A kinetic Study of the Oxidation of two
Simple Sugars by Peroxydisulphate
A thesis submitted
For
The degree of M.Sc. in chemistry
By
Nasma Dafalla Eljack
(B.Sc. U of K 1998)
Department of Chemistry
Faculty of Science
University of Khartoum
Khartoum, Sudan
August 2004
To my parents,
Husband,
Daughter
And Son.
Nasma
Many thanks and sincere gratitude to my
supervisor Professor Ali-M-Kheir for helpful discussions,
valuable
advice
and
constant
encouragement
throughout this work.
Thanks
to
my
family,
colleagues
and
friends for their assistance, encouragement
and moral support during the course of my
study.
I
am
grateful
to
the
University
of
Khartoum for providing me the opportunity to
conduct this work.
Thanks are also extended to Salwa Hamadto for
typing this work.
Last, but not least special thanks are
due
to
my
encouragement
study.
husband
and
Ali
Taha
support
Mahadi
for
throughout
his
this
Abstract
The present work deals with a kinetic study of redox reactions of two
simple sugars (Glucose and Maltose) with peroxydisulphate ion under
uncatalysed conditions over the temperature range (60ْ – 80ْ C).
It was found that these reactions follow consecutive two main paths:
(i)
Thermal decomposition of peroxydisulphate.
(ii)
Bimolecular reactions of these sugars with peroxydisulphate.
Two aspects of the study were involved:
(a) Kinetic study.
(b) Products analysis.
The progress of these reactions was followed by examining the
concentration of peroxydisulphate iodometrically at different times.
The study showed that the reaction rate follows a first order in
peroxydisulphate and a fractional order of (0.25) in the simple sugars
concentrations.
Thermodynamic parameters, activation energy, frequency factor,
entropy change and free energy change of the redox reactions were
calculated from rate constants obtained at different temperatures.
The rate laws of the reactions were determined and free radical
mechanisms were proposed.
The analysis of the reaction products revealed the presence of
formaldhyde and formic acid in the volatile fraction as main products and
these were confirmed by spot tests.
The
amounts
of
formaldhyde
were
determined
spectro-
photometrically and those of formic acid were determined titrimetrically.
‫ﻣﻠﺨﺺ اﻷﻃﺮوﺣﺔ‬
‫ﻫﺫﺍ ﺍﻟﺒﺤﺙ ﻴﺘﻨﺎﻭل ﺩﺭﺍﺴﺔ ﺤﺭﻜﻴﺔ ﻭﺩﻴﻨﺎﻤﻴﻜﻴﺔ ﺍﻟﺘﻔﻜﻙ ﺍﻟﺤﺭﺍﺭﻱ ﻟﺒﻴﺭﻭﻜﺴﻲ ﺜﻨﺎﺌﻲ‬
‫ﻜﺒﺭﻴﺘﺎﺕ ﺍﻟﺒﻭﺘﺎﺴﻴﻭﻡ ﻭﺃﻜﺴﺩﺓ ﺃﺜﻨﻴﻥ ﻤﻥ ﺍﻟﺴﻜﺎﻜﺭ ﺍﻟﺒﺴﻴﻁﺔ )ﺠﻠﻜﻭﺯ ﻭﻤﺎﻟﺘﻭﺯ( ﺒﻔﻌل ﺃﻴﻭﻥ‬
‫ﺒﻴﺭﻭﻜﺴﻲ ﻜﺒﺭﻴﺘﺎﺕ ﺍﻟﺒﻭﺘﺎﺴﻴﻭﻡ ﺩﻭﻥ ﺍﺴﺘﺨﺩﺍﻡ ﻋﺎﻤل ﻤﺴﺎﻋﺩ‪ ،‬ﻓﻰ ﺩﺭﺠﺔ ﺤﺭﺍﺭﺓ ﻤﺎ ﺒﻴﻥ‬
‫)‪(80ْ-60‬ﻡ‪ ،‬ﺇﺫ ﺃﻥ ﺍﻟﺘﻔﺎﻋل ﻻ ﻴﺘﻡ ﻓﻰ ﺩﺭﺠﺔ ﺤﺭﺍﺭﺓ ﺍﻟﻐﺭﻓﺔ ﺍﻟﻌﺎﺩﻴﺔ ﻭﺘﻡ ﺘﺘﺒﻊ ﻫﺫﻩ ﺍﻟﺘﻔﺎﻋﻼﺕ‬
‫ﺒﻘﻴﺎﺱ ﺘﺭﻜﻴﺯ ﺒﻴﺭﻭﻜﺴﻲ ﺜﻨﺎﺌﻲ ﺍﻟﻜﺒﺭﻴﺘﺎﺕ ﺒﺎﻟﻤﻌﺎﻴﺭﺍﺕ ﺍﻷﻴﻭﺩﻭﻤﻴﺘﺭﻴﺔ ﻓﻰ ﻓﺘﺭﺍﺕ ﺯﻤﻨﻴﺔ ﻤﺨﺘﻠﻔﺔ‪.‬‬
‫ﻭﻗﺩ ﺃﺜﺒﺘﺕ ﺍﻟﺩﺭﺍﺴﺔ ﺍﻟﺤﺭﻜﻴﺔ ﻟﻬﺫﻩ ﺍﻟﺘﻔﺎﻋﻼﺕ ﺃﻨﻬﺎ ﻤﻥ ﺍﻟﺭﺘﺒﺔ ﺍﻷﻭﻟﻲ ﺒﺎﻟﻨـﺴﺒﺔ ﻟﺘﺭﻜﻴـﺯ‬
‫ﺒﻴﺭﻭﻜﺴﻲ ﺜﻨﺎﺌﻲ ﺍﻟﻜﺒﺭﻴﺘﺎﺕ ﻭﻤﻥ ﺍﻟﺭﺘﺒﺔ ﺍﻟﻜﺴﺭﻴﺔ )‪ (0.25‬ﺒﺎﻟﻨﺴﺒﺔ ﻟﺘﺭﻜﻴﺯ ﻜل ﻤـﻥ ﺍﻟﺠﻠﻜـﻭﺯ‬
‫ﻭﺍﻟﻤﺎﻟﺘﻭﺯ‪.‬‬
‫ﻜﻤﺎ ﺘﻡ ﺤﺴﺎﺏ ﺍﻟﺜﻭﺍﺒﺕ ﺍﻟﺜﻴﺭﻤﻭﺩﻴﻨﺎﻤﻴﻜﻴﺔ ﻭﻫﻲ ﻁﺎﻗﺔ ﺍﻟﺘﻨﺸﻴﻁ‪ ،‬ﻋﺎﻤل ﺍﻟﺘـﺭﺩﺩ‪ ،‬ﺘﻐﻴـﺭ‬
‫ﺍﻻﻨﺘﺭﻭﺒﻴﺎ )ﺍﻟﻔﻭﻀﻰ ﺍﻟﺠﺯﺌﻴﺔ( ﻭﺍﻟﻁﺎﻗﺔ ﺍﻟﺤﺭﺓ ﻭﺫﻟﻙ ﻤﻥ ﺜﻭﺍﺒﺕ ﺴﺭﻉ ﺍﻟﺘﻔﺎﻋﻼﺕ ﻋﻨﺩ ﺩﺭﺠـﺎﺕ‬
‫ﺤﺭﺍﺭﺓ ﻤﺨﺘﻠﻔﺔ‪.‬‬
‫ﻭﻗﺩ ﺘﻡ ﻭﻀﻊ ﻗﺎﻨﻭﻥ ﻟﻠﻤﻌﺩل ﻟﻜل ﻫﺫﻩ ﺍﻟﺘﻔﺎﻋﻼﺕ ﻜﻤﺎ ﺘﻡ ﺍﻗﺘﺭﺍﺡ ﻤﻴﻜﺎﻨﻴﻜﻴﺔ )ﺁﻟﻴﺔ( ﻫـﺫﻩ‬
‫ﺍﻟﺘﻔﺎﻋﻼﺕ‪.‬‬
‫ﻭﺒﺈﺠﺭﺍﺀ ﺘﺤﻠﻴل ﻟﻨﻭﺍﺘﺞ ﻫﺫﻩ ﺍﻟﺘﻔﺎﻋﻼﺕ ﻭﺠﺩ ﺃﻨﻬﺎ ﺘﺘﻜﻭﻥ ﻤﻥ ﻨﻭﺍﺘﺞ ﻁﻴﺎﺭﺓ ﻭﺃﺨﺭﻯ ﻏﻴﺭ‬
‫ﻁﻴﺎﺭﺓ )ﻤﻭﺍﺩ ﻤﺘﺒﻘﻴﺔ(‪ .‬ﻭﺍﻟﻤﻭﺍﺩ ﺍﻟﻁﻴﺎﺭﺓ ﻫﻲ ﺍﻟﻔﻭﺭﻤﺎﻟﺩﻫﻴﺩ ﻭﺤﻤﺽ ﺍﻟﻔﻭﺭﻤﻴﻙ ﻭﺘﻡ ﺍﻟﺘﺄﻜﺩ ﻤـﻥ‬
‫ﻫﺫﻩ ﺍﻟﻨﻭﺍﺘﺞ ﺒﺈﺠﺭﺍﺀ ﺍﺨﺘﺒﺎﺭﺍﺕ ﺍﻟﻁﻴﻑ‪ ،‬ﻜﻤﺎ ﺘﻡ ﺘﺤﺩﻴﺩ ﻫﺫﻩ ﺍﻟﻨﻭﺍﺘﺞ ﻜﻤﻴﹰﺎ‪.‬‬
Contents
Page No.
Dedication …………………………………………………………
I
Acknowledgment ………………………………………………….. II
Abstract (English) …………………………………………………. III
Abstract (Arabic) ………………………………………………….. IV
Chapter (I)
1- Introduction……………………………………………………..
1
1-1 Peroxydisulphate oxidation …………………………………
1
1-2 The structure of peroxydisulphate …………………………..
1
1-3 Decomposition of peroxydisulphate ion …………………… 2
1-4 Classification of oxidation by peroxydisulphate…………….
6
1-5 Types of peroxydisulphate oxidation mechanisms ………… 7
1-6 Specific oxidations …………………………………………
1-6-a Peroxydisulphate ion oxidation of some inorganic
10
10
compounds …………………………………………………
1-6-a-I Reaction of thiosulphate with peroxydisulphate ion…... 10
1-6-a-II Reaction of peroxydisulphate ion with manganous ion. 10
1-6-a-III Reaction of peroxydisulphate ion with iodide ion…… 11
1-6-a-IV Reaction of several inorganic ions with S2O8= ………
11
1-6-b Peroxydisulphate oxidation of organic compounds ………
12
1-6-b-I Oxidation of Alcohols ………………………………..
12
1-6-b-II Oxidation of aldhydes ……………………………….
12
1-6-b-III Oxidation of cyclo ketones …………………………
13
1-6-b-IV Oxidation of carboxylic acids ………………………
14
Chapter (II)
Experimental ……………………………………………………..
15
2- Kinetic measurements …………………………………………
15
2-1 Estimation of peroxydisulphate ……………………………
16
2-1-2 The rate laws …………………………………………..
16
2-1-3 Preparation of solutions ………………………………
17
2-2 Thermal decomposition of peroxydisulphate (uncatalysed )... 19
2-2-1 Introduction ……………………………………………..
19
2-2-2 Rate dependence on peroxydisulphate concentration …… 19
Chapter (III)
3-A The Glucose – peroxydisulphate reaction ………………….
3-A-1 Results of kinetic measurement …………………….
3-A-1-1 Rate dependence on peroxydisulphate concentration ..
27
27
27
3-A-1-2 Rate dependence on glucose concentration ………… 36
3-A-1-3 Estimation of rate law at 60ْ C. ………………………. 42
3-A-1-4Effect of temperature on the rate of reaction of glucose 43
3-A-1-5 Calculation of physical parameters …………………...
46
3-A-1-6 Estimation of rate law at higher temperature …………
49
3-A-2 Analysis of the reaction products ………………………..
55
3-A-2-1 Spot tests ……………………………………………..
55
3-A-2-2 Spectrophotometry ……………………………………
56
3-B Maltose – peroxydisulphate reaction ……………………
59
3-B-1 Results of kinetic measurement ………………………
59
3-B-1-1 Rate dependance on peroxydisulphate concentration …
3-B-1-2 Estimation of rate law at 60ْC …………………………
59
66
3-B-1-3 Rate dependance on maltose concentration …………… 66
3-B-1-4 Effect of temperature on rate of reaction of maltose ….. 74
3-B-1-5 Calculation of some physical parameters for maltose
peroxydisulphate reaction ……………………………..
77
3-B-1-6 Analysis of the reaction products ……………………...
78
Chapter (IV)
Discussion and Conclusion ……………………………………….
81
References ………………………………………………………… 85
Chapter (1)
Introduction
1-1 Peroxydisulphate oxidation:
The Peroxydisulphate ion is one of the strongest oxidizing agents
known in aqueous solutions. The standard oxidation- reduction potential for
the reaction:
2 So4
2−
( aq )
2−
→ S 2 o8 + 2e ..........(1.1)
is estimated to be 2.01 volt(1). Reactions involving this ion, however,
are generally slow at ordinary temperatures
(2)
and many peroxydisulphate
oxidations have been studied kinetically. The term “peroxydisulphate” is
used by the chemical abstracts, also the International Union of Pure and
Applied Chemistry (3) recently has recommended the name peroxydisulphate.
The trivial name “persulphate” is also in common use.
1-2 The structure of peroxydisulphate ion:
It has been shown by x- ray analysis of ammonium and calcium
peroxydisulphate that the configuration of this ion is:
Sulphar
Oxygen
The angle SO-O is 128˚, the distance S-O is 1.5A˚, and the distance
O-O is 1.31 A˚. The mid point between central oxygen atoms is the centre of
symmetry. The Raman spectra of sodium and ammonium peroxydisulphate
were investigated by Simon and Richter (4) and their results confirm the
above structure.
In 1910 d’Ans and Friederich(5)found that the crystals of anhydrous
peroxydisulphuric acid are slightly decomposed at about 65 ˚ C. They can be
kept for months at ordinary temperature, but oxygen is slowly evolved. They
are hygroscopic and hydrolysed when dissolved directly in water. At
ordinary temperatures, aqueous solutions of the alkali metal
peroxydisulphate show an appreciable decomposition after a few days; and
with increasing temperature the rate of the decomposition is quite fast.
It was found that (6) crystals of anhydrous peroxydisulphuric acid can
melt with a slight decomposition at about 65˚ C. They are hygroscopic and
easily hydrolysed when dissolved in water forming sulphuric and
permonosulphuric acids.
1-3 Decomposition of peroxydisulphate ion:
The thermal decomposition of peroxydisulphate ion has been the
subject of study of many workers. Morgan and Christ (7) showed that
photochemical decomposition of potassium peroxydisulphate in aqueous
solution proceeds at 30˚ C, which is similar to the thermal decomposition.
The reaction proceeds according to a zero-order law in more concentrated
solution.
K2S2O8 + H 2O → 2KHSO4 + ½O2 ........
(1.2)
Levi and Migliorini (8) found that the decomposition of sodium and
potassium peroxydisulphate, was unimolecular in nature; they also observed
that, their solutions, which were stable at 35˚ C, were decomposed
catalytically by hydrogen and hydroxy ions as well as platinum black, and
lead (9). Since the decomposition of peroxydisulphate produces hydrogen
sulphate ions and it is catalysed by hydrogen ions, it should show autocatalysis, which is however not observed. To account for this, Green and
Mason (10) assumed that the acid sulphate produced from the reaction (1.2)
ionized into metal and hydrogen sulphate ions and no hydrogen ions are
formed to act as catalyst. They suggest the following rate law for the
reaction:
2−
− d { S 2dtO 8 } = k o { S 2 O 8 2 − }......... (1 . 3 )
Where ko is the observed specific rate constant.
Fronaeus and Ostman (11) showed that the rate constant was
independent of ionic strength, but in acid solution there is a negative salt
effect. Kolthoff and Miller (12) studied the decomposition of
peroxydisulphate ion using water enriched with oxygen at different pH
values, they found that in dilute acid solutions, oxygen evolved in the
reaction came from water. This, together with their observation on the
negative salt effect in acid solution enabled them to conclude that the
decomposition occurs by means of two reaction paths:
a/ An uncatalysed reaction in which there is a symmetrical rupture of
the oxygen- oxygen bond of peroxydisulphate ion to form two
.
sulphate free radicals ( S O 4 ) which are removed by reaction with
water according to the following scheme:
S 2O8−2 → 2 S O .4 .......... ..........
.......( 1 .4 )
.
S O .4 + H 2O → HS O 4 + O H ..........(1 .5 )
2 OH . → H 2O + ½ O 2 .................... .(1 .6 )
the hydroxyl group OH is a free radical.
b/ An acid catalysed reaction in which there is an asymmetrical rupture
of the oxygen- sulphur bond of HS2O8 ion to yield sulphur tetraoxide
(SO4) and hydrogen sulphate ion (HSO-4). Sulphur tetraoxide
decomposes to yield oxygen in the following manner:
S 2 O 82 − + H
+
→ HS 2O 8 .........( 1 . 7 )
HS2 O 8 → SO 4 + HS O4 ........( 1 . 8 )
SO
4
→ SO
3
+ ½ O2 ...............( 1 . 9 )
In aqueous media sulphur tetraoxide hydrolyses to give Caro’s acid.
SO4 rate
+ Hconstant
1.10
2O → H
5 ..........
) can
thus be.....(
split
into) two parts, one
The observed
(k2oSO
for the catalysed and the other for the uncatalysed.
k o = k1 + k 2
{H+} .………. (1.11)
The silver ion catalysed decomposition of proxydisulphate ion was
studied by Bawn and Margerison (13). Their study showed that the silver ion
catalysed decomposition is first order in both silver and peroxydisulphate
ions which occurred according to:
.
Ag+ + S2O 82-→ Ag 2++ SO 4 + SO42-............(1.12)
2+
+OH. → + + .............................(
Ag
Ag OH
(1.13)
The following rate law was thus suggested:
d{S 2O82 − }
−
= k 0 {S 2O82 − }..............(1.14)
dt
where
ko = k 1 + k 2 { Ag + }
k1 and k2 being the rate constants for the uncatalysed and silver ion
catalysed reactions.
It has been noted by many workers, that the system Ag+S2O82- is a
much more powerful oxidizing agent than S2O82- alone. Copper ion was
shown to be much less efficient than silver ion as a catalyst for
decomposition; the kinetics of the copper- catalysed reaction were complex
and the reaction occurred by a mechanism different from that found with
silver ion.
In the case of silver ion catalysed oxidations of the reducing agents,
the observed rate law is:
{
}
d S 2 O82−
= k 2 S 2 O82− Ag + …………………………..(1.15)
dt
{
}{
}
The velocity at any instant is proportional to the peroxydisulphate
concentration and to the constant silver ion concentration and is independent
on the reductant concentration. Thus the reactions are essentially
bimolecular and have a common rate determining step as follow:
Ag+ + S2O82− → Ag3+ + 2SO42−...........(1.16)
A mechanism has been postulated involving an equilibrium which is
followed by intermolecular rate determining steps.
2−
8
S2O
−
→ 2SO . 4 (a) ⎯rapid
⎯
⎯→........................(1.17)
−
2SO . 4 + Ag + → 2SO42 − + Ag 3+ (b) ⎯slow
⎯→
⎯ ...(1.18)
The rate of reaction will be given by:
{
}
d S 2 O82−
= k b SO4.−
dt
{
} {Ag }(c)
2
+
…………………(1.19)
From the equilibrium (a)
−
=
K
2
SO
−
Where K is the equilibrium
constant.
.
Substitution of (d) in (c) gives:
{
}
S
d
d S 2O82−
2−
= kb K S 2O8 Ag +
O4
dt
= kobs S 2O82− Ag +
{
{
}{ }
}{ }
………….(1.21)
Where kobs= kbK and hence kb= kobs/K
Similarly, both copper (II) and copper (I) ion catalysed
peroxydisulphate oxidation and the reactive spices is copper (III).
1-4 Classification of oxidation
by peroxydisulphate:
{
For many reducing agents, oxidation by peroxydisulphate does not
proceed at convenient rate at 25˚ C, unless a catalyst is present (14). The most
thoroughly investigated catalyst is silver (I) ion(15), although reactions
involving the copper (II) ion (16) and the iron (II) ion (17) also have been
studied. The rates of reaction catalysed
by these ions are independent of the
}
reductant concentrations but depend on the first power of the
peroxydisulphate and catalyst concentration.
Thus kinetic studies on the oxidation by peroxydisulphate can be
divided into two classes.
a/ Catalysed oxidation.
b/ Uncatalysed oxidation.
The uncatalysed oxidation can be further subdivided into
2
I.
Reactions whose rates depend on the first power of the
peroxydisulphate concentration and are independent of the reducing
agent concentration.
II.
Reactions whose rates depend on the first power of both the
reductant and the oxidant concentration.
2
8
III.
Reactions of intermediate nature, involving fractional powers of
the reactants, changes of order, or auto catalytic characteristics (18).
1.5 Types of peroxydisulphate oxidation mechanisms:
Oxidation mechanism of both inorganic and organic substrates by
peroxydisulphate have been reviewed by Wilmarth and Haim (19).
Peroxydisulphate oxidises a variety of organic compounds, including
alcohol, aldhydes, ketones, phenols, amines and carbohydrates. The majority
of the oxidations proceed via free- redical chain mechanism, but a few are
ionic in character, namely the oxidation of phenols and aromatic amines (20).
1.5.1 Mechanisms for the uncatalyzed Reaction:
Bartlett and Cottman(62) suggested for the uncatalyzed thermal
decomposition the following chain mechanism :
S 2 O82 − ⎯
⎯→ 2SO 4.− …………………………………(1.22)
SO4.− + H 2O ⎯
⎯→ HSO4− + OH . …………………(1.23)
OH . + S 2O82− ⎯
⎯→ HSO4− + 1 O2 + SO4.− ………..(1.24)
SO4.− + OH . ⎯
⎯→
(11)
Ostman
1
2
2
O 2 + HSO4− ………………(1.25)
suggested another mechanism
for the uncatalyed
decomposition.
S 2 O82 − ⎯
⎯→ SO 4.− + SO 42 − ……………………………(1.26)
SO4.− + H 2 O ⎯
⎯→ SO42− + OH . + H + ……………….(1.27)
S 2 O82− + H 2 O ⎯
⎯→ HSO 4− + SO4.− + OH . ……………(1.28)
SO4.− + H 2 O ⎯
⎯→ HSO4− + OH . ………………………...(1.29)
1
2OH . ⎯
⎯→ H 2 O + 2 O ………………………………(1.30)
Livilt
(58
) suggested that the uncatalyzed decomposition of S2O2-8 in
aqueous solution proceeds through hetrolytic cleavage :
...........................………(1.31)
S 2 O82− ⎯
⎯→ SO4.− + SO42− ………………………(1.31)
Further, SO4 reacts with water as follows :
SO4 + H 2 O → SO42− + OH + H + ............................(1.32)
Equation (1.29) postulates retardation of the reaction by added
sulphate ions.
It was observed that in the presence of many reducing substrates the
rate of the reaction is increased. If the primary step is the production of two
sulphate free radicals :
S 2O82− ⎯
⎯→ 2SO4.− ……………………………..(1.33)
followed by the rapid attack of these radicals on the reducing
molecules as suggested by House (59) in the following mechanism :
S O
−k1
⎯2⎯→
−
2 SO 4. ……………………(1.34)
−
k2
SO 4. + H 2O ⎯⎯→
HS O 4 + OH . …..(1.35)
2.
8
2−
−
k3
−
.
−
k4
SO 24− + SO4−+ X ………..(1.37)
X + S 2O82−⎯⎯→
.
.
In this scheme, peroxydisulphate is decomposed by steps(1.34) and
(1.37) leading to the rate expression :
−d
[S O ]= k [S O ]+ k [S O ]
22
2−
8
1
2
2−
8
4
2
2−
8
−
.
[X]
……………………(1.38)
X dtis a substrate ion such as oxalate ion and [X-.] is its free radical ion
concentration. The value of the free radical ion concentration may vary from
reductant to another, giving rise to the variation in the rate constants of the
different systems.
The chain mechanisms are characterized by the following; the
reactions proceed at a faster rate than the spontaneous decomposition of
peroxydisulphate, fractional orders with respect to peroxydisulphate and the
substrate are often encountered and the reactions are often susceptible to
catalysis or inhibition by impurities particularly metal ions and dissolved
oxygen (58).
Reproducible kinetic results are often difficult to obtain due to the use
of solutions containing unknown amounts of oxygen or metal ions.
The possible types of chain mechanisms for peroxydisulphate
oxidation have been classified by Wilmarth and Haim to the dominant
initiation steps, and the relative importance of sulphate radical ions and
hydroxyl radicals in the propagation steps (59).
1.5.2 Mechanisms of oxidation of carbohydrate (RHO):
The oxidation of some carbohydrates such as sucrose , rhamnose and
dextrose were studies by Wood and Walker(60) in the presence of silver ion
as catalyst and found that , concordant velocity constants for the oxidation of
aldoses can be obtained and they concluded that the aldoses were
quantitatively converted to aldonic acids.
Vasudeva
(61)
studied the uncatalysed oxidation of some simple sugar
by peroxydisulphate ion and found that the reaction follows first order in
S2O82- and fractional order in substrate, he proposed the following rate
equation:
[
]
[
R = k 1 S 2O82 − + k 2 S 2
2−
8
] [ substrate ]
(1 39)
Where k1 is the rate constant of the thermal decomposition of
peroxydisulphate alone, k2 being the rate constant of the bimolecular
reaction between substrate and peroxydisulphate ion, he proposed the
following mechanism:
S 2O82−
.
2 SO 4−....
k1
.
SO 4− + H 2O
RHO +
.
.
SO 4−
1
O ....(
2 2
Fast)
........………..(1.42)
.
−
R C = O + HS O 4 ........……..(1.43)
k3
R C = O + H 2O
........……..(1.40)
HSO 4−+ OH . (Slow)........…..(1.41)
k2
2OH . → H 2O +
( Initial step )
k4
RCOOH + H . ........…..(1.44)
O
......(1.46)
.......(1.47)
1.6 Specific oxidations:
(a) Peroxydisulphate ion oxidation of some inorganic compounds:
(i)
Reaction of thiosulphate with peroxydisulphate ion:
Kinetic studies (21,22,23,24) of the reaction between thiosulphate ion
and peroxydisulphate represented by the equation
S2O82- + 2 S2O32-
S2O62- + 2 SO2-4
(1.49)
have shown that the rate is first order in peroxydisulphate
concentration and zero order in thiosulphate concentration over a wide range
of reactant concentration(25). The reaction is very sensitive to traces of
impurities in water and the rate constant is markedly increased by catalytic
amounts of cupric ions. The proposed mechanism leads to the rate law is:
k1
S 2 O8= ⎯⎯→
2 SO4.− ………………………………(1.50)
k2
SO4.− + H 2O ⎯⎯→
HSO4− + OH . …………….(1.51)
OH . + S 2O3= ⎯
⎯→ OH − + S 2O3− …………fast (1.52)
k3
S 2O3− + S 2O8= ⎯⎯→
S 4O6= + SO4= …………….(1.53)
[
]
[
]
− d S 2O82−
= (k1k 2 k3 ) S 2O82− .................................(1.54)
dt
which agrees with experimental rate law.
Reaction of Mn (II) with S2O82- in the presence of Ag (I):
(ii)
In the presence of silver ion the oxidation of Mn (II) by
peroxydisulphate takes place at measurable rate at 25ْ C according to the
equation (26,27).
S2O82- + Mn (II) + 2H2O
Ag(1) 2SO2- + 4 H+
4
( 1. 55)
At higher temperatures and in strong acid, the oxidation can proceed
to the permanganate stage (28, 29).
2Ag(1)
2MnO4 + 10SO4 + 16 H+ …..(1.56)
5 S2O82- + Mn (II) + 8H2O
At intermediate concentrations manganese (II) is formed (30)
2Ag(1)
2 Mn(II) + 2SO4 ……..……. (1.57)
S2O82- + 2Mn (II)
(iii)
2-
Reaction of iodide ion with S2O8:
The stoichometric equation for the reaction between peroxydisulphate
and iodide ion is:
2-
S2O8 + 2I-
I2 + SO2-4 ……………………….(1.58)
The oxidation is usually considered to be bimolecular(31).
The velocity constants are calculated from the experimental
(32,33,34,35,36)
by means of the equation.
results
[
]
[
][ ]
− d S 2O82−
= k 2 S 2O82− I − ..................................(1.59)
dt
The slow oxidation of iodide ion by peroxydisulphate was first
observed by Marshall
(2)
and the first kinetic study was made by Price
(34)
who followed the extent of the reaction by taking samples at known time
intervals determining the amount of iodine produced by titration with
Thiosulphate (37,38).
(iv)
The reactions of several other inorganic ions with S2O82-:-
The reactions of inorganic ions like Cr(III)(39), Cu(I)(40), Ce(II)(44) Am(42) and
T1(43) have been studied, and suitable mechanisms are proposed in each case
which lead to the experimental rate laws.
(b) Peroxydisulphate oxidation of organic compounds:
(i)
Oxidation of Alcohols:
Subbaraman and Santapa
(44)
made a systematic study of the oxidation
of Methanol(45), Ethanol, Butanol, and Isobutanol by peroxydisulphate in
aqueous solution in the presense of air. The orders of reaction with respect
to peroxydisulphate and alcohols were 3/2 and ½ respectively.
Edwards et.al
(46)
, showed that in the absence of the oxygen the
reaction is zero- order with respect to Methanol. They proposed the
following mechanism.
S2O82-
k1
.
.
2S O 4 …………………………….(1.60)
.
S O 4 + CH3OH
k2
CH2OH + HSO4………. (1.61)
CH2OH + S2O82-
k3
HCHO + HSO4- + S O 4 …(1.62)
.
.
2 CH2OH
k4
.
products ……….………..(1.63)
The steady state approximation gives:
[
]
[
1
− d S 2O82−
= [k1 / k 2 ] 2 k3 S 2O82−
dt
]
3
2
....................(1.64)
The rate of oxidation of ethanol by peroxydisulphate was measured (56)
and the reaction rate follows the relation.
[
] [
]
− d S 2O82−
= k S 2O82− ........................................(1.65)
dt
(ii)
Oxidation of Aldehydes:
The kinetics of oxidation of propanaldehyde
(47)
, acetaldehyde
(48,49)
,
and formaldehyde by peroxydisulphate had been studied. The reactions were
found to be first order with respect to both peroxydisulphate, and silver ion
concentrations, and zero order with respect to aldehydes.
Subbaraman and Santappa
(49)
studied the oxidation of the above
aldehydes in the absence of silver ion, at low aldehyde concentration. The
following rate law was observed.
3
− d S 2 O 82 −
= k S 2 O 82 − 2 [R − CHO
dt
[
]
[
]
]1 2 ..........
.......... ....( 1 . 66 )
The oxidations of propanaldehyde, acetaldehyde, and acetone by
peroxydisulphate were studied. (50) It was found that these reactions followed
first order with respect to peroxydisulphate and zero order with respect to
the above substrates. A radical mechanism was proposed as follows:
.
k1
S2O82-
2S O 4
.
S O 4 + H2O
.
.
k2
…………………………(1.67)
.
.
.
HS O 4 +OH…………………(1.68)
X + OH
k3
X+ S2O82-
k4
products ………………………(1.70)
k5
products ……………………(1.71)
.
X + SO 4
X + H2O………………………(1.69)
Where X is the substrate. On applying steady state treatment to the
above mechanism, the following rate law is deduced.
[
]
[
]
− d S 2O82−
= k 0 S 2O82− .......................................(1.72)
dt
This agrees very well with the experimental rate law.
(iii)
Oxidation of Cyclo ketones:
Kinetics of silver ion catalysed oxidation of cyclo pentanone, cyclo
hexanone, and cyclo heptanone by peroxydisulphate ion has been studied
(51)
. These reactions were first order in peroxydisulphate, and silver (I)
ion, and zero order in cyclo ketones.
(iv)
Oxidation of Carboxylic acids:
The uncatalysed reactions of peroxydisulphate with that of carboxylic
acid such as malonic
(52)
, citric
(53)
and glyxolic
(54)
were studied free radical
mechanism have been proposed which lead to the experimental rate laws.
It has been found that these reactions follow first order with respect to
peroxydisulphate and zero order in the acids concentration.
The kinetics of copper (II) catlaysed oxidation of formic acid
(57)
and
oxalic acid by peroxydisulphate has been studied and a free radical
mechanism is postulated in each case.
Chapter (11)
Experimental
This work includes a study of the redox reactions between pottasium
peroxydisulphate and some simple sugars [Glucose and Maltose], in the
absence of catalysts.
Two aspects of the study are involved :
1.A kinetic study.
2.An investigation of the reaction products.
2-1 Kinetic measurements :
The aim of carrying out the kinetic study was to establish a rate law
for each reaction of the two substrates.
This required that sets of runs to be carried out in which the
concentration of one of the reactants was changed while keeping the
concentration of the other reactant constant, this kinetic run was carried out
as follows :
The solution of the substrate in deionized water was placed in the
thermostat for 20 minutes before the experiment was started the required
volume of peroxydisulphate solution was then transferred to reaction flask
(three – necked round bottom 250 ml) .
Special care was taken for cleaning the reaction vessel in a uniform
manner. Each time ,the flask was cleaned with chromic acid then washed
with deionized water, steamed and left to dry at 100 ˚C, the shape of the
reaction vessel was found to effect the rate of the reaction .
2-1 Estimation of Peroxydisulphate :
The iodometric method employed for the analysis of unreacted
peroxydisulphate was a modification of the method used by Bartett and
Cotman and Rosin(62). In this method, sodium bicarbonate and sulphuric acid
were used to maintain pH value of 7.1 to 7.2. The solution to be analysed
for peroxydisulphate consisted of the following components :
(I)
5ml of the reaction mixture.
(II)
5ml of 4% sodium bicarbonate solution.
(III) one ml of (IN) sulphuric acid solution.
(IV) 5ml of approximately 30% pottasium lodide solution.
The librated iodine was titred against sodium thiosulphate using starch
as indicator. The strength of sodium thiosulphate solution was of little
importance in different kinetic runs but through the same run its composition
was the same .
The persence of unreacted substrate was found not to interfere with
this method of estimation.
The addition of a reducing substrate to peroxydisulphate at higher
temperature 60˚C increased the rate of its decomposition due to the redox
reaction .
2-1-1 The Rate Law :
The measurement of the rates and rate constant aimed at getting at the
end an expression for the rate law of each reaction investigated. This
required calculation of the rates of reaction at different reaction
concentrations. A common feature presisted throughout the systems
investigated that rates of different reactions increased linearly with increase
in peroxydisulphate concentration when the substrate concentration was
kept constant. This showed that the order of all reactions was unity in
peroxydisulphate and this fact seems to be common for many other
peroxydisulphate reactions . It will be seen in the next chapter that the order
of the reaction in substrates concentration is fractional and this does not
seem uncommon in peroxydisulphate reactions .
2-1-2 Preparation of Solutions :
The redox reaction of peroxydisulphate with organic compounds is
highly susceptible to impurities in the solution and that often leads to errotic
values. King and Stein back(64)showed that the variations observed in their
results were due to chloride ions in the distilled water they used. This
required that the water used in our work to be deionized water . This water
was used to prepare different solutions and as medium for the kinetic runs :
The following paragraphs summerize the procedure followed in
preparing the solutions :
(I)
Glucose :
The required weight of glucose (Hopkin and Williams reagent grade)
was dissolved in the approriate amount of deionized water to give (0.1)
molar solution.
(II)
Maltose :
(0.1) molar solution was prepared by dissolving the required wieght of
Maltose (Hopkin and Williams reagent grade) in the approporiate amount of
deionized water.
Each of the above sugars was prepared one day before kinetic
measurments for optical equilibrium . Every kinetic run required a new
sugar solution.
(III) Potassium peroxydisulphate :
The required amount of (K2S2O8) (Riedle-Dehenag–Seelze Hannover.
Analar grade) was dissolved in the approporiate amount of deionized water
to give (0.1) molar solution . It was used within two days from the time of its
prepration and was kept at room temperature in the dark.
(IV) Sodium thiosulphate :
The required concentration of Na2S2O3 was prepared by dissolving a
calculated amount of Na2S2O3 (Metlab,labrotary reagent)in the approporiate
amount of hot water to give (0.1) molar solution . The concentration of
Na2S2O3 was adjusted to peroxydisulphate conc. of every kinetic run. It was
kept as a stock for 4-5 days.
(V)
Potassium Iodide :
An approporiate amount of (KI) (B.D.H,A.R) was dissolved in the
required volume of deionized water to give 30% (w/v) solution of (KI) , it
was freshly prepared.
(VI) Sodium bicarbonate :
A 4% (w/v) solution of NaHCO3 (Merck., A.R ) in the required
amount of deionized water was kept as stock.
(VII) Sulphuric acid :
(IN) H2SO4 acid was prepared by diluting the calculated volume of
98% H2SO4 acid (Meltlab . Labrotary reagent ) with deionized water . It was
kept as stock..
(VIII) Starch :
One gram of soluble starch (B.D.H,A.R) was mixed with a little water
to form a paste then poured into 100ml of boiled water with constant stirring
until cooled. It was freshly prepared every time
2-2 Thermal Decomposition of Peroxydisulphate (uncatalysed) :
2-2-1 Introduction
A number of runs were made on thermal decomposition of
peroxydisulphate in the absence of any reducing agent.
The purpose of doing such studies varied with different workers .
There
is
general
agreement
that
the
thermal
decomposition
of
peroxydisulphate obeys a first order rate law(63) . But there are, however
some disagreements as to the manner in which such reaction takes place.
Also to calculate the order of the bimolecular reaction this required
that few sets of experiments to be carried out on the thermal decomposition
of peroxydisulphate under similar conditions to those carried out at different
concentration of peroxydisulphate at 60˚C as follow:
2-2-2 Rate dependence on peroxydisulphate concentration :
Table No (2.1) :
Time
Titre
M
∆M
∆t
Sec
ml
mole / l
mole / l
sec
106 R
mole /l/Sec
105 ko
sec
[S2 O=8]
0
9.90
0.01000
3600
9.65
0.00989
0.00011
3600
0.030
2.55
7200
9.40
0.00988
0.00018
3600
0.027
0.52
10800
9.20
0.00987
0.00009
3600
0.025
0.53
14400
9.15
0.00980
0.00007
3600
0.019
0.51
18000
8.90
0.00975
0.00005
3600
0.013
0.57
21600
8.50
0.00971
0.00004
3600
0.011
0.53
[S2 O=8] = 0.01 mole / l
Temperature = 60˚ C
Table No (2.2) :
Time
Titre
M
Sec
ml
mole / l
∆M
∆t
106 R
105 ko
mole / l
sec
mole/ l/Sec
sec-1
[S2 O=8]
0
9.75
0.0200
3600
8.80
0.0182
0.0018
3600
0.500
3.78
7200
8.50
0.0176
0.0006
3600
0.166
1.29
10800
8.30
0.0174
0.0002
3600
0.055
0.62
14400
7.60
0.0165
0.00009
3600
0.025
1.99
18000
7.45
0.0158
0.00007
3600
0.019
0.96
21600
7.10
0.0141
0.00005
3600
0.013
0.98
[S2O=8] = 0.02 mole/l
Temperature 600C
Table No (2.3) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
0
11.50
0.04000
3600
11.15
0.03844
0.00156
3600
0.433
3.03
7200
10.80
0.0373
0.00145
3600
0.402
0.68
10800
10.60
0.0366
0.00120
3600
0.333
0.65
14400
10.25
0.0331
0.00091
3600
0.252
0.60
18000
9.60
0.0248
0.00064
3600
0.177
0.59
21600
8.90
0.0216
0.00037
3600
0.102
0.61
25200
7.50
0.0204
0.00015
3600
0.041
0.67
[S2O=8] = 0.04 mole / l
Temperature = 60˚C
Table No (2.4) :
Time
Titre
M
∆M
∆t
Sec
ml
mole / l
mole / l
sec
106 R
mole /l/Sec
105 ko
sec-1
[S2 O=8]
0
13.90
0.06000
3600
13.55
0.0550
0.0050
3600
1.388
2.10
7200
13.10
0.0521
0.0029
3600
0.805
0.56
10800
12.65
0.0512
0.0009
3600
0.270
0.46
14400
12.30
0.0510
0.0002
3600
0.250
0.61
18000
11.85
0.0490
0.0001
3600
0.220
0.50
27000
8.50
0.0489
0.00009
3600
0.190
0.63
20600
7.65
0.0488
0.00005
3600
0.133
0.58
[S2O=8] = 0.06 mole / l
Temp. = 60˚ C
Table No (2.5) :
Time
Titre
M
∆M
∆t
Sec
ml
mole / l
mole / l
sec
106 R
mole /l/Sec
105 ko
sec-1
[S2 O=8]
0
11.45
0.1000
3600
11.20
0.0973
0.0027
3600
0.750
2.51
7200
10.70
0.0961
0.0012
3600
0.333
0.60
10800
10.30
0.0951
0.0010
3600
0.277
0.49
14400
9.95
0.0940
0.0009
3600
0.250
0.46
18000
8.20
0.0733
0.0007
3600
0.190
0.50
21600
7.60
0.0628
0.0005
3600
0.133
0.62
15200
6.40
0.0537
0.0001
3600
0.021
0.56
[S2O=8] = 0.1mole/l
Temp = 60˚C
Table No (2.6) :
Values of m, R and ko from table (2-5)
M [S2O=8] mole/l
106 R mole/l/sec
105 kosec-1
0.0973
0.750
2.51
0.0961
0.333
0.60
0.0951
0.277
0.49
0.0940
0.250
0.46
0.0733
0.190
0.50
0.0628
0.133
0.62
0.0537
0.021
0.56
Figure (2.1) represents the plot of the rate (R) in mole/ L/sec against
the peroxydisulphate mean concentration . The linear plot statifies the
relationship;
R = k0 [S2O=8] …………………………………. (2.1)
Where k0 is the observed rate constant of the reaction equation
(2.1) shows that the reaction is first order in peroxydisulphate
concentration the slope of the plot Figure (2.1) has a value equal to k0
= 0.55х 10-5 sec-1
Figure (2.2) represents the plot of (k0) observed constant a
gainst the peroxydisulphate mean concentration , and this plot is linear
and paralleled to the concentration axis.
− d
[S
2
O
dt
=
8
]=
k
0
[S
2
O
=
8
]
Chapter (III)
3.A The Glucose – peroxydisulphate reaction
3.A.1 Results of the kinetic measurement
3.A.1.1 Rate dependence on peroxydisulphate concentration :
The reduction of peroxydisulphate by glucose is very slow at ordinary
temperature. The half-life of peroxydisulphate – glucose reaction at room
temperature is nearly one month. It was found that at 60˚C the reaction
proceeds at measurable rate and for this reason the temperature range (60˚ to
80˚C) was chosen in this study. A set of runs were carried out in which the
concentration of glucose was kept constant at 2 x 10-2 mole/l and the
concentration of peroxydisulphate was varied from 1 x 10-2 to 10 x 10-2
mole/l.
Table No (3.A.1)
Time
Titre
M
∆M
∆t
106 R
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
105 ko
[S2 O=8]
0
4.75
0.0100
3600
4.60
0.00968
0.00032
3600
0.088
5.9
7200
4.55
0.00955
0.00013
3600
0.036
1.2
10800
4.50
0.00947
0.00008
3600
0.022
1.5
14400
4.25
0.00942
0.00005
3600
0.013
3.2
18000
4.20
0.00940
0.00002
3600
0.005
3.6
21.600
4.05
0.0939
0.00001
3600
0.002
3.5
[S2O=8] = 0.01 mole / l
Temp. 60˚ C
Table No (3.A.2) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
6.15
0
9.50
0.0200
3600
9.45
0.0196
0.0004
3600
1.11
0.71
7200
9.25
0.0193
0.0003
3600
0.83
2.10
10800
8.85
0.0181
0.0002
3600
0.55
4.50
14400
8.15
0.0190
0.0001
3600
0.27
4.80
18000
8.10
0.0189
0.00005
3600
0.13
3.52
27000
8.00
0.0188
0.00003
3600
0.11
2.11
[S2O=8] = 0.02 mole/l
[ Substrate ] = 0.02 mole/l
Na2S2O3 = 0.01 mole /l
Temp. = 60˚ C
Table No ( 3.A.3) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
0
14.7
0.030
3600
14.3
0.024
0.006
3600
1.666
6.92
7200
13.9
0.021
0.003
3600
0.833
1.90
10800
13.5
0.019
0.002
3600
0.555
2.11
14400
12.4
0.018
0.001
3600
0.277
3.79
18000
12.3
0.0175
0.0005
3600
0.138
4.30
27000
12.1
0.0171
0.0004
3600
1.111
2.61
[S2O=8] = 0.03 mole /l
Temp. = 60˚C
Table No ( 3.A.4) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
0
39.5
0.050
3600
38.0
0.045
0.205
3600
1.388
7.1
7200
37.2
0.041
0.003
3600
0.833
1.7
10800
35.5
0.039
0.002
3600
0.555
3.1
14400
34.1
0.037
0.001
3600
0.277
3.5
18000
33.6
0.036
0.0005
3600
0.138
4.2
27000
32.9
0.040
0.0003
3600
0.083
4.5
[S2O=8] = 0.05 mole / l
Temp. = 60˚C
Table No (3.A.5 ) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
0
41.10
0.070
3600
39.05
0.063
0.007
3600
1.944
6.84
7200
37.50
0.060
0.003
3600
0.833
3.1
10800
35.95
0.058
0.002
3600
0.550
4.1
14400
33.60
0.057
0.001
3600
0.277
2.2
18000
30.20
0.0565
0.0005
3600
0.138
4.1
21600
28.10
0.0561
0.0004
3600
0.111
2.8
[S2O=8] = 0.07 mole / l
Temp. = 60˚C
Table No ( 3.A.6) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole / l
sec
mole/l/Sec
sec-1
[S2 O=8]
0
45.2
0.080
3600
43.9
0.071
0.009
3600
2.500
0.999
7200
42.05
0.065
0.006
3600
1.666
1.230
10800
41.3
0.062
0.003
3600
0.833
1.888
14400
40.1
0.060
0.002
3600
0.550
2.311
18000
38.2
0.059
0.001
3600
0.277
3.33
21600
37.5
0.0595
0.0005
3600
0.138
2.87
[S2O=8] = 0.08 mole/l
Temp. = 60˚C
Table No (3.A.7 ) :
The values of m, R, ko from table (3-A.5)
102 [S2O=8]mole/l
106 R mole/l/Sec
105 kosec-1
0.063
1.944
6.84
0.060
0.833
3.1
0.058
0.550
4.1
0.057
0.277
2.2
0.0565
0.138
4.1
0.0561
0.111
2.8
3.A.1.2 : Rate Dependence on Glucose conc. :
Table (3.A.8) to (3.A.12) include experimental data for the run to
investigate the dependence of the rate on Glucose concentration at 60˚C.
A set of runs were carried out in which the concentration of
peroxydisulphate was kept constant at 2х 10-2 mole /l and the concentration
of Glucose was varied from 2х10-2 to 10х10-2 mole/l.
Table No (3.A.8) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
6.9
0.020
3600
5.6
0.0162
0.0038
3600
1.05
15.1
7200
4.7
0.0136
0.0026
3600
0.722
12.1
10800
4.2
0.0121
0.0015
3600
0.416
12.6
14400
3.2
0.0108
0.0013
3600
0.361
13.1
18000
3.0
0.0098
0.0010
3600
0.277
12.8
21600
2.9
0.0090
0.0008
3600
0.22
12.9
25200
2.5
0.0084
0.0006
3600
0.166
12.5
[Glu] = 0.02 mole/l
[S2O=8] = 0.02 mole/l
Temp. = 60˚C
Table No (3.A.9) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
7.0
0.020
3600
6.3
0.0161
0.0039
3600
1.08
16.9
7200
6.2
0.0135
0.0026
3600
0.722
11.7
10800
5.5
0.0117
0.0018
3600
0.660
11.9
14400
4.9
0.0099
0.0015
3600
0.505
12.1
18000
3.75
0.0088
0.0011
3600
0.305
12.3
21600
3.10
0.0078
0.0010
3600
0.277
11.7
[Glu] = 0.04 mole/l
[S2O=8] = 0.02 mole/l
Temp. = 60˚C
Table No (3.A.10) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
7.5
0.0200
3600
7.35
0.0154
0.0046
3600
1.27
15.1
7200
6.9
0.0121
0.0033
3600
0.916
12.3
10800
6.2
0.0092
0.0029
3600
0.805
12.4
14400
5.55
0.0081
0.0011
3600
0.305
13.3
18000
5.3
0.0080
0.0010
3600
0.277
13.6
21600
4.5
0.0071
0.0009
3600
0.250
12.5
[Glu] = 0.06 mole/l
[S2O=8] = 0.02 mole/l
Temp. = 60˚C
Table No (3.A.11) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
8.10
0.020
3600
7.5
0.0141
0.0059
3600
1.63
16.14
7200
7.15
0.0138
0.0003
3600
0.83
13.39
10800
6.30
0.0136
0.0002
3600
0.55
14.72
14400
6.05
0.0135
0.0001
3600
0.27
13.05
18000
5.40
0.0134
0.00008
3600
0.22
14.98
21600
5.20
0.0133
0.00005
3600
0.13
13.16
[Glu] = 0.08 mole/l
[S2O=8] = 0.02 mole/l
Temp. = 60˚C
Table No (3.A.12) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
10.05
0.020
3600
9.80
0.0127
0.0073
3600
2.02
17.25
7200
9.15
0.0072
0.0055
3600
1.52
12.9
10800
8.50
0.0032
0.0040
3600
1.11
13.8
14400
7.50
0.0022
0.0010
3600
0.277
13.7
18000
7.80
0.0015
0.0007
3600
0.194
14.1
21600
6.90
0.0009
0.0006
3600
0.166
14.5
[Glu] = 0.1 mole/l
[S2O=8] = 0.02 mole / l
Temp. = 60˚C
Table No (3.A.13) :
Summary of tables (3.A.8) to (3.A.12) using the mean values of M , R and
ko ;
102 [Glu]
102 [S2O=8]
106 R
105 Ko
mole/l
mole/l
mole/l/Sec
Sec-1
2
0.0141
1.63
16.14
4
0.0138
0.83
13.39
6
0.0136
0.55
14.72
8
0.0135
0.27
13.05
10
0.0134
0.22
14.98
3.A.1.3 Estimation of the order and rate law for Glucose and
peroxydisulphate at 60˚C
R = k1 [S2O=8] + k2 [S2O=8][substrate]n
Where R represent the rate of the bimolecular reaction.
k1 represent the rate constant of the decomposition of
peroxydisulaphate alone.
k2 represent the rate constant of the bimolecular reaction.
n is the order of the reaction.
By integration of the above equation we can calculate n:-
n=
n=
Log R- log k1 – log [S2O=8]
Log k2 + Log [S2O=8] + Log [sub]
Log 1.04 - log 0.225 – log [0.02]
Log 2.24 + Log [0.02] + Log [0.02]
n = 0.25
∴
the rate expressions:-
R = k1 [S2O=8] + k2 [S2O=8] [ sub] 0.25
3.A.1.4 : Effect of Temperature at the Rate of Reaction of Glucose :
Table No (3.A.14) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
0
10.25
0.02
3600
9.40
0.018
0.002
3600
5.55
12.66
7200
8.60
0.017
0.001
3600
2.77
8.34
10800
8.25
0.0165
0.0005
3600
1.38
9.20
14400
7.90
0.0162
0.0003
3600
0.38
9.55
18000
7.20
0.0160
0.0002
3600
0.55
9.97
21600
6.80
0.0159
0.0001
3600
0.27
10.11
[S2O=8] = 0.02 mole/l
[Glu]
= 0.02 mole /l
Temp. = 65˚C
Table No (3.A.15) :
Time
Titre
[S2O=8]
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
0
10.70
0.0200
3600
8.55
0.01740
0.0026
3600
7.22
15.61
7200
5.20
0.01720
0.0020
3600
5.55
13.84
10800
3.60
0.0154
0.0018
3600
5.00
12.31
14400
2.05
0.0139
0.0015
3600
4.66
12.70
18000
1.80
0.0128
0.0011
3600
3.05
12.44
21600
1.20
0.0120
0.0008
3600
2.22
11.98
[S2O=8] = [Glu] = 0.02 mole/l
Temp. = 70˚C
Table No (3.A.16) :
Time
Titre
105 k0
Time
Titre
105 ko
Sec
ml
Sec–1 (75˚C)
Sec
ml
sec-1(80˚C)
0
11.20
0
18.00
1800
10.05
41.45
900
16.15
62.15
3600
6.10
37.91
1800
14.00
58.24
5400
3.60
37.36
2700
10.70
59.19
7200
1.90
35.99
3600
7.34
59.99
9000
0.60
34.06
5400
2.47
59.78
12600
0.25
36.12
7000
1.50
58.36
[S2O=8] = [Glu] = 0.02 mole/l
Temp. = 75˚, 80˚C
Table No (3.A.17) :
TC˚
T K˚
103 1/T K-1
105 koSec-1
105 log koSec-1
60˚
333
3.003
2.40
0.38
65˚
338
2.958
9.27
0.96
70˚
343
2.915
12.81
1.10
75˚
348
2.875
35.79
1.55
80˚
353
2.833
59.11
1.77
3.A.1.5 : Calculation of Physical Parameters [ Ea , A, ∆ G and ∆ S]
for Glucose – Peroxydisulphate reaction:
1/ Activation energy (Ea) in J/mole ;
Ea = - Slope х 2.303 х 8.314 J / mole
II/ Frequency factor (A) in Sec-1 ;
A = 2.97 х 10-5 Sec-1 [from the intercept of y-axis]
III/ Free Energy Change (∆ G) kJ/mole ;
( at different temperatures 65˚ , 70˚ , 75˚ , 80˚C )
∆ G = ∆ Ea - T∆ S kJ/ mole
IV/ Entropy change (∆ S) J / K
(at different temperatures 65˚ , 70˚ , 75˚ , 80˚C )
∆ S = 2.303 R ( log A – log RT(Nh) J/ K
Whore R/N is Boltzman’s gas constant = 1.3806 х 10-23 Jdeg-1 and h is
Plank's constant= 6.626 х 10-34 Js
3.A.1.6 : Estimation of Rate Law at higher Temperature
Table No (3.A.19): Rate dependence on Glucose conc. :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
10.25
0.0200
1800
8.70
0.0610
0.01390
1800
2.8
14.20
5400
7.10
0.0042
0.011
3600
2.5
13.66
9000
4.90
0.0031
0.009
3600
2.05
12.99
14400
3.05
0.0027
0.004
5400
1.11
12.78
16200
2.30
0.00025
0.002
1800
0.55
12.51
18000
1.00
0.00024
0.001
1800
0.277
13.11
[S2O=8] = 0.02 mole/l
[Glu] = 0.02 mole / l
Temp = 70˚C
Table No (3.A.20) :
Time
Titre
M[S2O=8]
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
0
10.15
0.020
1800
7.98
0.0154
0.0046
1800
2.55
16.31
5400
6.55
0.0114
0.0040
1800
2.22
13.22
9000
5.20
0.0083
0.0031
1800
1.72
13.01
14400
4.00
0.0058
0.0025
1800
1.38
13.31
16200
3.10
0.0038
0.0020
1800
1.11
12.74
18000
1.60
0.0028
0.0010
1800
0.55
12.69
[Glu] = 0.04 mole/l
[S2O=8] = 0.02 mole/l ,
Temp. = 70˚C
Table No ( 3.A.21) :
Time
Titre
M[S2O=8]
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
0
10.50
0.0200
1800
6.80
0.01351
0.00649
1800
3.61
20.7
3600
5.65
0.0116
0.00191
1800
1.06
17.6
5400
4.70
0.0101
0.0015
1800
0.833
16.8
7200
3.50
0.0091
0.0010
1800
0.55
15.7
9000
2.80
0.0083
0.0008
1800
0.44
15.9
10800
1.35
0.0079
0.0004
1800
0.222
15.2
[Glu] = 0.06 mole/l
[S2O=8] = 0.02 mole/l
Temp.70˚C
Table No (3.A.22) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
12.90
0.0200
1800
10.55
0.0119
0.0081
1800
4.50
15.8
3600
7.80
0.0055
0.0064
1800
3.55
13.7
5400
5.05
0.0025
0.005
1800
2.77
11.11
7200
3.20
0.0005
0.003
1800
1.66
12.03
9000
1.60
0.0003
0.002
1800
1.11
12.11
10800
1.15
0.0002
0.001
1800
0.55
11.69
[Glu] = 0.08 mole/l
[S2O=8] = 0.02 mole/l
Temp = 70˚C
Table No (3.A.23) :
Summary of tables (3.A.19) to (3.A.23) R dependence on [Glu] at
70˚C
102 m
106R
105 ko Sec-1
[Glu]0.25
[Glu]0
mole/l/Sec
2
1.23
12.87
0.457
4
1.50
13.35
0.525
6
1.51
16.50
0.569
8
2.35
12.30
0.603
3.A.2 Analysis of the reaction products :
A mixture of glucose (20g) and potassium peroxydisulphate (20g) in
deionized water (250ml) was refluxed for 4 days at 70 C
until
peroxydisulphate was used up. A reflux condenser was used while cold
water (10˚C) circulated throughout the heating so as to prevent the escaping
of the volatile products, at the end of the reflux the reaction mixture, was
evaporated to dryness under reduced pressure. The resulting distillate was
identified as a mixture of formaldehyde and formic acid by spot tests and
stiochiometry as follows:
3.A.2.1 Spot tests :
(A) Formaldehyde :
To ½ ml of the distillate was added 2ml of concentrated sulphuric
acid and 2 ml sodium chromotropate in a test tube. The mixure was heated
for 10 minutes on a water bath at 60˚ C, aviolet colour appeared, indicating
the presence of formaldhyde , to confirm this, its 2,4dinitrophenyl hydrazone
derivative was prepared by adding 2,4 dinitrophenyl hydrazine to fraction of
distillate , when recrystalized from ethanol gave a m.P.(166-167)0C (65).
(B) Formic acid :
To 2ml of the resulting distillate , 1ml of HgCl2 solution was added
and the mixture was heated .A white precipitate turning to grey-black
was obtained indicating the presence of formic acid. To confirm the
presence of formic acid, its amide derivative was prepared by adding 2ml
of the distillate to 2g of pCl3 in a porcelain basin. The contents were
mixed until the mixture became liquid then 20ml of NH4OH solution was
added drop by drop. The mixture was then filtered and washed with cold
water.
The derivative was recrystallized from water and dried , gave a m.p. (195
– 196)0C(65).
Titemetric estimation of formic acid : It was determined by titration,
against 0.01m NaOH using phenolphthalin as indicator.
3.A.2.2 Spectrophotometry :
The estimation of unreacted substrates in a reaction mixture and the
estimation of formaldhyde required a spectrophotometric determination of
these compounds by means of coloured complexes which had an absorbance
at a known wavelength . The complexes formed obeyed Beer’s law.
Spectrophotometric estimation of formaldhyde :
A solution
containing very small amount of formaldhyde when
heated with sulphuric and chromotropic acid gives rise to purple coloured
complex with maximum absorbance at 580 nm. The absorbances of different
solutions were found to be linear in formaldhyde concentration.
To one ml portions of solution a,b,c, d and e (taken) in different
flasks, 10ml chromotropic acid was added and allowed to stand in water bath
(90˚-98˚C) for 30 minutes. The solution was then left to cool to room
temperature, and the absorbances of the solution were determined using
perkin-Elmer UV/vis spectrophotometer model 550S . The absorbances of
the standar solution of formaldhyde are given in table (3.A.24) and
graphically represented by figure (3.A.10). Unknown amounts of
formaldhyde were determined with the help of figure (3.A.10) .
Table No (3.A.24) :
Solutions
10-4concentration mole/l
Absorbance at 580 nm
a
5
0.451
b
10
0.978
c
15
1.396
d
20
2.031
e
25
2.317
3.B The Maltose – Peroxydisulphate reaction :
3.B.1 Result of the kinetic Measurement :
3.B.1.1 Rate depend on peroxydisulphate concentration; Table
(3.B.1) to(3.B.5) include the data of kinetic run for the redox reaction at 60˚c
in which the concentration of Maltose was constant at 2х 10-2 mole/L and
that for peroxydisulphate was varied from 1х10-2 to 10 х 10-2 mole/L.
Table No (3.B.1) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
5.5
0.01
3600
5.05
0.0092
0.0008
3600
2.22
8.4
7200
4.85
0.0088
0.0004
3600
1.11
5.7
10800
4.45
0.0085
0.0003
3600
0.83
5.99
14400
4.25
0.0083
0.0002
3600
0.55
6.15
18000
4.00
0.0082
0.0001
3600
0.27
6.31
21600
3.20
0.0073
0.00009
3600
0.25
5.8
[S2O=8] = 0.01mol/l
Temp.= 60˚C
[Maltose] = 0.02 mole /l (in all cases)
Table No (3.B.2) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
9.60
0.020
3600
9.45
0.0196
0.0004
3600
1.11
9.5
7200
9.10
0.0193
0.0003
3600
0.83
6.2
10800
8.50
0.0191
0.0002
3600
0.55
6.3
14400
8.20
0.0090
0.0001
3600
0.27
6.9
18000
7.95
0.0089
0.00008
3600
0.22
7.1
21600
5.60
0.0088
0.00005
3600
0.13
6.2
[S2O=8] = 0.02 mole/l
Temp 60˚C
Table No (3.B.3) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
16.05
0.050
3600
15.7
0.043
0.007
3600
1.94
9.9
7200
15.2
0.037
0.006
3600
1.66
7.2
10800
14.8
0.033
0.004
3600
1.11
7.3
14400
14.5
0.031
0.002
3600
0.55
7.1
18000
14.15
0.030
0.001
3600
0.277
6.1
21600
15.20
0.029
0.0009
3600
0.211
6.4
[s2o8=] =.0.05mole/l
Temp.=600C
Table No (3.B.4) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
27.1
0.070
3600
25.9
0.060
0.010
3600
2.77
10.11
7200
22.7
0.052
0.008
3600
2.22
6.03
10800
21.3
0.046
0.006
3600
1.66
5.85
14400
20.65
0.043
0.003
3600
0.83
6.75
18000
20.05
0.041
0.002
3600
0.55
7.10
21600
19.55
0.040
0.001
3600
0.277
6.5
[S2O=8] = 0.07 mole/l
Temp. = 60˚C
Table No (3.B.5) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
35.15
0.1000
3600
34.70
0.090
0.010
3600
2.77
11.50
7200
33.90
0.081
0.009
3600
2.50
6.21
10800
32.05
0.074
0.007
3600
1.94
5.80
14400
30.85
0.069
0.005
3600
1.38
6.01
18000
29.10
0.067
0.002
3600
0.55
7.31
21600
28.25
0.061
0.0009
3600
0.52
6.42
[S2O=8] = 0.1 mole/ l
Temp = 60˚C
Table No (3.B.6.):
The values of M, R, ko from table (3-B.5)
M [S2O=8] mole/L
10-6 R mole/L/Sec
10-5 KoSec-1
0.090
2.77
11.50
0.081
2.50
6.21
0.074
1.94
5.80
0.069
1.38
6.01
0.067
0.55
7.31
0.061
0.52
6.42
3.A.1.3
Estimation of the order and rate law for Maltose and
peroxydisulphate at 60˚C
R = k1 [S2O=8] + k2 [S2O=8][substrate]n
Where R represent the rate of the bimolecular reaction.
k1 represent the rate constant of the decomposition of
peroxydisulaphate alone.
k2 represent the rate constant of the bimolecular reaction.
n is the order of the reaction.
By integration of the above equation we can calculate n:-
n=
n=
Log R- log k1 – log [S2O=8]
Log k2 + Log [S2O=8] + Log [sub]
Log 1.61 - log 0.225 – log [0.02]
Log 2.87 + Log [0.02] + Log [0.02]
n = 0.25
∴
the rate expressions:-
R = k1 [S2O=8] + k2 [S2O=8] [ sub] 0.25
3.B.1.3 Rate dependence on Maltose concentration :
Table (3.B.7) to (3.B.11) include experimental data for the run to
investigate the dependence of the rate on maltose concentration at 60 ˚ C. A
set of runs were carried out in which the concentration of peroxydisulphate
was kept constant at 2х10-2 mole/l and the concentration of maltose was
varied from 5х10-3 to 8х10-2 mole/l.
Table No (3.B.7) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/Ll/Sec
sec-1
[S2O=8]
0
14.50
0.0200
3600
13.70
0.01780
0.0022
3600
0.611
6.71
7200
12.05
0.01580
0.00200
3600
0.555
2.23
10800
11.60
0.0140
0.0018
3600
0.500
3.31
14400
10.00
0.0124
0.00160
3600
0.440
3.14
18000
9.70
0.0128
0.00153
3600
0.425
3.42
21600
8.55
0.0098
0.0009
3600
0.275
4.21
[S2O=8] = 0.02 mole / l
[maltose] = 5х10-3 mole/L
Temp.=600C
Table No (3.B.8) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
25.7
0.0200
3600
22.00
0.0161
0.0039
3600
1.08
5.12
7200
20.10
0.0126
0.0035
3600
0.972
3.21
10800
19.25
0.0096
0.0030
3600
0.75
3.15
14400
18.00
0.0069
0.0027
3600
0.58
2.42
18000
17.45
0.0048
0.0021
3600
0.416
2.10
21600
15.80
0.0033
0.0015
3600
0.311
2.56
[S2O=8] = 0.02 mole/l
[maltose] = 0.01 mole/l
Temp. = 60˚C
Table No (3.B.9) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
16.75
0.020
3600
15.05
0.0186
0.0014
3600
0.38
6.92
7200
13.10
0.0174
0.0012
3600
0.33
3.89
10800
11.00
0.0164
0.0010
3600
0.277
3.81
14400
9.20
0.0084
0.0008
3600
0.222
4.11
5.60
0.0034
0.0005
3600
0.123
3.69
3.60
0.0014
0.0002
0.055
3.58
18000
21600
3600
[Maltose] = 4х10-2 mole/l
[S2O=8] = 2 х 10-2 mole/l
Temp. = 60˚C
Table No (3.B.10) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
15.50
0.020
3600
12.95
0.0115
0.0085
3600
2.30
8.12
7200
10.00
0.0035
0.0080
3600
2.20
5.50
10800
8.15
0.0025
0.006
3600
1.60
5.31
14400
7.00
0.0020
0.005
3600
1.30
4.97
18000
5.20
0.0015
0.003
3600
0.80
4.51
21600
4.60
0.0005
0.001
3600
0.20
5.11
[maltose] = 5х10-2 mole/l
[S2O=8] = 2 х 10-2 mole/l
Temp. = 60˚C
Table No (3.B.11) :
Time
Titre
M
∆M
∆t
106 R
105 ko
Sec
ml
mole / l
mole/ l
sec
mole/l/Sec
sec-1
[S2O=8]
0
17.00
0.0200
3600
15.15
0.0165
0.0035
3600
9.11
7.87
7200
12.40
0.0145
0.0020
3600
5.55
4.21
10800
9.10
0.0135
0.0010
3600
2.77
4.15
14400
6.25
0.0126
0.0009
3600
2.11
4.06
18000
3.05
0.0119
0.0007
3600
1.25
3.59
21600
2.50
0.0115
0.0004
3600
1.01
4.12
[maltose] = 8х10-2 mole/l
[S2O=8] 2 х10-2 mole/l
Temp. = 60˚C
Table No (3.B.12) :
Summary of tables (3.B.7) to (3.B.11)
103[maltose]mole/l
106 R mole/l
105 koSec-1
[Maltose]0.25
5
0.467
3.33
0.3465
10
0.680
2.68
0.3981
40
0.231
3.91
0.5251
50
1.40
4.89
0.5492
80
3.63
4.07
0.6034
3.B.1.4 Effect of Temperature at the Rate of Reaction of Maltose :
Table No (3.B.13) :
Time
Titre
105ko (60˚C)
Time
Titre
105 ko
Sec
ml
Sec-1
Sec
ml
(70˚C)Sec-1
0
15.50
7.11
0
10.80
3600
14.15
5.23
3600
8.75
24.10
7200
12.20
5.12
7200
6.50
20.32
10800
7.95
4.99
10800
3.10
19.54
14400
4.60
4.81
14400
1.60
19.13
18000
3.55
5.15
18000
0.50
19.61
mean
4.91
20.22
Table No (3.B.14) :
Time
Titre[S2O=8] 105ko (75˚C)
Sec-1
Time
Titre[S2O=8]
105 ko
Sec
ml
(80˚C)Sec-1
0
9.10
Sec
ml
0
9.80
1800
6.75
35.71
900
7.90
50.87
3600
4.20
33.12
1800
4.35
47.23
5400
2.10
34.41
2700
2.20
44.11
7200
1.05
35.28
3600
1.35
45.02
10800
0.50
34.11
5400
1.00
46.10
mean
34.21
45.75
Table No (3.B.15) :
TCo
10 T Ko
103 I/T K-1
105 koSec-1
Log ko Sec-1
60
333
3.003
4.91
0.69
70
343
2.915
20.22
1.30
75
348
2.875
34.21
1.53
80
353
2.833
45.75
1.66
3.B.2 Analysis of the reaction products :
A mixture of Maltose (20g) and potassium peroxydisulphate (20g) in
deionized water (250ml) was refluxed for 4 days at 70 C, until
peroxydisulphate was used up. A reflux condenser was used while cold
water (10˚C) circulated throughout the heating so as to prevent the escaping
of the volatile products, at the end of the reflux the reaction mixure, was eva
porated to dryness under reduce pressure. The resulting distilate was
identified as a mixture of formaldehyde and formic acid by spot tests and
stiochiometry as follow:
3.B.2.1 Spot tests :
Formaldehyde: The same as in (3.A.2.1) the derivative gave a.m.p (1650C).
Formic Acid: The same as in (3.A.2.1) the derivative gave a.m.p (1960C).
3.B.2.2 Spectrophotometry:
The estimation of unreacted substartes in
reaction
estimation
mixture
and
the
of
formaldhyde
required
a
spectrophotometric determination of these compounds by means of coloured
complexes which had an absorbance at a known wavelength. The complexes
formed obeyed Beer’s law.
Spectro photometric estimation of formaldhyde:
A solution containing very small amount of formaldhyde when heated
with sulphuric and chromotropic acid gives rise to purple coloured complex
with maximum absorbance at 580 nm. The absorbances of different solution
were found to be linear in formaldhyde concentration.
To one ml portions of solution a,b,c,d and e (taken) in different flasks,
10ml chromotropic acid was added and allowed to stand in water bath (900 –
980C) for 30 minutes. The solution was then left to cool to room
temperature, and the absorbances of the solution were determined using
Perkin-Elmer UV/ vis spectrophotometer model 550S. The absorbances of
the stander solution of formaldhyde are given in table (3.B.17) and
graphically represented by figure (3.B.6). Unknown amounts of formaldhyde
were determined with the help of figure (3.B.6).
Table No (3.B.17):
10-4 concentration
Absorbance at 580
mole/L
nm
a
5
0.131
b
10
0.250
c
15
0.416
d
20
0.552
e
25
0.725
Solutions
Proposed mechanism and rate law:
It was showed that the Glucose and Maltose-Peroxydisulphate
reaction can be represented by two term rate law, one of the thermal
decemposition of peroxydisulphate alone and the other for the bimolecular
reaction, also the reaction has a characterstics of chain reactions, so three
paths exist in this reactions;
Path (I) can be represented by the following scheme:
−
k1
S 2 O8= ⎯⎯→
2 SO4. ……………………………………..(3.1)
−
k2
SO4. + H 2 O ⎯⎯→
HSO4− + OH . ………………………..(3.2)
2OH . ⎯
⎯→ H 2O + 12 O2
…………………………….(3.3)
Path (II) gives rise to aldonic and aldaric acids, the formation of
aldonic acid takes place by the oxidation of the terminal alcoholic group of
Glucose or Maltose as follows:
.−
4
.
RCH 2 OH + SO ⎯⎯→ R C HOH + HSO4− ………………(3.4)
k3
.
.
−
k4
R C HOH + S 2O8= ⎯⎯→
R C HO + HSO4− + SO4. ………...(3.5)
.
−
4
.
R C HO + SO ⎯⎯→ R − C O + HSO4− ……………………(3.6)
k5
.
k6
R − C O + H 2O ⎯⎯→
R C OOH + H . ………………….…(3.7)
.
.
k7
H + O H ⎯⎯→
H 2O ……………………………………….(3.8)
The formation of aldaric acid takes place by the oxidation of the
second terminal alcoholic group as follows:
.
.
k8
RCH 2 OH + O H ⎯⎯→
R C HOH + H 2 O ……………….(3.9)
.
.
−
k9
R C HOH + S 2O8= ⎯⎯→
R C HO + HSO4− + SO4. …….(3.10)
.
.
−
k10
R C HOH + SO4. ⎯⎯→
R C HO + HSO4− ……………(3.11)
Path (III) give rise to formaldehyde and formic acid in which the
process of C-C oxidative cleavage proceeds until all the sugar molecule is
completely converted to formaldhyde and formic acid.
At the steady state the rate of disappearance of peroxydisulphate ion
can be represnted by the following equestion:
{
}
− d S 2 O8=
⎧ .
⎫
= k1 S 2 O8= + (k 4 + k 9 )⎨R C HOH ⎬ S 2 O8= …………(3.12)
dt
⎩
⎭
{
}
{
}
.
The rate of change in the concentration of the radicals ⎧⎨ R C HOH ⎫⎬ ,
⎩
⎭
{SO } , {OH }, {H }, ⎧⎨⎩R C = O⎫⎬⎭ and ⎧⎨⎩R C HO⎫⎬⎭ with time equals zero as follow:
.−
4
.
.
.
.
⎧ .
⎫
− d ⎨ R C HOH ⎬
⎩
⎭ = k ⎧ R C. HOH ⎫ S O = + k ⎧ R C. HOH ⎫ S O = +
⎬ 2 8
⎬ 2 8
4⎨
9⎨
dt
⎩
⎭
⎩
⎭
{
}
{
{ }
}
{ }
⎧ .
⎫
⎧ .⎫
…(3.13)
k10 ⎨ R C HOH ⎬ SO4. − k 3 {Substrate} SO4. − k8 {Substrate}⎨O H ⎬ =........
0
⎩
⎭
⎩
⎭
{ } = k {SO }{H
−
− d SO 4.
dt
2
{SO }− k {S
.−
4
{
.−
4
1
2
}
2 O } + k 3 {Substrate
{
}{SO 4.
}
−
}+ k
{
{
}
⎫
⎧ .
⎫
⎧ .
O 8= − k 4 ⎨ R C HOH ⎬ S 2 O 8= − k 9 ⎨ R C HOH ⎬ S 2 O 8= = 0
⎭
⎩
⎭
⎩
}
⎫⎧
⎧
d OH .
⎫
⎧ . ⎫⎧ . ⎫
⎧ . ⎫
= k 7 ⎨ H ⎬⎨O H ⎬ + k 8 {Substrate}⎨O H ⎬ − k 2 ⎨SO4.− ⎬⎨ H 2 O ⎬ = 0
dt
⎩ ⎭⎩
⎭
⎩
⎭
⎭
⎭⎩
⎩
{ }
d H.
⎫
⎧ . ⎫⎧ . ⎫
⎧ . ⎫⎧
7
= k 7 ⎨ H ⎬⎨O H ⎬ − k 6 ⎨ R C O ⎬⎨ H 2 O ⎬ = 0
dt
⎩ ⎭⎩
⎭
⎩
⎭⎩
⎭
}
⎫
⎧ .
⎫
⎧ .
.−
R
C
HO
SO
k
+
⎬
⎨
5
4
10 ⎨ R C HOH ⎬
⎭
⎩
⎭
⎩
.......…(3.14)
......…(3.15)
.....…(3.16)
.
⎫
⎧
d ⎨R − C O⎬
⎭ = k ⎧ R − C. O ⎫{H O} − k ⎧ R C. HO ⎫⎧SO .− ⎫ = 0
⎩
⎬⎨ 4 ⎬
⎬ 2
6⎨
5⎨
dt
⎭⎩
⎩
⎭
⎩
⎭
{
}
.
.
⎫
d RC . HO
⎫⎧
⎧
⎫
⎧
= k 5 ⎨ R CHO ⎬ SO4.− + k 4 ⎨ R CHO H ⎬⎨S 2 O8= ⎬
dt
⎭⎩
⎩
⎭
⎩
⎭
{
....…(3.17)
}
.
.
⎫
⎧
⎫⎧
⎧
⎫
- k 9 ⎨ R CHO H ⎬⎨S 2 O8= ⎬ − k10 ⎨ R CHO H ⎬ SO4.− = 0
⎩
⎭⎩
⎩
⎭
⎭
{
}
...…(3.18)
After rearangement of these equations we get:
k3 k 4 {substrate}+ k 2 k 4 {H 2 O}+ k3 k 9 {substrate}
⎧ .
⎫
⎨ R C HOH ⎬ = 4
2k 4 k10 + 2k 9 k10
⎩
⎭
.…(3.20)
Because of the great abudance of H2O we can neglect it.
.
If we substitute for the value of ⎧⎨ R C HOH ⎫⎬ in equation (3.12) for the
⎭
⎩
rate law we get:
{
}
{
{
}
{
− d S 2 O8=
0.25
= k1 S 2 O8= + (k 4 + k 9 ) S 2 O8= {substrate}
dt
}
{
}
− d S 2O8=
0.25
= k1 S 2 O8= + k 2 S 2 O8= {substrate}
dt
}
{
}
…(3.21)
…(3.22)
Where k2 = k4 + k9
k1 represent the rate constant for the thermal decomposition of the
peroxydisulphate alone.
k2 represent the rate constant for bimolecular reaction.
Substarte stand for Glucose or Maltose.
Chapter (1V)
Discussion and Conclusion
The present work deals with a systematic study of the redox reaction
of some simple sugars (Glucose and Maltose) with peroxydisulphate under
uncatalyzed conditions.
It was found that the reactions of these substrates with
peroxydisulphate do not proceed at room temperature. Therefore, the
temperature range chosen was (60-800 C).
It
was
decomposes
found
even
substrates.
paths;
decomposition
path
(ii)
reaction
of
in
There
path
(i)
of
that
the
are,
peroxydisulphate
absence
two
represents
main
of
reaction
the
thermal
peroxydisulphate,
represents
the
peroxydisulphate
these
while
bimolecular
with
reducing
substrate. In the present work, no attempt
was made to estimate the extent of path(i)
and
path (ii).
Two aspects of the study were involved:a) Kinetic study:
i)
The progress of each reaction was followed by examining the
concentration of potassium peroxydisulphate in the reaction mixture
at different time intervals by iodometric titration method. The
reaction rate was measured as a function of concentration of
potassium peroxydisulphate (oxidant) and Glucose and Maltose
(substrates).
ii)
The Rate Law: In all reactions involving peroxydisulphate, the
initial step is its decomposition into two sulphate free radicals. The
sulphate radicals can either react with water to give oxygen radicals,
or with the organic substrates to give their free radicals which further
attack peroxydisulphate and enhance its decomposition. A systematic
kinetic study of the reactions showed that these followed first order in
peroxydisulphate concentration, and fractional order (0.25) in sugar
concentration. Vasudeva
(15)
showed that the reaction of aldoses with
peroxydisulphate was first order in peroxydisulphate satisfying the
equation:
{
=
}
R = k 0 S 2 O8 .....................................(4.1)
When the concentration of peroxydisulphate was gradually increased,
the plot of the mean rates against peroxydisulphate was linear passing
through the origin at zero peroxydisulphate concentration. The slopes of
the lines give the values of the observed rate constants (k0) according to
equation (4.1). The initial rate constans were usually not included in the
calculation of the mean (k0) for each experiment, these higher rate
constants in the initial stages of the reaction can be due to the calatlysing
role which the surface of the reaction vessel plays between
peroxydisulphate and hydroxyl radicals. With the progress of reaction the
concentration of products becomes appreciable and if these products
occupy a part of the surface by adsorption this will reduce the catalysing
effect of the surface. After a while from the start of the reaction a state of
equilibrium between product adsorption and surface catalysing occur and
the observed rate constant (k0) tends to be constant.
When the concentration of substrate was raised to a power of 0.25
the plots became linear.
The rate of the reaction may be expressed by the following equation.
R = R − + R = .......... .......... .......... .........( 4.2)
where R represent the observed rate of substrate peroxydisulphate
R − represent
reaction.
the
rate
of
thermal
decomposition
of
=
peroxydisulphate alone. R represent the rate of bimolecular reaction
between substrate and peroxydisulphate.
We therefore, suggest the following rate law for these reactions.
{
=
} {
R = k1 S 2O8 + k 2 S 2O8
=
}{substrate}
0.25
................................(4.3)
The above equation when integrated gave the following relationship:
k0 = k1 + k 2 {substrate} ........................................(4.4)
0.25
where
k0
is
the
observed
rate
constant
of
the
substrate
peroxydisulphate reactions, k1 is the rate constant of the thermal
decomposition of peroxydisulphate alone. k2 is the rate constant of the
bimolecular reaction between substrate and peroxydisulphate.
iii)
Temperature Effect:The variation of the observed rate constant depended on the change of
temperature (60-80ْ C). From the value of k0 at different temperatures, the
energy of activation (∆Ea ) , the frequency factor (A), the entropy change
(∆S ) and free energy change (∆G ) were calculated.
iv)
Proposed Mechanisms:-
The mechanisms proposed explain the experimental rate law for both
glucose and maltose reactions with peroxydisulphate, and account for the
simple decomposition of peroxydisulphate is free radical mechanism. The
rate of the reaction obtained is the same as the experimental rate law. These
mechanism can be generalized as follows:
k1
S2O8-2
k2
SO-.
4 + H2O
.
k3
X + OH
.
X + S2O8-2 k4
.
-.
X + SO4
2 SO-40
.
HSO-. + OH
4
.
X + H2O
products
products
where X = substrate.
b) Investigation of the reactions products:
The analysis of the reaction products revealed the presence of
formaldhyde and formic acid in volatile fractions and a non-volatile fraction,
this fraction consisted mainly of potassium hydrogen sulphate, unreacted
substrate and other products of oxidation. Volatile products were confirmed
by spot tests, and the amount of formaldhyde was determined
spectrophotometrically using Perkin-Elmer spetrophotometer model 550S,
while those of formic acid were determined titmetricaly.
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n
Curriculum Vitae
Personal Information:
Name:
Nasma Dafalla Eljack
Date of Birth:
18/7/1974
Place of Birth:
Karkoj – Sinnar State
Nationality:
Sudanese
Religion:
Muslim
Social status :
Married
Languages:
Arabic (Fluent)
English (Written and Spoken)
Address:
Khartoum North – Tel: 85338317
Mobile: 0912961987
Academic Qualifications:- Sudanese higher secondary school certificate (1992).
- B.Sc. honour degree – first class with university price for best
performance in chemistry – faculty of science – University of
Khartoum (1998).
- M.Sc. in Physical chemistry (The Oxidation of two simple sugars
by peroxydisulphate (2004).
Training and courses:- Course of computer science at U of K computer centre.
- Course of Internet at U of K Internet centre.
Table No (3.A.18) : Some physical parameters for glucose – peroxy
disulphate reaction :
Physical
Activation
Frequency
Entropy change (∆S±)J/K
Free energy (±∆G) kJ/mole at
Parameter
energy(±Ea)
Factor
at different temperatures
different temperatures
kJ/mole
(A)Sec-1
43.15
2.97 х 10-5
found
65˚
70˚
75˚
80˚
15.21 12.85 12.46 12.23
60˚
70˚
75˚
80˚ C
37.94
38.6
38.7
38.91
3.B.1.5 Some physical parameters for maltose – peroxy disulphate
reaction :
Table No (3.B.16) :
ysical
Activation Energy
Frequency
Free Energy change
Entropy change (±∆S) in k
ameter
(Ea±) in kJ/mole
factor (A) in
(±∆G) in kJ /mole at
at different temperature
Sec-1
different temperatures
ound
60˚
27.63
2.91 х10-5
70˚
75˚
80˚
13.21 12.79 12.35 12.06
60˚
70˚
75˚
23.24
23.26
23.34
2