUNDS
The peroxo oxidants are the derivatives of hydrogen peroxide
(H-O-O-H) formed by the replacement of hydrogen atom by the groups of
sulphate, phosphate and carbonate. The weak peroxide bond (–O-O-) makes
the peroxides highly reactive with easily oxdisable molecules. They
spontaneously decompose in solution leading to more stable products. The
–O-O- linkage undergoes cleavage during the reaction and makes sensitive
towards trace amount of catalysts and promoters, which can accelerate the
decomposition. Hence the kinetics and mechanism of oxidation reactions of
amino acids with peroxo compounds are reviewed in the following section.
Peroxo oxidants such as peroxomonosulphate (PMS), peroxodisulphate
(PDS), peroxomonophosphate (PMP) and peroxomonocarbonate (PMC) have
gained paramount importance due to their utilisation as auxiliary reagents in
organic synthesis (Swern 1971, Sosnovsky and Rawlison 1971 and Adam
et al 1992).
1.5.1 Peroxomonosulphate (PMS) (Caro’s acid)
Among the oxidant, PMS (HSO5 ) is a dibasic acid having two
ionisable protons, one of them resembling sulphuric acid proton and the other
peroxide proton (Ball and Edwards 1956). The first pKa value of the peroxide
proton was found to be 9.4±0.2 which is lower than that of hydrogen peroxide
25
(pKa = 11.65). It is also significantly higher than that of peroxoacetic acid. IR
studies revealed that the -O-O- stretching frequency is higher than that of
H2O2 and the two -OH groups are structurally different (Arnau and Giguere
1970).
HSO4 + H2O HSO5 +2H+ + 2e (1.22)
The standard electrode potential (E°) for the couple HSO4 /HSO5
was estimated to be -1.75 V (Spiro 1979). Later Stele and Appleman (1982)
estimated the value as -1.82 V. This high potential value suggested that many
room temperature oxidation can be carried out with PMS. The potentiality of
PMS as a powerful oxidant is brought by its ability to oxidize many organic
and inorganic compounds (Kennedy and Stock 1960). The high oxidation
potential of PMS and the propensity of HSO5 to react via oxygen transfer
(Trost and Curran 1981, Johnson and Balahura 1987 and Edwards and Marsh
1989) make this molecule as a favourable one for the oxidation of various
organic compounds in aqueous solution.
PMS is a two equivalent oxidant and in a number of systems,
oxygen atom transfer occurs from the terminal peroxide position following
nucleophilic attack (at the peroxide moiety) by the substrate (Fortnum et al
1960, Edwards and Muller 1962, Johnson and Edwards 1966, Secco and
Venturini 1976 and Thomson et al 1979). Moreover the superior ability of
HSO4 to act as a leaving group permits HSO5 oxidation proceeded at a
faster rate than other peroxides. PMS could be a viable alternative to H2O2
and S2O82 for the control of sulphide induced corrosion in concrete sanitary
sewers and wastewater treatment facilities. The effectiveness of HSO5 for
H2S oxidation in wastewater from wastewater treatment plant was studied and
compared with H2O2 and S2O82 oxidation. In this report, PMS was found to
26
be more effective based on its molar stochiometry of oxidant used in the total
sulphide oxidation.
1.5.1.1 Decomposition of peroxomonosulphate
PMS in aqueous acetonitrile readily converted aryl and thio
benzoates to carboxylic acid and sulphonic acid respectively (Bunton et al
1995). The rate of reaction increased with increase in the water content in
acetonitrile. Studies on the oxidative hydrolysis of phosphorous(V) esters of
thiols (Blasko et al 1997) by PMS revealed that HSO5 converted
phosphorous(V) esters of thiols into phosphorous(V) and sulphuric acid.
Kinetic and mechanism of decomposition of Caro’s acid was
carried out in alkaline medium (Ball and Edwards 1956). The rate of
decomposition of Caro’s acid is given as in equation (1.22).
-5- d[HSO ]
dt = k [HSO5
-] [SO52-] (1.22)
The rate limiting step is the nucleophilic attack of SO52 on the
peroxide oxygen of HSO5 . Studies on the decomposition of PMS were also
reported by Goodman and Robson (1963). The results revealed that the
kinetics of decomposition was identical with peroxyaromatic acids. The
activation energy was calculated to be 11.9 Kcal/mole. The rate limiting step
proposed is the nucleophilic attack of SO52 on the sulphur atom of HSO5 .
The intermediate then decomposed to give S2O82 and HO2 .
The decomposition of PMS over a wide range of pH was
investigated by Kyrki (1963). The results confirmed the direct interaction
between SO52 and OH in highly alkaline medium. From the results
obtained, the rate equation (1.23) can be derived.
27
5-d[HSO ]dt
= k [SO52 ] [OH ] (1.23)
The mechanism proposed in acid medium involved the formation of
H2O2 by hydrolysis of HSO5 as given in equation (1.24).
HSO5 + H3O+ HSO4 + H2O2 + H+ (1.24)
The rate equation can be given as
5-d[HSO ]dt
= k [H+] [HSO5 ] (1.25)
The rate of decomposition of PMS in 0.5 M perchloric acid is very
slow at room temperature. But the decomposition rate of PMS is remarkably
enhanced by the presence of moderate concentration of Ag(I) (Thompson
1981). Plots of [PMS] vs. time were strictly linear for at least 90% of the total
reaction. The rate expression is
5-d[HSO ]dt
= k (1.26)
Detailed investigation was made on the formation of singlet
molecular oxygen for the ketone- catalyzed decomposition of
peroxomonosulfuric acid by infrared phosphorescence measurements (Lange
and Brauer 1996). Janakiram et al (1998) investigated the mechanism of
oxidation of aliphatic acetals by PMS. The reaction was found to be first order
in [PMS] and [acetal]. The oxidation reaction was independent of dielectric
constant of the medium. A hydride ion shift was proposed in the mechanism.
Kinetics and mechanism of thermal decomposition of PMS by phase transfer
catalyst (PTC) viz., tetrabutylammonium chloride (TBAC) and
tetrabutylphosphonium chloride (TBPC) was reported by Balakrishnan and
28
Kumar (2000). The effects of [PMS], [PTC], ionic strength of the medium (µ)
and temperature on the rate of decomposition of PMS was first order in
[PMS] for TBAC and half order for TBPC.
The kinetics and mechanism of decomposition of Caro’s acid in
aqueous sodium hydroxide in the presence of -cyclodextrin have been
reported. The rate of decomposition of PMS was enhanced by the formation
of -cyclodextrin peroxyanion by the interaction between SO52 and
-cyclodextrin anion. The self decomposition of PMS in alkaline medium
catalyzed by -cyclodextrin involved molecular intermediate rather than free
radical intermediate (Ramachandran and Lathakannan 2002). The kinetics of
ketone-catalyzed decomposition of Caro’s acid was studied in aqueous
alkaline medium at 25 ºC. The rate followed simple second-order kinetics,
first order each in [ketone] and [PMS]. The rate constant values were
independent of hydroxide ion concentration over the range from 0.05 to
0.15 M. The experimental results suggested that the nucleophilic addition of
SO52 ion at the carbonyl carbon led to the formation of oxirane intermediate
in the rate-determining step. Oxirane reacted rapidly with another SO52 to
give the parent ketone, oxygen and SO52 (Selvarani et al 2005). PMS is an
inexpensive, safe and environmentally benign oxidant for C-H bond
oxygenation. It could catalyze acetoxylation, etherification of arene and
alkane C-H bonds. Hence PMS is proved to be effective in acetic acid and/or
methanol. So these transformations are applied to a wide variety of substrates
(Desai et al 2006). The kinetics and mechanism of Mn(II) catalyzed
decomposition of PMS in highly alkaline medium was attempted by Sundar
et al (2008). The value of pseudo first order rate constant (kobs) decreased with
increase in [PMS]0 and increased with increase in [OH-] and [Mn(II)].
Analysis of the results revealed that the formation of manganese peroxide
(MnO2) as molecular intermediate which decomposed rapidly. The rate law is
29
-2
252 5
-- d[SO ] -= k [complex] + k [SO ]1dt (1.27)
1.5.1.2 Oxidation of organic substrates by PMS
The oxidation of halide (Fortnum et al 1960), nitrite (Edwards and
Muller 1962) and chlorite ion (Johnson and Edwards 1966) by PMS was
reported in the literature. The rate limiting step in these oxidation reactions
was found to be nucleophilic attack on the peroxide oxygen. The oxidation of
nitrite by PMS was also studied by Sharma et al (1992a). The product of
oxidation was identified as nitrate.
HSO5 + NO2 SO42 + NO3 + H+ (1.28)
Oxidation of aldehyde by PMS in the presence of alcohols
produced high yield of esters of corresponding acids (Nishihara and Kubota
1968). Further studies confirmed that the rate equation (1.27) is invalid in the
presence of H2O2, S2O82 and metal ions. Isotopic labeling studies revealed
that the peroxo bond in HSO5 was broken in alkaline medium (Koubek et al
1964). Review on the self decomposition of PMS and peroxides were reported
in the literature (Curci and Edwards 1970).
The oxidation of azide by peroxomonosulphate was studied in acid,
neutral and basic solutions, and the oxidation of [Cr(NH3)5N3]2+ by PMS was
studied in acidic medium. The reaction with azide obeyed the rate law:
k1 [N3 ] [HSO5 ] in acid and neutral solution and the rate law: k2 [N3 ]
[SO52 ] in base. The reaction with [Cr(NH3)5N3]2+ obeyed the rate law:
k3 [Cr(NH3)5N3]2+ [HSO5 ] (Thompson et al 1979). This study revealed that
the reaction of azide complex with PMS is the convenient route for the
synthesis of nitrosyl complexes. The reaction proceeded through transfer of
30
terminal peroxide oxygen of PMS to reductant, resulting in the formation of
N2, N2O in azide and [Cr(NH3)5NO]2+in [Cr(NH3)5N3]2+ respectively.
The kinetics of Baeyer-Villiger oxidation of biacetyl and benzyl by
PMS and PMPA was studied in different pH range at 308 K. The rate of
oxidation was strongly pH dependent and the rate increased with increase in
pH. From the pH data the reactivity of different peroxo species in the
oxidation was determined by Panda et al (1988). The detailed kinetic study of
PMS with dimethyl sulphoxide was studied by Pandurangan and
Maruthamuthu (1981). The following products viz., aldehydes (Renganathan
and Maruthamuthu 1986a, Ramachandran et al 1986 and Naseruddin et al
1987), ketones (Manivannan and Maruthamuthu 1986 and Panda et al 1988),
sulphides (Betterton 1992 and Bacaloglu et al 1992), dimethyl aniline (Ogata
and Tabushi 1958), cyclohexene and cyclooctene (Taquikhan et al 1990),
indoles (Balon et al 1993 and Carmono et al 1995) and aniline (Jameel and
Maruthamuthu 1998) were reported. Oxidation of aldehydes to the
corresponding acids by oxone (Caro’s acid) in aqueous acetone was also
reported by Webb and Ruszkay (1998).
The first order (k0) and heterogenetic (k1) rate constants showed
first order dependence on [PMS] and 1/[H+]. The reaction was catalyzed by
the addition of chelating ligand glycine and k1 showed first order dependence
on [glycine] at a fixed pH. This catalysis was ascribed to complexation
whereby the redox potential for Mn (gly)n(2-n)+ was lower than for Mnaq
2+ ,
facilitating the oxidation (Lawrence and Ward 1985). Oxidation of Mnaq2+ to
Mn(II) by PMS in acetate buffer was autocatalytic and obeyed the rate
expression of the general form
0 1 x- d[Mn(II)] = k [Mn(II)] + k [Mn(II)] [MnO ]
dt (1.29)
31
The kinetics of oxidation of aliphatic aldehydes such as
formaldehyde, acetaldehyde, propionaldehyde, n-butyraldehyde and
trichloroacetaldehyde by peroxomonosulphate was carried out in aqueous
perchloric acid medium at a constant ionic strength of 1.2M in the
temperature range 283-333 K (Renganathan and Maruthamuthu 1986a). The
reaction of all aldehydes was found to be a total second order kinetics, first
order each with respect to [PMS] and [aldehyde]. The rate law is expressed as
+5a b
--d[HSO ] = k [PMS] [aldehyde] [H ] + k [PMS] [aldehyde]dt
(1.30)
The kinetics of oxidation of hydroxylamine by PMS in acetate
buffered solution was investigated by Sharma et al (1992b). The reaction
proceeded through a free radical mechanism. Under the experimental
conditions, PMS produced two pairs of possible species viz., OH + SO4.
and OH + SO42 . The reactivity order was S2O8
2 < HSO5 > H2O2. The
kinetic reaction and rate law are as given in equation 1.31 and 1.32.
2 HSO5 +NH3OH+ 2 SO42 + N2 + 3 H2O + 3H+ (1.31)
5-- d[HSO ]
dt= k K [PMS] [NH3OH+] (K + [H+])-1 (1.32)
The oxidation of bisulphate ion by PMS to form sulphate ion was
studied in the pH range 3.8-7.9 (Connick et al 1993). The proposed
mechanism involved the formation of pyrosulphate ion (S2O72-) as an
intermediate, which undergoes hydrolysis to form sulphate and hydrogen
ions. The overall reaction is given as
HSO5 + HSO3 2SO42 + 2H+ (1.33)
32
The rate law derived is given in equation (1.34)
5-- d[HSO ]
dt = (kaaH
+ + kbaH-1 +kc) + [HSO5 ] [HSO3 ] (1.34)
where aH+ is the activity of hydrogen ion and aH
-1 is activity of hydrogen
anion.
Oxidation of hypophosphorous acid by PMS in acid medium was
investigated by Dubey et al (2002). The reaction was first order with respect
to [hypophosphorous] and [PMS]. The oxidation of indole-3-acetic acid by
PMS in acetonitrile medium was also reported by Chandramohan et al (2002).
The reaction followed total second order, first order each with respect to
[indole-3-acetic acid] and [PMS]. The rate of the reaction was not affected by
[H+]. Variation of ionic strength (µ) did not show any significant effect on the
reaction rate. Increase in the percentage of acetonitrile decreased the rate. The
rate of the reaction proceeded through a non-radical pathway. Kinetics of
oxidation of ascorbic acid (AH2) by PMS was determined in acidic, neutral
and alkaline conditions over the temperature range 286-301 K (Raja et al
2003). The reactions were found to obey total second order kinetics, first
order each with respect to [PMS] and [ascorbic acid].
- 2- d[AH ]dt
= k2 [PMS] [AH2] (1.35)
The stoichiometry of the reaction revealed the absence of self
decomposition of PMS. The addition of neutral salt (NaClO4) found to
increase the reaction rate. A suitable mechanism was proposed based on the
formation of free radical intermediates such as hydroxyl, sulphate and
ascorbate. Accelerated bleaching, photobleaching and mineralization of non-
biodegradable azo-dye, Orange II was observed with PMS in the solution of
Co2+ ions. The results revealed that the bleaching rate of Orange II in the dark
33
was found to follow first order kinetics with respect to [Co2+] and reaction
was found to proceed by a chain radical branched mechanism (Fernandez et al
2004). PMS, an efficient oxidant for the photo catalyzed degradation of a
textile dye, acid red 88 was studied by Madhavan et al (2009). The rate of
photodegradation of dye decreased with increase in dye concentration. Total
organic carbon (TOC) analysis revealed rapid mineralization of acid red 88 in
the presence of PMS.
Oxidation of indole by PMS in aqueous acetonitrile was studied by
Meenakshisundaram and Sarathi (2007). Analysis of the results showed that
HSO5 and SO52 are the respective electrophile in acidic and basic media.
Nucleophilic attack of the ethylene bond on the persulphate oxygen was
envisaged to explain the reactivity. The reaction failed to initiate
polymerization and hence radical mechanism was ruled out. The values of
thermodynamic parameters suggested a bimolecular process. Catalytic
activity was not significantly observed for the reaction system in the presence
of Ag(I), Cu(II) and heteroaromatic N-bases. Oxidation of tris (1, 10-
phenanthroline) Fe(II) with PMS ion was first order with respect to both
substrate and oxidant. The rate is accelerated by alkali metal ion due to the
formation of an ion pair between M+ and SO52- ions and not between M+ and
HSO5- ions (Mehrotra and Mehrotra 2008).
1.5.1.3 Oxidation of organic substrates by PMS in the presence of
metal ions
Evidences have been reported for the catalytic decomposition of
Caro’s acid by specific substances in aqueous phosphate buffer. Co(II) and
Mo(VI) are capable and effective catalyst although other metal ions can also
act as catalyst for the decomposition. Cobalt catalyzed decomposition of
Caro’s acid was found to be second order kinetics (Ball and Edwards 1958).
34
Homogeneous catalyst systems are extensively used in the manufacture of
important chemicals. Compared to heterogeneous catalysts, the number of
chemicals manufactured using homogeneous catalyst are quite less. However,
there are reports for the catalytic oxidation of alkenes to aldehydes or ketones
in homogeneous reaction using Pd(II) and Cu(II) catalyst (Smidt et al 1962).
Studies on the uncatalyzed decomposition of PMS in the pH range
6-12 revealed that oxygen is evolved and terminal peroxide oxygen is
incorporated into this oxygen product (Koubek et al 1964). In strongly acidic
medium the product formed is hydrogen peroxide with both oxygen atoms
originating from the peroxide moiety in PMS. However, in the metal ion
catalyzed decomposition neither of the above modes are competitive and
instead a redox process is operative. Metal ion catalyzed decomposition of
PMS in acidic and weakly alkaline medium are widely reported in the
literature (Ball and Edwards 1958, Mariano 1968, Billing et al 1970,
Thomson 1981, Gilbert and Stell 1990a, b and Bennet et al 1991). Metal ions
are observed to influence the rate of decomposition of PMS to a greater
extent. The catalytic effect of different metal ions on the rate of
HSO5 decomposition was studied in acid medium. The results revealed that
the catalytic effect of Ag(I) is profound compared to other metal ions such as
Cu(II), Ti(III) and Co(II) (Rettmer et al 1970).
Detailed studies on Ce(IV) induced decomposition of PMS in
acidic medium (Billing et al 1970) revealed that the peroxy radical generated
from HSO5 is SO5. rather than HSO5
. as shown below.
Ce (IV) + HSO5. Ce (III) + H+ + SO5
. (1.36)
SO5. is proposed as an intermediate in the radiolysis of S2O8
2
(Atkins et al 1963) and in other radiolysis reactions as well (Maruthamuthu
35
and Neta 1977, Huie and Neta 1984 and Deister and Warneck 1990).
However it is reported that SO4. is formed during the decomposition of PMS
by low valency metal ions such as Fe(II) and Ti(III) (Gilbert and Stell 1990a).
Mn+ + HSO5 Mn+1 + OH + SO4. (1.37)
OH. is the proposed intermediate during the decomposition of PMS
at pH ~ 2.0 by Cu(I) which is produced in situ by the reaction between Ti(III)
and Cu(II) as shown in equation (1.38).
Ti (III) + Cu (II) Ti (IV) + Cu (I) (1.38)
However, Fe(II) and Ti(III) react with HSO5 also produce SO4. (Gilbert and
Stell 1990a).
Kinetics of oxidation of some sulphoxides with oxone in the
presence of Ru(III) have been reported (Meenakshisundaram and
Sathiyendiran 2000). The catalytic activity of Ru(III) was demonstrated with
several diaryl, dialkyl and alkyl aryl sulphoxides, all of which were found to
undergo oxidation under homogenous condition. The rate increased
substantially with increasing water content in aqueous acetic acid. The rate
data were consistent with the mechanism involving electron transfer from
electrophilic perhydroxyl oxygen of oxone to sulphoxide.
The kinetics of oxidation of glycolic acid and -hydroxy acid by
PMS were studied in the presence of Ni(II) and Cu(II) ions in the acidic pH
range of 4.05-5.89. The metal glycolate and not glycolic acid was oxidized by
PMS. The rate was found to be first order in [PMS] and metal ion
concentrations. The oxidation of nickel glycolate is zero order in glycolic acid
and inverse first order in [H+]. The increase of glycolic acid decreased the rate
in copper glycolate, and the rate constants initially increased and then
36
remained constant with pH. The results suggested that metal glycolate reacts
with PMS through a metal-peroxide intermediate, which transformed slowly
into a hydroperoxide intermediate by the oxygen atom transfer to hydroxyl
group of the chelated glycolic acid (Shailaja and Ramachandran 2009).
The kinetics of redox reactions of PMS with Fe(IV), Ce(III),
chloride, bromide and iodide ions were reported (Gabor Lente et al 2009).
Ce(III) is only oxidized upon illumination by UV light and Ce(IV) is
produced in the photoreaction. Fe(II) and V(IV) are most probably oxidized
through one electron transfer producing sulphate ion radicals as intermediate.
The halide ions are oxidized by two electron process, which most likely
include oxygen-atom transfer. Comparison with literature data suggested that
activation entropies may be used as indicators distinguishing between
heterolytic and homolytic cleavage of the peroxo bond in the redox reaction
of PMS. PMS ion in the presence of Ni(II) lactate and formaldehyde in the pH
between 4.0 and 5.9 undergo self-decomposition and evolve oxygen. The
reaction is second order in [PMS]. Experimental results revealed that
hemiacetal of Ni(II) lactate catalyzed the self-decomposition (Murugavelu
et al 2009).
1.5.2 Literature Review on the Oxidation of Amino acids by PMS
and Metal Ions
The kinetics of oxidation of amino acids by PMS in the presence
and absence of formaldehyde were studied. Analysis of the results revealed
that the rate can be represented at constant [H+] and in the absence of
formaldehyde as
5-d[HSO ]dt
= ka [amino acid] [PMS] (1.39)
37
in the presence of formaldehyde and at constant [H+], the rate law can be
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