KINETICS AND MECHANISMS OF PEROXODISULPHATE OXIDATION OF CERTAIN AMINO ACIDS IN ACIDIC AQUEOUS SOLUTION

A THESIS SUBMITTED FOR THE DEGREE

OF PH.D IN CHEMISTRY

BY

MOHAMMED AWAD ALI KHALID B.SC (HONOURS) JULY 2000

DEPARTMENT OF CHEMISTRY FACULTY OF SCIENCE

UNIVERSITY OF KHARTOUM

November 2006

I

CONTENTS DEDICATION …………………………………………….…... VIII

ACKNOWLEDGEMENTS …………………………….… IX

ABSTRACT /English……………………………….................... X

ABSTRACT /Arabic……………………………………………. XI

Chapter (1) Introduction 1.1 Peroxodisulphate as an oxidizing agent…………………...…… 1

1.2 Aqueous decomposition of Peroxodisulphate ………….…… 2

1.3 The first-order uncatalyzed peroxodisulphate oxidation 6

1.4 Oxidation of amino acids in vitro ……………………………… 9

1.5 Oxidative decarboxylation of amino acids……………….……. 20

Chapter (2) Experimental 2.1 Preparation of solutions………………………………………… 22

2.2 Apparatus……………………………………………………….. 24

2.3 Washing procedure……………………………………………... 24

2.4 Identification methods for the reaction products……………... 25

2.5 Estimation of the residual peroxodisulphate……………….. 25

2.6 Kinetic measurements…………………………………………... 26

Chapter (3) Results 3.1 Preliminary Studies…………………………………………….. 29

3.1.1 Stoichiometrey……………………………………………... 29

3.1.2 Catalytic metal ions in deionized water…………………... 32

3.2 Kinetic Results…………………………………………………... 35

3.2.1 Effect of Varying Peroxodisulphate and amino acid concentrations on the Reaction Rate…………………..…

35

(3.2.1.1) Alanine-Peroxodisulphate system…………… 36

(3.2.1.2) Argnine-Peroxodisulphate system………… 39

(3.2.1.3) Asparagine -Peroxodisulphate system………… 42

II

(3.2.1.4) Aspartic acid -Peroxodisulphate system……… 45

(3.2.1.5) Cysteine -Peroxodisulphate system…………… 48

(3.2.1.6) Glutamic acid -Peroxodisulphate system……… 51

(3.2.1.7) Glutamine -Peroxodisulphate system………… 54

(3.2.1.8) Glycine -Peroxodisulphate system………… 57

(3.2.1.9) Histidine -Peroxodisulphate system………… 60

(3.2.1.10) Leucine -Peroxodisulphate system…………… 63

(3.2.1.11) Lysine -Peroxodisulphate system…………… 66

(3.2.1.12) Metheonine -Peroxodisulphate system……… 69

(3.2.1.13) Phenylalanine -Peroxodisulphate system…… 72

(3.2.1.14) Proline -Peroxodisulphate system……… 75

(3.2.1.15) Serine -Peroxodisulphate system………… 78

(3.2.1.16) Threonine-Peroxodisulphate system……… 81

(3.2.1.17) Tyrosine -Peroxodisulphate system………… 84

(3.2.1.18) Valine -Peroxodisulphate system…………… 87

3.2.2 Effect of Temperature…………………………..……….. 90

3.2.3 Test for free radicals……………………………..………. 103

3.2.4 Effect of solvent composition…………………….……… 103

3.2.5 Effect of Catalysts……………………………………..…. 108

3.2.6 Effect of ionic strength…………………………….…….. 109

3.3 Identification of reaction products……………………….…….. 113

Chapter (4) Discussion References…………………………………………….………….. 165

III

APPENDICES APPENDIX (A.1) …………………………………………………… 135 APPENDIX (A.2)……………………………………………………. 135

APPENDIX (B.1)……………………………………………………. 136

APPENDIX (B.2)……………………………………………………. 136

APPENDIX (C.1)……………………………………………………. 137

APPENDIX (C.2)……………………………………………………. 137

APPENDIX (D.1) …………………………………………………… 138

APPENDIX (D.2) …………………………………………………… 138

APPENDIX (E.1) …………………………………………………… 139

APPENDIX (E.2) …………………………………………………… 139

APPENDIX (F.1) …………………………………………………… 140

APPENDIX (F.2) …………………………………………………… 140

APPENDIX (G.1) …………………………………………………… 141

APPENDIX (G.2) …………………………………………………… 141

APPENDIX (H.1) …………………………………………………… 142

APPENDIX (H.2) …………………………………………………… 142

APPENDIX (I.1) ……………………………………………………. 143

APPENDIX (I.2) ……………………………………………………. 143

APPENDIX (J.1) ……………………………………………………. 144

APPENDIX (J.2) ……………………………………………………. 144

APPENDIX (K.1) …………………………………………………… 145

APPENDIX (K.2) …………………………………………………… 145

APPENDIX (L.1) …………………………………………………… 146

APPENDIX (L.2) …………………………………………………… 146

APPENDIX (M.1) ………………………………………………….. 147

IV

APPENDIX (M.2) ………………………………………………….. 147

APPENDIX (N.1) …………………………………………………… 148

APPENDIX (N.2) …………………………………………………… 148

APPENDIX (O.1) …………………………………………………… 149

APPENDIX (O.2) …………………………………………………… 149

APPENDIX (P.1) …………………………………………………… 150

APPENDIX (P.2) …………………………………………………… 150

APPENDIX (Q.1) …………………………………………………… 151

APPENDIX (Q.2) …………………………………………………… 151

APPENDIX (R.1) …………………………………………………… 152

APPENDIX (R.2) …………………………………………………… 152

APPENDIX (S.1) ……………………………………………………. 153

APPENDIX (S.2) ……………………………………………………. 153

APPENDIX (S.3) ……………………………………………………. 154

APPENDIX (S.4) ……………………………………………………. 154

APPENDIX (S.5) …………………………………………………… 155

APPENDIX (S.6) ……………………………………………………. 155

APPENDIX (S.7) ……………………………………………………. 156

APPENDIX (S.8) ……………………………………………………. 156

APPENDIX (S.9) ……………………………………………………. 157

APPENDIX (T.1) …………………………………………………… 158

APPENDIX (T.2) …………………………………………………… 158

APPENDIX (T.3) …………………………………………………… 159

APPENDIX (T.4) …………………………………………………… 159

APPENDIX (U.1) …………………………………………………… 160

APPENDIX (U.2) …………………………………………………… 160

V

APPENDIX (U.3) …………………………………………………… 161

APPENDIX (U.4) …………………………………………………… 161

APPENDIX (V.1) …………………………………………………… 162

APPENDIX (V.2) …………………………………………………… 162

APPENDIX (V.3) …………………………………………………… 163

APPENDIX (V.4) …………………………………………………… 163

APPENDIX (W) …………………………………………………….. 164

VI

LIST OF ABBREVIATIONS

S2O82- Peroxodisulphate

AA Amino acid

Sch Schiff Base

Ala Alanine

Arg Arginine

Asn Asparagine

Asp Aspartic acid

Cys Cysteine

Glu Glutamic acid

Gln Glutamine

Gly Glycine

His Histidine

Leu Leucine

Lys Lysine

Met Methionine

Phe Phenylalanine

Pro Proline

Ser Serine

Thr Threonine

Tyr Tyrosine

Val Valine

Ǻ Angstrom

T Temperature

kobs Observed Rate Constant

VII

E Activation energy

Ř Universal Gas Constant

R Rate of reaction

m/l Mole per Liter

w/v Weight per volume

µ Ionic strength

µL Microlitre

ppm Part Per Million

∆H Enthalpy change

∆S Entropy change

∆G Free energy change

DL Dextro-Levo (Racemic Mixture)

L Levo-

A.R Analytical Reagent

VIII

DEDICATION

TO

MY

MOTHER, *****

FATHER, *****

BROTHERS, *****

AND

FRIENDS *****

WITH

ALL MY LOVE

MOHAMMED

IX

Acknowledgements All praise is due to '' ALLAH", without whose help and guidance this work

would not have been fulfilled.

I would like to express my sincere gratitude to my supervisor Prof. Ali Mohammed Kheir

For his keen guidance and continuous encouragement, thoughtful guidance, unlimited support as well as giving me such precious assistance and feedback throughout the development of my thesis. Thanks and deepest gratitude are extended to

Dr Mohammed Abd Algaphar Osman

for his invaluable advice, comments, and directions. My gratitude also goes

to each member of Chemistry Department's staff for their invaluable

assistance.

I am so grateful to my sponsor, the Deutscher Akademischer Austausch Dienst E.V. Programme (DAAD), for providing this invaluable opportunity to follow my MSc degree and after upgrading this work to Ph.D at the University of Khartoum, Faculty of Science, Department of Chemistry. Heartfelt and sincere thanks are extended to my brothers, all friends and colleagues at Khartoum University for their continuous encouragement.

Mohammed Awad Ali Khalid University of Khartoum

November 2006

X

ABSTRACT Kinetics and mechanism of oxidation of eighteen α-amino acids by

peroxodisulphate ion have been studied in an aqueous acidic medium

(sulphuric acid) at the temperature range 60-800 C. The rate shows first order

dependence on peroxodisulphate concentration, and zero order dependence

on amino acid concentration.

The rate law observed is: –d[S2O82-]/dt = kobs [S2O8

2-][Amino acid]0.

An autocatalytic effect has been observed in amino acids oxidation due to

formation of Schiff base between the formed aldehyde and parent amino

acid. A decrease in dielectric constant of the medium by adding acetic acid

(5-15% v/v) results in decreasing the rate in all cases studied. Reactions

were carried out at different temperatures (60-800C) and the thermodynamic

parameters viz. activation energy, free energy change, entropy change, and

enthalpy change have been calculated. The rate of amino acids oxidation is

greater in presence of mixture of Ag+ and Cu2+ than in presence of only Ag+

or Cu2+. At higher concentrations of the added neutral salt, the logarithm of

the rate constant is linearly interrelated to the square root of the ionic

strength. The reaction mechanism involving the formation of free radicals

viz. amino acid free radical, sulfate ion free radical, Schiff base free radical,

hydroxyl free radical, has been suggested. The steady state hypothesis was

applied to suggested mechanism. This is found in agreement with the

experimental results.

XI

الخالصةبواسطة ايونات α - من نوعا امينياض ثمانية عشر حمأآسدةتمت دراسة حرآية وميكانيكية تفاعل

درجة مئوية 60ئى محمض بحمض الكبرتيك عند درجة حرارة فى وسط ما بيراآسوثنائى الكبريتات

بالنسبة لترآيز ايونات األولى الرتبة التفاعل منأن الحسابات أآدت . درجة مئوية80إلى

آما من الرتبة الصفرية بالنسبة لترآيز الحمض االمينىآذلك فان التفاعل بيراآسوثنائى الكبريتات و

يلي آما هوهد السرعة المشاقانون معدل أنثبت

–d[S2O82-]/dt = kobs [S2O8

2-][Amino acid]0

أثناءيف تبعا لتكون قواعد ش امينية وذلك أحماض لعدة تحفيزا ذاتيا للتفاعل قد حدثأن أتضحوقد

.المتبقي والحمض االمينى الناتجالتفاعل بين االلدهيد

مقادير من حمض الخل إضافةعند بالنسبة لوسط التفاعل الكهربيعزل قيم ثابت الإنقاص آذلك فان

. تمت دراستهاالتي االمينية األحماض لكل ةبالنسب التفاعل سرعة معدلإنقاص إلىؤدى ي

آمكنقد و درجة مئوية80 درجة مئوية إلى 60 منحرارةال درجات عند دراسة حرآية التفاعل تمت

فيالتغير , في الطاقة الحرةالتغير ( األخرىناميكية حساب طاقة التنشيط وبقية القيم الثيرمودي

.)نتروبيا قيمة االفي و التغير الحراريالمحتوى

+Ag الفضة وجود خليط من ايوناتفي أسرع االمينية األحماض أآسدة معدل أن الدراسة أآدتقد و

.منفردة يوناتاال هذه وجود حالةفي هو عليه جتمعة ممام+Cu2 النحاسو

تنتج اريثم ثابت المعدل غنسبة لترآيز عالى من محلول ملح متعادل من آبريتات البوتاسيوم فان لوبال

. للقوة االيونية الجذر التربيعىقابل مالبياني لرسمل عالقة خطية عنه

الجذر الحر تم افتراض ميكانيكية للتفاعل تحتوى على تكوين جذور حرة اثناء التفاعل وهى هذا وقد

اليون الجذر الحراعد شيف و و لقالجذر الحر, ن الكبريتاتو اليالجذر الحر, ينىللحمض االم

. مع النتائج التجريبيةتتفق أنهاوقد وجد الهيدروآسيل

1

1. Introduction 1.1 Peroxodisulphate as an oxidizing agent

Peroxodisulphates are the most chemically active of the peroxygens, with

great utility in a variety of chemical processes. The term Peroxydisulphate is

used by chemical abstract although the international union of pure and applied

chemistry (IUPAC) has recommended the name Peroxodisulphate(1), the trivial

names, Persulphate, peroxydisulfate and peroxodisulfate were used in the

literature. The structure of Peroxodisulphate ion is the subject of several

investigations. Its structure was established as a result of X-ray analysis of

ammonium and cesium Peroxodisulphate(2) which showed four oxygen atoms

arranged in approximately tetrahedral fashion for each sulfur atom.

The Peroxodisulphate group consists of two sulfate groups linked by a covalent

bond between two oxygen atoms, the distance between these two oxygen atoms

is 1.46 Ả, the distance S-O is 1.50 Ả, while S-O-O inter bond angle is 1280 and

the axis of symmetry is pass through the mid-point between the central oxygen

atoms. This was approved by the results of Raman spectra of sodium and

ammonium peroxodisulphate(3).

S

OO

O O OO

O O

SAxis of symmetry

Figure ︵1.1 ︶Structure of peroxodisulphate group. It was found that crystals of anhydrous peroxydisulphuric acid [H2S2O8] melt

with a slight decomposition at about 650C. It is hygroscopic and easily

hydrolyzed when dissolved in water forming sulphuric acid and permonosul-

phuric acid:

H2S2O8 + H2O H2SO4 H2SO5+

Peroxydisulphuric acid is also considered as an oxidizing agent, peroxodisul-

phate ion, S2O82-, is an excellent and versatile oxidant for a variety of organic

2

and inorganic compounds. The standard oxidation-reduction potential for the

following reaction was estimated to be –2.01 volts(4).

SO4 S2O82-

aq 2e-2-2 +

Marshal observed that reactions involving peroxodisulphate ion, generally are

slow at room temperature.

1.2 Aqueous decomposition of Peroxodisulphate

The following equations represent the hydrolysis of peroxodisulphate ion in

neutral solution,

S2O8-2-

2+ H2O HSO4 O21/2+

alkaline solution,

H2O HSO4 O21/22-

S2O8 + +-

dilute acid solution,

H2O HSO4 H2O22-

S2O8 + +-2 2

and concentrated acidic solution, respectively.

H2O H2SO5 SO42- 2-

S2O8 + +

The rate of decomposition at any pH value is adequately expressed by the first

order rate low (5).

-d[S2O82-]/dt = kobs[S2O8

2-]

The rate of the aqueous thermal decomposition of peroxodisulphate ion can be

catalyzed by hydrogen ion(6), and the values of activation energy can be

calculated from Arrhenius equation:

k = A exp(-E/ŘT)

Where k is the rate constant, A is the frequency factor, E is the activation

energy, Ř is the universal gas constant and T is the absolute temperature.

By using this equation, the values of activation energy were calculated to be

33.3, 32.5 and 32.1 Kcal.mol-1 at pH values 13.0, 9.5 and 8.0 respectively, in

alkaline solution (k1)(6,7) and in neutral solution (kobs.) to be 28.39 and 30.8(8.9),

3

but in acid solution (k2) to be 26.0 at pH=1 and 27.7 at pH=1.7(9.10). It was

concluded that the activation energy increased with increasing pH of solution.

Bawn and Magerison(8) calculated the value of activation energy, it was 23.39

Kcal.mol-1, from which they found the frequency factor to be 7.4x1013 sec-1.

Kolthoff and Miller(7) studied the kinetics of the decomposition in water

labled with oxygen18 at different pH values, and found that in acid solution

0.5M HClO4, all oxygen produced came from peroxodisulphate, but in alkaline

solution 0.1M NaOH, the oxygen came from water. This observation led them

to postulate different mechanisms for the hydrogen catalyzed and uncatalyzed

decompositions. This observed from minimizing the energy of activation in acid

solution. The first mechanism is the hydrogen ion independent reaction

S2O8 SO4

H2O

O2

.

18

2

1/2

2- .

-

-

-.SO4+ HSO4 + OH

OH.

2 H2O+

︵1 ︶︵2 ︶

︵3 ︶

1818

1818

Which showed that the oxygen produced came from water. Equations (1), (2)

and (3) represent the simplest unimolcular decomposition of peroxodisulphate

ion in aqueous solution. Bartlett and Cotman(10) have pointed out, however, that

a chain mechanism also can give rise to the observed kinetics:

S2O8 SO4

H2O

O2+

.2

1/2

2- .

-

-

-.SO4+ HSO4 + OH

OH.

︵1 ︶︵2 ︶

︵3 ︶

2-S2O8 SO4

.-HSO4

-+ +

SO4.-

+ .OH -HSO4+1/2O2 ︵4 ︶

Thus equations (1), (2), (3) and (4) lead to first order kinetics, as step (4)

accounts for chain termination, rather than step (3) in the previous mechanism.

Bartlett and Cotman(10) rejected the chain decomposition, though they

considered it in some detail, on the grounds that the non-chain process was the

simplest; but from consideration of more recent evidence a reviewer believes

4

that a chain mechanism is better suited to explain the results of

peroxodisulphate oxidation. An alternative mechanism for the hydrogen ion

independent decomposition has been suggested by Fronaeus and Ostman(11) in

which only one sulfate free radical is formed in the initial step.

S2O8 H2O HSO4 SO4 OH

O22

- .+

1/2

2- -+ +

.

-.SO4+H2O -HSO4+

.OH

.OH H2O+

︵1 ︶︵2 ︶

︵3 ︶ On the other hand the mechanism of the acid catalyzed reaction, in dilute

solution is:-

HS2O8

SO3

︵4 ︶

H+2-S2O8+ -

-HS2O8-HSO4 + SO4

.-

SO4.-

+1/2O2

SO3 H2O+ H2SO4

︵3 ︶︵2 ︶

︵1 ︶

In strong acid solution is:

HS2O8H+2-S2O8+ -

-HS2O8-HSO4 + SO4

.-

SO4.-

H2O+ H2SO5 ︵3 ︶︵2 ︶

︵1 ︶

Little data have been published on this aspect of the decomposition and the

evidence for the sulfur tetroxide molecule as an intermediate rests on the

detailed study made by Kolthoff and Miller(6). The formation of sulfur tetroxide

explains the observation that the oxygen produced in acid solution comes from

the peroxodisulphate, but Bawn and Magerison(8) found that if SO4 was present

in the pH range 3–7, it did not show active radical characteristics, in that it did

not attack the free radical capture agent.

5

Also, Saxena and Singhal suggested that acid catalyzed reactions are caused by

the reactions

H+ .+ HSO4SO4

.-

and although they may take place, the oxygen production from the peroxodisul-

phate is not explained.

.HSO4 H2O+ HSO4

- OH.

+ +H+

Measurements have been made to determine the amounts of peroxymono-

sulfuric acid and hydrogen peroxide produced when more concentrated acid

solutions are used (greater than 2M). These indicate that the peroxodisulphate

first decomposes to give peroxymonosulfuric acid which hydrolyzes further to

hydrogen peroxide. The above result may be explained by the reactions:-

HS2O8

︵4 ︶

H+2-S2O8+ -

-HS2O8 H2O+ H2SO5 HSO4-+

H2SO5 H2O+ H2SO4 H2O2+

H2O2 +1/2O2H2O︵3 ︶

︵2 ︶︵1 ︶

Equation (1) is the rate determining step, producing the HS2O-

8 ion which,

because of the influence of the hydrogen ion, decomposes unsymmetrically by

reaction with water, and involves the breaking of an O–S bond.

SO

O

O

O O

O

O

OS

H

SO

O

O

OS

O

O

O

OH2O +OH HH

equation (1) involved the reaction between two oppositely charged ions

accounts for the negative salt effect. The peroxymonosulfuric acid formed by

equation (2) is further hydrolyzed to give hydrogen peroxide in equation (3),

which is thermally decomposed, producing oxygen from the peroxodisulphate

in equation (4). The difference in activation energy between the catalyzed and

6

uncatalyzed decompositions may be due, in part, to the difference in the energy

of O-O and O-S bonds.

Investigation on the silver ion catalyzed decomposition of peroxodisul-

phate(9,12,13) showed that the rate can be expressed in the form:-

[S2O8]2-2-

=/dt-d[S2O8] kobs

[Ag+]k2+= k1kobs

where k1 and k2 are rate constants for the non-catalyzed and silver ion catalyzed

decompositions respectively.

Table (1.1), values of k2 at 250C.

T0C k2 L.mol-1.min-1 µ m.l-1

25 1.290 0.015

25 0.210 0.300

25 0.225 1.288

The activation energy is found(8) to be 17.9 Kcal.mol-1, with a frequency factor

of 6.2x1011 L.mol-1.min-1

1.3 The first-order uncatalyzed peroxodisulphate oxidation

A major problem arising from the first order uncatalyzed oxidation is to

account for the increased rate of peroxodisulphate decomposition on addition of

a reducing agent, whereas the rate is independent of concentration of these

substrates. If the primary step is decomposition of peroxodisulphate into sulfate

free radicals, followed by a rapid attack of these in the reducing agent, then the

rate of oxidation should be the same for all reducing agents in this class.

SO4.-

S2O82-

2

Rate determining step These sulfate free radicals produced by the decomposition of peroxodisulphate

is evident from the fact that in polymerization studies using peroxodisulphate

7

labeled with sulfur35 as an initiator, polymer fragments containing radioactive

sulfur groups have been isolated(14,15). The breakdown of peroxodisulphate into

free radicals probably is irreversible, as it seems likely that sulfate free radicals

would exchange electrons with sulfate ions at a rate comparable with their

recombination to peroxodisulphate. However no exchange of sulfur35 between

sulfate and peroxodisulphate has been observed(16) under the condition of

kinetic experiments, and consequently any proposed mechanism of decomposi-

tion involving an equilibrium between peroxodisulphate and sulfate or sulfate-

ion radicals is probably incorrect. for example:-

2-SO4SO4S2O8

2- 2+2-

S2O8 SO4.-

and

Bartlett and Cotman(10) have suggested that radicals produced in the

peroxodisulphate decomposition in aqueous solution can not induce the

peroxodisulphate decomposition, and based their statement on the fact that

autoca-talysis is not observed in the thermal decomposition and the reaction is

first order in peroxodisulphate concentration. However, to explain the increased

rate on addition of an oxidizable substrate, it is necessary to postulate that

radicals produced from the reducing agent can induce peroxodisulphate

decomposition, and this suggests that a similar radical mechanism operates in

the absence of the reducing agent. The first order dependence can be explained

by a chain decomposition mechanism(10) , analogous to that suggested for the

uncatalyzed reaction between peroxodisulphate and oxalate as substrate:

Initial step:

22-S2O8 SO4

.-

OH.

+HSO4-H2O+SO4

.-

k1

k2

︵1 ︶︵2 ︶

Step (2) is slow, but faster than step (1).

︵3 ︶OH.

+2-

S2O8 SO4.-

+HSO4-k3 +1/2O2

Step (3) is fast, but governed by step (2).

8

OH.

+SO4.- k4 O21/2+-HSO4 ︵4 ︶

Step (4) is the chain termination. The primary step (1) is characteristic of all peroxodisulphate oxidations and

may be initiated by impurities present in the solution or light, as the aqueous

decomposition of peroxodisulphate is known to be photosensitive(17). In step (2)

the sulfate free radicals react with water to produce hydroxyl free radicals. This

in turn rapidly decomposes the oxidizing ions present by step 3 and 4.

Application of the steady state hypothesis to the radicals in the above scheme

leads to the rate law

-d[S2O82-]/dt = (k1k2k3/k4)1/2[S2O8

2-]

This mechanism scheme is readily extended to include cases when an oxidiz-

able substrate is added in a finite concentration and the observed rate law is

-d[S2O82-]/dt = -d[Xn-]/dt = kobs.[S2O8

2-]

Where [Xn-] is the concentration of the ion added, and kobs. is the observed

velocity constant. In such cases the rate constant is greater than that observed in

the aqueous decomposition. The proposed mechanism(18) is:-

2-SO4

+

︵2 ︶︵1 ︶

k2

k1

-.SO4+H2O -HSO4+

.OH

-.SO4S2O8

2-2

OH.

X2-

OH-X+ .-

2-S2O8

-.X+ + SO4

.-+ -X

-.X SO4

.-+ SO4

2-X-+

k3

k4

k5

︵5 ︶︵3 ︶

︵4 ︶

Here peroxodisulphate is decomposed by steps (1) and (4) leading to the rate

expression:-

-d[S2O82-]/dt = k1[S2O8

2-]+k4[ ] [S2O82-] = (k1+k4[ ] )[S2O8

2-]

where [ ] is the concentration of the free radical produced from the reducing

agent.

X.

X. X

.

9

By the steady state hypothesis the concentration of free radicals is constant.

Thus:-

kobs = k1 + k4 [ ]

Although the value of the constant free radical concentration may vary from

reactant to reactant giving rise to the variation in rate constant in the different

substrates.

1.4 Oxidation of amino-acids in vitro

The problem of oxidation of amino acids inside the body is one still waiting

satisfactory solution. The mechanism of the oxidation is not accessible to direct

experimentation. The close relationship among the amino acids, ketonic acids,

and hydroxy acids furnishes a basis for the interpretation of the mode of

utilization of amino acids by the body. Experiments on perfusing the liver of

various animals with different amino acids have demonstrated the presence of

ketonic acids in the perphusate(19).

The deamination of the amino acids in the liver and the production of urea are

also well established; amino acids may also be decarboxylated in the organism,

as is shown by the presence in the body fluids of Histamine and Tyramine (19).

The oxidation of the amino acids in-vitro has been the subject of many

investigations, both from the purely chemical standpoint and from the point of

view of its bearing on the mechanism of the amino acid metabolism in the

organisms.

Amino acids are resistant to hydrolyzing agents; they do not react readily

with reducing agents, but are very reactive with oxidizing agents in general(20).

The production of aldehydes from amino acids has been demonstrated by many

authors with a diversity of oxidizing agents(21).

Langheld was the first who studied the effect of sodium hypochlorite on the

amino acids and found that carbon dioxide, ammonia, and “C-armere aldehyde”

were produced in the reaction.

Dakin(22) who worked on the reaction of amino acids and sodium hypochlorite

followed Langheld. Dakin used sodium N–Chloro-p-toluenesulfonamide

X.

10

(chloramine–T) and found that although this reagent contain no free sodium

hypochlorite, it acted in the same manner as would a solution of sodium

hypochlorite. Chlorination of the amino group to give N-chloro amino acid

being the first product which decomposes yielding ammonia, carbon dioxide

and corresponding aldehyde; Similarly Made Gowda(23, 24) and his co-workers

studied the oxidation of amino acids by sodium N–chloro-p-toluenesulfon-

amide (chloramine–T) in acid media and in alkaline media. They investigated

the kinetics of oxidation of α-amino acids, glycine, valine, alanine, and

phenylalanine in HClO4 medium at 300C. The rate showed first-order

dependence on both chloramine–T and amino acid concentrations and an

inverse first-order on [H+], in alkaline media they investigated the kinetics of

oxidation of argnine, histidine, and threonine by chloramine-T at 350C. The rate

is first order in both chloramine-T and amino acid concentrations, and inverse

fraction order in OH- concentration for argnine and histidine oxidation and the

rate is independent of OH- concentration for threonine oxidation. The

thermodynamic parameters studied by Gowda et al., for the system amino acid-

chloramine-T in acid and alkaline media can be summarized bellow.

Table (1.2) Values of the thermodynamic parameters of oxidation of some

amino acids by chlorammine-T in acidic and alkaline media.

Alkaline media (24) Acidic media (23)

Thr. His. Arg. Phe. Ala. Val. Gly. Parameter

119.7 95.8 124.7 65.7 72.2 73.6 95.8 Ea (KJ.Mol-1)

117.1 93.2 121.9 63.2 69.7 71.0 93.3 ∆H#(KJ.Mol-1)

76.6 3.8 85.5 -99.0 -77.0 -73.3 -6.8 ∆S#(J.Mol-1.K-1)

All experiments were carried out at strict conditions provided in each reference.

The values of activation energies of oxidation of amino acid by chloramine-T

were relatively high in alkaline media as well as with values in acidic media.

Generally chloramine-T (CAT) undergoes a two-electron change in its

reactions resulting in the formation of the reduction products, p-toluenesulf-

11

onamide (PTS) (p-CH3C6H4SO2NH2) and sodium chloride. The oxidation

potential of the CAT–PTS couple varies with pH of the medium (its 1.139V at

pH 6.5, 0.778V at pH 7.0, and 0.614V at pH 9.7) and the following mechanism

was proposed by Gowda, for the oxidation of amino acid by chloramines-T:

H3C S N

O

O

H

HCl k1

fastCl H

H

O

O

NSH3C ++

+RNH2Cl or CAT RNHCl

........ ︵i ︶

The fast pre-equilibrium step (i) involving deprotonation of RNH2Cl+ forms the

reactive oxidant species, RNHCl. In the slow step, an electrophilic attack by Cl+

of RNHCl on the carboxylate anion of the substrate, resulting in the formation

of an N-Chloro-O bridged transition state, X, is invisaged.

RNHCl

CH3SNH

Cl O

O

CO

OC HN

HHH R

+

RHHH

N HCO

OC

O

O

ClHN S CH3k2

slow

...... .....δ δ

︵S ︶ ︵X ︶

.... ︵ii ︶

CH3SNH

H

O

OOCCR

NHH

O Cl

+fast ....... ︵iii ︶

︵X ︶ ︵RNH2 ︶

-

After the transient state X is formed, the chlorine atom probably undergoes a

Walden-type of inversion as the carboxylate anion attacks and kicks out the

sulfonamide anion species, which is resonance stabilized and which later

abstracts the acidic proton off the amino acid nitrogen and forms the

sulfonamide product as in fast step (iii).

..

H

HN H

O

OCC

R'

Cl

︵X' ︶

︵fast ︶ R' C NH2 CO2 ClH

+ + -

︵iv ︶

12

HNH2C

R'

OH H

︵fast ︶ RC

H

NH3

O H

︵fas t ︶ RC

HO NH4+ ︵v ︶

:

In the subsequent fast steps (iv) and (v), the complex X' undergoes intramo-

lecular rearrangements and a nucleophilic attack by water to form the end

products including the aldehyde. Gowda found that by application of the

steady-state concept to the intermediate, RNHCl, in Scheme above leads to the

following rate law:

Rate = -d[CAT]/dt = k1k2[S][CAT]/[H+]

Gowda(25) also used Sodium N-Chlorobenzene Sulfonamide (chloramine-B) to

oxidize alanine and phenylalanine in hydrochloric acid medium at 300C in two

ranges of acid concentrations, and found that the reactions follow identical

kinetics for both amino acids. At low acid concentration (0.03-0.1M),

simultaneous catalysis by H+ and Cl- ions is noticed. The rate shows a first-

order dependence on CAB concentration, but is independent of substrate

concentration. At higher acid concentration > 0.2 M, the rate is independent of

H+ concentration, but shows a first-order dependence in CAB concentration and

a fractional-order dependence on substrate concentration. The thermodynamic

parameters studied by Gowda et al for the system amino acid-chloramine-B in

low acid and high acid concentrations are summarized bellow.

Table (1.3), values of thermodynamic parameters at low and high acid concent-

rations for the oxidation of some amino acids by chloramines-B.

High acid (>0.2 M)(25) Low acid (0.030.1M)(25)

Phe. Ala. Phe. Ala. Parameter

119.7 88.9 67.0 68.0 Ea (KJ.Mol-1)

117.1 86.4 64.4 65.4 ∆H# (KJ.Mol-1)

115.1 8.3 -60.0 -68.7 ∆S# (J.Mol-1.K-1)

81.5 83.7 83.2 86.8 ∆G# (KJ.Mol-1)

All experiments were carried out at strict conditions provided in each reference.

13

Schmidt et al.(26) discussed the degradation of the amino acids, poly–peptides

and diketopiperazines with sodium hypobromite and showed that this was

similar to the reaction with sodium hypochlorite, aldehydes and nitriles being

formed. The tendency to react with hypobromite at higher temperatures is

especially obvious with substances which do not simply form nitriles or

aldehydes, but with substances, such as uric acid, in which the oxidation

products are of more complex nature. The following scheme for the oxidation

of amino acids by sodium hypobromite embodies Langheld’s, Dakin’s and

Schemidt’s results;

R C C OH NaOBr R C CHNH

OHOH

O

HNBr

H

NeutralNaOBr

Alkaline

︵M+

︶︵OH-

︶

BrNBr

OR C C OH

H HN OR C C OH

H2O

R CO

H CO2 NH3

NaOBr

N2 H2O

1.5

1/2 3/2

+

++R C N HBr CO2++

+ The first step of the reaction is bromination of the amino group. In alkaline

solution there is greater tendency for this group to split off hydrogen bromide

with the probable formation of an imine, which further hydrolyzed to the

aldehyde, ammonia and carbon dioxide. The formed ammonia reacts further

with 1.5 mol NaOBr of the reagent. In less alkaline or neutral solution there is

greater tendency for the formation of the dibromo substitution product and

subsequent nitrile formation.

Lal(27), Reddy(28), and Ramachandran(29) studied the oxidation of amino

acids by N-bromoacetamide; Lal(27), studied the oxidation of amino acids by N-

bromoacetamide in aqueous perchloric acid solution, and found that the main

14

products of the oxidation were the corresponding carbonyl compounds, and the

reaction is of first-order with respect to the oxidant and the amino acid, the

oxidation of deuteriated glycine indicated the absence of a primary kinetic

isotope effect. Reddy(28) studied the oxidation of amino acids in acid and

alkaline media, and noticed that the order with respect to amino acid was

dependent on the nature of the medium, its 0.3-0.8 order in perchloric acid and

alkaline media, and zero order in aqueous acetic acid; Irrespective of the

medium, the reaction was first-order in oxidant concentration, and the rate of

oxidation increase with OH- and decrease with H+ concentrations.

Ramachandran(29) studied the oxidation of amino acids by N-Bromoacetamide

in aqueous buffered medium at 35°C, the rate of decreasing oxidant

concentration was catalyzed by the Br- produced from the reduction of N-

Bromoacetamide; Analysis of the autocatalyzed reaction gave the kinetic data

for the oxidation of bromide ion by N-Bromoacetamide.

The results suggested that the protonate N-Bromoacetamide reacted with Br- to

form Br2, which rapidly oxidized amino acids.

Amino acids were also oxidized by hydrogen peroxide. It has been shown that

the action of hydrogen peroxide on casien dissolved in formic acid causes a

selective oxidation of its tryptophan, methionine and (partly) cysteine units. As

regards the effect of other acid oxidizing agents on the natural amino acids,

Williams and Woods(30) stated that among 16 samples tested only cysteine,

tyrosine and tryptophan were oxidized by iodic acid [at 1000C]; Nicolet and

Shinn(31) showed that periodic acid selectively attacks tryptophan, methionine

and cysteine, as well as the α–hydroxy amino acids. Since, however, according

to the same authors, the latter type of compound seemed to be protected against

the oxidation by acylation or by peptide formation through the amino acid

group. The action of performic acid (the product of interaction of hydrogen

peroxide and formic acid) which showed α–hydroxy amino acid units to be

resistant, had been obtained on unhydrolyzed protein, it became of interest to

examine the action of performic acid of the free amino acids.

15

Radhy(32), Joaquin(33), Insausti(34), and Arrizabalaga(35), studied the oxidation

of amino acids by permanganate; Radhy(32) studied the kinetics of homogeneous

acid-catalyzed oxidation of certain amino acids by potassium permanganate in

moderately concentrated acidic media. He found that the rate of oxidation of

glycine, DL-alanine, DL-valine, and DL-leucine by potassium permanganate in

aqueous sulphuric and perchloric acid solution was proportional to the

concentration of amino acid. For each amino acid the total order of the reaction

was two at a given concentration of sulphuric and perchloric acid. The rate of

oxidation of amino acid was greater in sulphuric acid than in perchloric acid for

the same concentration. There is no primary salt effect, but at a higher

concentration of added neutral salts, the logarithm of the rate constant is

linearly related to ionic strength.

Joaquin(33) studied the permanganate oxidation of L-valine in neutral aqueous

solution by visible spectrophotometry. He found that under these conditions,

both the zwitterionic and anionic forms of the amino acid are oxidized; the

reaction being autocatalyzed by soluble colloidal manganese dioxide.

Insausti(34) studied the oxidation of glycine by permanganate in phosphate

buffered solutions, and found that the product obtained was identified as a

soluble form of colloidal manganese dioxide. The influence of glycine on the

colloidal product showed that changes in glycine concentration alters the

extinction coefficients of the colloidal product, and only when the influence of

glycine has been eliminated, the theoretical and experimental reaction rates

coincide. Arrizabalaga(35) studied the kinetics of the oxidation of L-α-amino-n-

butyric acid by permanganate ions in buffered acid medium at pH=1-3 using

spectrophotometric technique, and found that an autocatalytic effect had been

observed in all cases due to Mn2+ ions formed as a product of the reaction.

A first-order reaction with respect to the amino acid and the permanganate ions

in both processes, catalyzed and uncatalyzed, was obtained; and studied the

influence of pH, temperature, ionic strength, and reactants concentration on the

reaction rate.

16

Table (1.4) gives summary of the thermodynamic parameters determined for

the oxidation of certain amino acids by permanganate ion.

Amino acid ∆S# (J.Mol-1.K-1)

∆H# (KJ.Mol-1)

Ea (KJ.Mol-1) Ref.

Glycine (H2SO4 Media)

Glycine (HClO4 media)

-128.74

-116.25

-

-

45.73

50.54 (32)

DL-Alanine (H2SO4 Media)

DL-Alanine (HClO4 media)

-28.63

-53.13

-

-

79.09

72.44 (32)

DL-Valine (H2SO4 Media) -105.29 - 53.38 (32)

DL-leucine (H2SO4 Media)

DL-leucine (HClO4 media)

-25.96

+22.61

-

-

72.44

95.30 (32)

L-α-Amino-n-butyric acid

Catalyzed process.

Uncatalyzed process.

-18.60

-210.00

66.3

37.4

69.1

40.1

(35)

All experiments were carried out at strict conditions provided in each reference.

Joaquin(33) observed that the colloidal manganese dioxide had been detected as

a product of the reduction of permanganate ion by amino acid, both as a

brownish yellow soluble colloid and as a brown precipitate, one day after

completion of the reactions. Furthermore, carbon dioxide, aldehydes, and

ammonium ion were considered as the products from the oxidation of α -amino

acids by permanganate ion. All these considerations, along with the kinetic

results led to propose the following mechanism for the uncatalyzed pathway,

involving the oxidation of both the zwitterionic and anionic forms of amino

acids

17

RCH ︵NH3+

︶CO2- MnO4

- RCH ︵NH3+

︶CO2 MnO42-.

.RCH ︵NH3

+

︶CO2

+ +

+

︵1 ︶

RCHNH3+

+

CO2 ︵2. Manganate

.RCHNH3

+ RCHNH2 H+.

RCH ︵NH2 ︶CO2- MnO4

-+

slow

.RCH ︵NH2 ︶CO2 + MnO4

2-s low

.CO2RCHNH2 +RCH ︵NH2 ︶CO2

.

RCHNH2.

MnO4-+ RCH NH2

+ MnO42-+

RCH NH2+ + H2O RCHO NH4

+

MnO42- H2O MnO4

- MnO2 OH-

+

3 + 2 2 + + 4

︵3

︵4

︵5

︵6

︵7

︵8 ︶ Transfer of one electron from the reducing agent to the permanganate ion,

results in the formation of manganate ion. The latter is known to be very

unstable in all but in strongly alkaline media manganese dioxide appears via

manganate ion dismutaion(36). Aldehydes and ammonium probably being

formed by hydrolysis of iminium ion (step 7).

Gowda(37, 38) , and Kamaluddin(39) studied the oxidation of amino acids by

manganese (III) ions; Gowda(37) studied the kinetics of oxidation of L-serine by

manganese (III) ions in aqueous sulphuric acid media at 500C. The dependence

of the reaction rate is one and a half-order on manganese (III) concentration,

first-order on serine concentration, and inverse first-order in H+ concentration,

and an inverse fractional order on manganese (II) concentration. Gowda(38)

studied the kinetics of oxidation of L-aspartic acid and L-glutamic acid by

manganese (III) ions in aqueous sulphuric acid, acetic acid, and pyrophosphate

media. The reaction showed a variable order in initial concentration of

manganese (III) ions. The order changed from two to one as the reactive

oxidizing species changed from an aqua ionic form to a complex form.

18

There is a first-order dependence of the rate on initial amino acid concentration

in all the three media, while the other common features include an inverse

dependence on H+ and on manganese (II) concentrations.

Kamaluddin(39) studied the kinetics of oxidation of DL-α-aminobutyric acid,

DL-isovalin, DL-n-valine, and L-leucine by manganese (III) ions in sulphuric

acid medium, all the amino acid were found to follow similar kinetics, the nature

of the reaction was very much dependent on the initial manganese (II)

concentration present in the reaction mixture. The reaction showed a first or

second order dependence on manganese (III) concentration depending on

whether the initial manganese (II) concentration in the reaction mixture is less

than 0.01M or greater than 0.15M. In both cases, the reaction showed a first

order dependence on the amino acid concentration; and an inverse first order

dependence upon H+ concentration, and an inverse first order dependence on

manganese (II) concentration was observed in the presence of high

concentration of manganese (II). Gowda(37) assumed that Mn(III) species present

in aqueous sulphuric acid medium were Mn3+(aq), Mn(OH)2+(aq), and

MnSO4+(aq), due to the existence of the following equilibrium:

MnOH2+

︵aq ︶ HSO4-

︵aq ︶ + MnSO4+

︵aq ︶ H2O+

Since the effect of HSO4- on the reaction is negligible, Mn3+

(aq) and

Mn(OH)2+(aq) species become important. In view of this, the proposed

mechanism is shown bellow:

MnOH R CHNH3

COOHk1 R CH C OH

NH2 O

Mn

H2O+ +2+aq

3+

....... ︵i ︶

+

3+

Mn

ONH2

OHCCHR k2

k -2

R CH C OHNH2 O

Mn

.++

2+......... ︵ii ︶

19

R CHNH2

CO

OH k3 R CH+

NH2

Mn H+ CO2

.+

+ Mn3+ .+++ 2+ ........ ︵iii ︶

++.

NH2

CH+R H2O R CHO NH4+ ....fast ︵iv ︶

By Applying of the steady state approximations to the above mechanism the

rate law can be derived as:

Rate = k1k2k3[Mn(III)]2[Amino acid]/[H+]{k-2[Mn(II)]+k3[Mn(III)]}

The rate law is consistent with the observed kinetic data, first-order in [amino

acid], one and a half-order in [Mn(III)], a negative first-order in [H+], and a

negative fractional-order in [Mn(II)]. Nalwaya(40) studied the oxidation of

α-amino acids by pyridinium bromochromate (PBC) in acetic acid–water

mixture containing perchloric acid. The reaction rate is first order in [PBC] and

inverse first order in [H+] and aldehyde being a product. The reaction rate

increases with a decrease in the polarity of solvent indicating an ion–dipole

interaction in the slow step. The reactions exhibit no primary kinetic isotope

effect. Nalwaya(40) studied the thermodynamic parameters of the above reaction

as shown in table (1.7) bellow.

Table (1.5), thermodynamic parameter values of the reaction amino acid

pyridinium bromochromate.

α-amino acid ∆S# (J.Mol-1.K-1) ∆G# (KJ.Mol-1) Ea# (KJ.Mol-1)

Glycine -106.1±4.2 95.3±2.9 63.2±2.8

Alanine -65.1±3.8 96.3±2.9 76.6±3.1

Valine -4.4±2.9 99.2±2.8 97.9±3.8

iso-Leucine -55.7±4.3 98.9±2.6 81.5±3.9

Nor-Leucine -115.3±4.0 101.3±3.0 66.4±2.9

Phenylalanine -97.5±2.6 98.5±3.2 68.9±2.9

The value of the activation energy is approximately 63 to 98 KJ.mol-1. The

entropy values are all negative and high (except valine) suggesting that the

20

transition state is more rigid and extensively solvated than the reactants. The

negative entropy also suggests the formation of cyclic intermediate from acyclic

species. The plot of ∆S# against Ea# is nearly linear suggesting that all the

amino acids studied follow a similar mechanism.

Amino acid are also oxidized by Bismuth (V)(41), Cerium (IV)(42), some free

stable hydrazyl radicals(43), alkaline hexacyanoferrate (III)(44), peroxomonop-

hosphoric acid(45), and peroxomonosulphate(46).

Ramachandran(46) , studied the oxidation of amino acids by peroxomono-

sulphate in the presence and absence of formaldehyde, analysis of the results

shows that the rate at constant H+ concentration and in the absence of

formaldehyde are first order in both amino acid and peroxomonosulphate

concentrations, and in constant H+ concentration and presence of formaldehyde

can be present as:

-d[PMS]/dt = kb[AA]HCHO][PMS] + kc[HCHO][PMS]

Perusal of the kinetic results shows that the formaldehyde catalyzed reaction

occurs ~ 105 times faster than uncatalyzed reaction and this is attributed to the

formation of Schif's base.

1.5 Oxidative decarboxylation of amino-acids

Spencer(47) suggested the following mechanism entailing electrophylic

attack by the oxidizing agent, formally represented as OH+, on the amino acid at

the α-carbon atom. This becomes transiently anionoid as a result of

simultaneous loss of the carboxyl group as carbon dioxide, to give a

carbinolamine (aldahyde-ammonia) as the primary product. This rapidly

decomposes to aldehyde and ammonia.

R CHNH3

CO

H+CO2

++

NH2

R CH+

OH O- --

OHCHRO

NH3-

The First stage of this mechanism is analogous to the heterolytic bimolecular

electrophilic substitution (SE2) mechanism of thermal decarboxylation of some

organic acids, discovered by Schenkel and Schenkel–Rudin(48).

21

Direct isolation of the postulated intermediate did not appear promising since

most aldehyde–ammonia compounds are unstable.

To test the proposed mechanism, α-N-diphenyl glycine was oxidized at room

temperature with one equivalent of potassium permanganate or potassium

ferricyanide in the presence of two equivalents of sodium hydroxide. In each

case the expected Schiff’s base, benzalalinine, identical with an authentic

specimen, was isolated in good yield. Previous attempts to oxidize the amino-

acid had failed to give recognizable products.

Ph CH COO-

+NH2 Ph

OH+

H+ CO2Ph CH OH

NH Ph

H2O Ph CH

N Ph- -,+ -

︵I ︶ ︵I I ︶ ︵I I I ︶ The end product could not have arisen by recombination of aniline and

benzaldehyde, if these had been the primary products of the oxidation, for cold

alkaline permanganate is instantly reduced by aniline, benzaldehyde, and

α-N-diphenylglycine, whereas benzaniline reacts only slowly; potassium

ferricyanide reacts quickly with aniline and amino acid, but only very slowly

with benzaldehyde and with anil. Isolation of the anil thus constitutes direct

evidence in favor of the proposed mechanism, since it could have been formed

only by dehydration of the carbinolamine (II).

It is not worthy that a mechanism similar to the above was postulated by

Robinson et al.(49) in the special case of the oxidative conversion of tryptophan

to Harman; Spencer(47) and his co-workers indicated that the rate of oxidative

decarboxylation of α-amino acids increased in general with decreasing pK1, i.e.

with increasing power of electron attraction of the substituent R group.

Correlation of the rate of reaction and pK2 is more difficult, but since N-mono

and N-di-alkylated amino acids are oxidized more readily than the parent

compound, and the N-alkylation increase the basic strength of an amine, the

reaction appears to be favored by an increase in pK2.

The fact that trialkylbetaines resist oxidation(50), suggests that the availability of

an ionizable proton on the amino group is a prerequisite for the reaction.

22

2. Experimental 2.1 Preparation of solutions

The reaction of peroxodisulphate with organic compounds is very sensitive to

impurities(51), King and Steinback(52) showed that the variations observed in

their results were due to chloride ion in the distilled water they used. Therefore

the water used in this investigations is deionized water, used to prepare different

solutions, as medium for the kinetic experiments and for washing the reaction

vessels.

(1) Sodium thiosulphate solution

The required concentration (0.01M) of sodium thiosulphate pentahydrate (A.R.)

was prepared by dissolving 2.4817g of sodium thiosulphate in 1.0L of

deionized water.

(2) Peroxodisulphate solution

Potassium peroxodisulphate (A.R.) was recrystalyzed from deionized water and

was dried in a vacuum desiccator. It was standardized using a method modified

by Rosin(53), 1.0g of dried recrystalized peroxodisulphate was added to a

solution of 25ml of 1.0M potassium iodide and followed by immediate addition

of 0.5g of sodium bicarbonate, 10ml of 10% sulfuric acid was added. During

and after the liberation of carbon dioxide, which provide a convenient means of

sweeping out the oxygen, the solution was left to stand in a glass-stoppered

flask for thirty minutes in the dark at room temperature. The liberated iodine

was titrated against standard thiosulphate solution.

(3) Potassium iodide solution

200g of potassium iodide (A.R) was dissolved in 1.0L of deionized water to

give 20% (w/v) solution.

(4) Amino acids solutions

The amount of 1.0g of chromatographically pure amino acid (Philip Harris

Biological) was dissolved in 100ml for L-Alanine, 100ml for L-Arginine,

100ml for L-Asparagine, 250ml for DL-Aspartic acid, 500ml for L-Cysteine,

250ml for L(+)-Glutamic acid, 100ml for L(+)-Glutamine, 100ml for Glycine,

23

100ml for L-Histidine.HCl, 100ml for L-Leucine, 250ml for Lysine, 100ml for

DL-Methionine, 100ml for DL-Phenylalanine, 100ml for Proline, 100ml for L-

Serine, 100ml for L-Threonine, 500ml for L-Tyrosine, 100ml for L-Valine], in

deionized water to give standard solutions. Freshly prepared solutions were

used every time.

(5) Sodium bicarbonate solution

4% (w/v) stock solution of sodium bicarbonate was prepared by dissolving 40g

of sodium bicarbonate (A.R) in 1.0L of deionized water.

(6) Sulfuric acid solution

98% sulfuric acid (A.R) was diluted to give 0.5M sulfuric acid using deionized

water. This was kept stoppered as a stock solution.

(7) Starch indicator

1.0g of soluble starch (A.R) was mixed with 100ml deionized water and 2.0g of

potassium iodide was added, a freshly prepared solution was used every time.

(8) 2,4 –Dinitrophenylhydrazine Reagent

* The first reagent is prepared by means of Brady’s method for aldehydes

produced from aromatic amino acids, 2.0g of 2,4-Dinitrophenylhydrazine was

treated with 4ml concentrated sulfuric acids and 30ml methyl alcohol was

added cautiously with cooling, 10ml of deionized water was added.

* The second reagent is saturated solution of 2,4-Dinitrophenylhydrazine in

2.0M HCl was prepared for aldehydes produced from aliphatic amino acids.

(9) Nessler's reagent

100g of mercury (II) iodide and 70 g of potassium iodide were dissolved in

100ml deionized water. The resulting solution was added with stirring to a

solution of 160g of sodium hydroxide in 700ml deionized water. This was then

diluted to one liter with deionized water. The precipitate was allowed to settle

for three days and the supernatant liquid was decanted and kept in a brown

bottle.

24

(10) Chromotropic acid solution

Chromotropic acid, disodium salt [4,5-Dihydroxynaphthalene-2,7-disulfonic

acid, disodium salt dihydrate [(HO)2C10H4(SO3Na)2.2H2O] was prepared by

dissolving 0.5g of sodium di-salt of the acid in 50ml of deionized water in

250ml volumetric flask, then made up to the mark with 70% (v/v) sulfuric acid.

2.2 Apparatus

* The reaction vessels (250ml Pyrex Round bottom flasks containing the

reaction mixtures) were kept at a constant temperature in a thermostatic water

bath. Thomas and Ephraim(54) showed that the results of oxidation reaction of

peroxodisulphate were not significantly different when the experiments were

conducted in the dark or light.

* The absorption spectra were measured at room temperature with Parkin-

Elmer UV/VIS spectrophotometer Model 550S.

* The aliquot portions of reaction mixture were withdrawn by Eppendorf 10ml

automatic pipette.

* The amounts of sodium bicarbonate and sulphuric acid were taken by 10ml

dispenser.

* Volumes of organic reagents were taken by 10ml safety pipette.

* I.R measurements of the hydrazones were measured by Fourier Transform

Infrared Spectrophotometer (FTIR) 8400S (Shimadzu).

* Titration of the residual peroxodisulphate with the thiosulphate solution was

carried out by 50ml titration apparatus (Aldrich acc to Dr. Schilling).

* pH measurements were carried out by microprocessor pH meter 211.

* Concentration of some cations was determined by Atomic Absorption

Spectro-photometer AA-6800.

2.3 Washing procedure

The reaction rate and the reproducibility of the results depend to some extent

on the method used for cleaning the reaction vessels. Therefore, special care

was taken for cleaning the reaction vessels:-

(1) Washing with soap (detergent) and tap water 5 times.

25

(2) Washing with deiniozed water (3 times) with vigorously shaking.

(3) The reaction vessels left to dry either in an oven or left overnight.

2.4 Identification methods for the reaction products

Product analysis was carried out under kinetic conditions. A mixture of slight

excess of peroxodisulphate over amino acid were made up to 50ml with

deionized water and kept in a water bath thermostat at 60 0C for approximately

five hours and the following tests were carried out:-

(1) Test for Ammonia

5.0ml of the reaction solution was diluted to 50ml and a few drops of Nessler's

reagent were added; a brownish color indicated the presence of ammonia.

(2) Test for Carbon dioxide

A small volume of reaction solution was re-evaporated and the gaseous product

formed was passed through freshly prepared lime water solution which turned

milky confirming the presence of carbon dioxide.

(3) Test for Aldehydes

The reaction mixture was treated with solution of 2,4-Dinitrophenylhydrazine

in methanol for products of aromatic amino acids and in solution of

2,4-Dinitrophenylhydrazine in 2.0M HCl for products of aliphatic amino acids.

The hydrazones derivatives were recrystalyzed from ethanol. Furthermore the

recorded infrared spectra of these hydrazones and authentic hydrazones were

obtained.

2.5 Estimation of the residual peroxodisulphate

Aliquot portions (10ml) of the reaction mixture were withdrawn at suitable time

intervals by means of 10ml automatic pipette and transferred into glass

stoppered conical flasks; The following reagents were added (in the same order)

(5ml of 4% sodium bicarbonate, 1ml of 0.5M Sulfuric acid, 5ml of 20%

potassium iodide), the liberated iodine after 15 minutes was titrated against

standard sodium thiosulphate solution using starch as indicator. The volume of

thiosulphate obtained at any time interval was taken to represent the residual

concentration of peroxodisulphate.

26

It should be noted that quenching of the amino acid-peroxodisulphate reaction

was accomplished by the much faster reaction of peroxodisulphate with iodide

ion. King and his co-workers(55) have shown that, under the conditions

employed, the peroxodisulphate iodide reaction was at least one hundred times

faster than peroxodisulphate-oxalate reaction.

2.6 Kinetic measurements

The reduction of peroxodisulphate by organic compounds was very slow

at ordinary temperature. It was found that at 600C the reaction proceeds at a

measurable rate. Thus the temperature range (60-800C) was chosen for the

kinetic runs. To establish appropriate rate laws for the reactions, sets of runs to

be carried out in which the concentration of one of the reactants was changed,

while the concentration of the other reactant was kept constant.

Kinetic runs were carried out by keeping the solution of calculated amounts of

amino acid, 0.25 M potassium sulfate (to keep the total ionic strength constant

at 0.25), and calculated volume of deionozed water to complete the total

volume to 100ml after addition of peroxodisulphate, in the thermostatic water

bath at 600C for 20 minutes before starting the experiment by addition of the

required volume of peroxodisulphate solution, peroxodisulphate solution was

usually not preheated to thermostat temperature before addition, therefore the

volumes of peroxodisulphate solution were taken as small amounts as possible

to decrease the effect of lowering in the reaction temperature, the time of

mixing the reactants was taken as zero, at suitable intervals of time portions of

reaction mixture (10ml) were withdrawn and the amount of unreacted

peroxodisulphate was estimated iodometrically, the titre values at any time were

taken to represent the residual peroxodisulphate concentrations at that time and

will be referred to as [S2O82-].

27

The observed rate constants were calculated by the integrated rate law for a first

order reaction, as follows:-

kobs = 2.303 x Log [S2O82-]0 (2.1)

t [S2O82-]

The study required a knowledge of the rate of the reaction (R) which was

calculated as follows; Suppose V0 (ml) was the initial titre value at the zero

time for a reaction mixture in which peroxodisulphate concentration was M0 (in

mol.L-1), if after a time interval (per minutes) the titre value for an other sample

of reaction mixture was V (ml), then the concentration of peroxodisulphate in

the reaction mixture after this time interval will be equal to: -

V/ V0 = M/ M0

Then:

M = V. M0/V0 (2.2)

therefore, the change in peroxodisulphate concentration will be,

M0 ـ M = ∆M mol.L-1

and the rate of the reaction (R) will be,

R = ∆M/∆t mol.L-1.min-1. (2.3)

Thus, for each concentration of peroxodisulphate it was possible to calculate the

value of (kobs) and (R). The kinetic measurements were carried out over the

temperature range (60 800 ـC), the plots of Log kobs against 1/T were linear and

the activation energy (Ea) was calculated from the slope of the plot from

Arrhenius equation:-

kobs = Ae-E/ŘT (2.4)

or:

Ln kobs = Ln Aـ(E/ŘT) (2.5)

where A, is the frequency factor; Ea, is the activation energy; Ř, is the gas

constant and is equal to 1.987Cal.deg-1.mol-1 or 8.314 J.deg-1.mol-1, and the

slope of the graph of Ln kobs against 1/T is

Slope = -E/Ř

E = - slopexŘ (2.6)

28

The values of frequency factor (A) obtained from the intercept of the straight

line with Y-axis. The change in entropy (∆S#) and the change in the free energy

(∆G#) were calculated using the following equations(56):

k = (RT/Nh) e-∆G/RT (2.7)

and k = (RT/Nh) e∆S/R .e-∆H/RT (2.8)

Where R/N is Boltzman’s gas constant which is 1.3806 x10-23 J.K-1 and h is the

Planck’s constant 11.03 x10-36 J.Min;

The factor e-∆H/RT can be replaced by the factor e-∆E/RT using the experimental

activation energy E, equation (2.8) can be written as

k = (RT/Nh) e∆S/R.e-∆E/RT (2.9)

Comparing equation (2.9) with Arrhenius equation (2.4) the value of A will be

A = (RT/Nh) e∆S/R (2.10)

Taking the logarithm of the equation (2.10)

Ln A = Ln(RT/Nh) + ∆S#/R

And thus

∆S# = R [Ln A ـ Ln (RT/Nh)] J.K-1 (2.11)

Substituting the values of R/N, Boltzman’s gas constant, and h Planck’s

constant in equation (2.11)

∆S# = R [Ln A ـ Ln 1.2517x1012T] J.K-1 (2.12)

And value of ∆G is obtained from the following equation, equation (2.10)

written as

∆G# = ∆H# - T∆S# (2.13)

29

3. Results 3.1 Preliminary Studies 3.1.1 Stoichiometrey Stoichiometrey of the peroxodisulphate-amino acids reactions were extensively

studied(57, 58, 59, 60, 61) by several methods who found that the reaction followed

a 1:1 stoichiometry.

This is a modification method in which Chromotropic acid and ninhydrine

were used and spectrophotometric techniques were applied to evaluate the

concentrations of remaining amino acid and aldehyde which proportional to

amount of amino acid reacted, this modification here is applied for simple

amino acid (Glycine) as example.

For this modification, excess of peroxodisulphate concentration in the

reaction mixture was avoid; In five cleaned, dried glass-stoppered conical flasks

10, 20, 30, 40, and 50ml of 0.13M glycine solution, 25ml of peroxodisulphate

solution (0.02M) were added. The volume in each flask was completed to

100ml with deionized water. These were kept in a thermostat at 80ºC for three

hours, and then cooled to room temperature. Assuming that the reaction

proceeds as 1:1 mole ratio between glycine and peroxodisulphate,

formaldehyde was formed in each reaction mixture with equal concentration

because the same quantity of peroxodisulphate was used, and the same

absorbance value was expected.

2.5ml of each reaction mixture were pipetted into 250ml conical flasks, and

10ml of chromotropic acid were added, the volume was completed to 50ml

using deionized water and allowed to stand in a water bath at 80ºC for 30

minutes; the solutions were cooled to room temperature; the absorbance of the

purple solutions were determined at λmax 545nm, using Perkin-Elmer UV/Vis

spectrophotometer model 550S.

30

Table (3.1) The absorbance values of aldehyde-chromotropic acid complex

produced from the reaction mixtures.

Flask number 1 2 3 4 5

Absorbance 0.437 0.428 0.425 0.428 0.435

From the above absorbance values, formaldehyde gave approximately the same

concentration in each reaction mixture. This is in agreement with the

assumption that peroxodisulphate reacts with glycine in mole ratio 1:1.

To determine the remaining glycine, calibration curve of ninhydrine-glycine

complex was constructed. 5ml of 2% ninhydrine solution in deionized water

was added to 2, 4, 6, 8, and 10ml volume of 0.13M glycine solution in five

250ml conical flasks, the volumes were completed to 100ml using deionized

water, then placed on a water bath at 100 ºC for 15 minutes, the flasks were

cooled to room temperature and the absorbence were measured using Perkin-

Elmer UV/Vis spectrophotometer model 550S.

Table 3.2 Absorbance of standard solutions of glycine-ninhydrine complex at

λmax 540 nm.

Flask

number

Glycine concentration

x103 ( M) Absorbance

1 0.513 0.287

2 1.030 0.469

3 1.540 0.720

4 2.050 1.090

5 2.570 1.228

The absorbance values were plotted against glycine concentration as is shown

in fig (3.1). To determine the remaining glycine, 10ml from reaction mixtures 1,

3, 5 were pipetted into 250ml conical flasks, 5ml of ninhydrine were added, and

the volumes were madeup to 100ml using deionized water, then placed on a

31

water bath at 100ºC for 15 minutes and then cooled to room temperature. The

absorbance was determined, and results were tabulated below.

Table (3.3) absorbance of remaining glycine in the reaction mixtures.

Concentration of the remaining glycine Flask

number Absorbance

theoretically (M) graphically (M)

1 0.033 0.06 x 10-3 0.07 x 10-3

3 0.159 0.325 x 10-3 0.33 x 10-3

5 0.287 0.59 x 10-3 0.6 x 10-3

The theoretical concentration of the remaining glycine was calculated using the

molar ratio glycine-peroxodisulphate 1:1, from the results of formaldehyde and

the remaining glycine, the reaction between peroxodisulphate and glycine

followed a 1:1 stoichiometry.

Figure (3.1) absorbance values against glycine concentration.

0

0.2

0.4

0.6

0.8

1

1.2

1.4

0 1 2 3

Absorpance

Gly

cine

con

cent

ratio

n x1

03

32

3.1.2 Catalytic metal ions in deionized water

Deionized water was collected from Amipharma Laboratories for Drugs, and

tested for presence of catalytic metal ions such as Fe2+, Cu2+, Mn2+, and Cr3+ by

atomic absorption spectrophotometer.

In typical experiments, standard solutions of these metal ions were prepared,

1.0g of Iron dust and 1.0g of Cupper metal were dissolved separately in 50ml

nitric acid 5.0M; 1.0g of Manganese and 1.0g of Chromium powder were

dissolved separately in 50ml concentrated hydrochloric acid, after complete

dissolving the volumes were completed to 1L in volumetric flasks with

deionized water, each solution then contains 1000ppm of metal ion. Standard

addition method was used to determine concentrations of these metal ions, this

involves the addition of known amounts of the ion to be determined to a

number of aliquots of the sample solution; and the solutions thus obtained were

diluted to the same final solution, set of five 100ml volumetric flasks was

prepared for each metal ion, 10ml of the unknown sample were added to each

volumetric flask and 20, 40, 60, 80, and 100µL of metal ion solution were

added using micropipette, completed to the mark with deionized water; Test

solution was 100ml deionized water in 100ml volumetric flak. The absorbance

of test solution was first measured, and then each of prepared solutions were

examined in turn, leading up to the solution of highest concentration.

Table (3.4) concentration of metal ions and corresponding absorbance values.

Absorbance Values True Value (ppm)

Fe Mn Cr Cu 0 0.0017 0.0008 0.002 0.0007 2 0.3145 0.544 0.103 0.5091 4 0.6766 0.9786 0.1687 0.8517 6 0.7932 1.2459 0.2345 1.4049 8 0.8781 1.4264 0.3012 1.5379

10 1.0017 1.5423 0.3566 1.7274

33

The absorbance values were then plotted against the added concentration

values, fig (3.2) to fig (3.5) showed the absorbance values of these metal ions.

Straight lines were obtained for each metal ion and the straight line was

extrapolated to the concentration-axis the point where the axis is cut gives the

concentration of test solution.

Fig (3.3), calibration curve for Cr3+

ion (concentrations ppm).

Fig (3.2), calibration curve for Mn2+

ion (concentrations ppm).

Fig (3.5), calibration curve for Fe2+

ion (concentrations ppm).

Fig (3.4), calibration curve for Cu2+

ion (concentrations ppm).

34

Cr3+

Fe2+

Mn2+

Cu2+

00.20.40.60.8

11.21.41.61.8

2

-2 -1 0 1 2 3 4 5 6 7 8 9 10 11

Concentrations (ppm)

Abs

orba

nce

Fig (3.6), absorbance against the metal ion concentrations, extrapolation to

concentration axis.

The actual concentrations of these metal ions were calculated from the

extrapolation of the straight lines to cut at concentration axis.

Table (3.5), actual concentration of the metal ions in deionized water.

Metal ion Concentration (ppm)

Mn2+ 1.4

Fe2+ 1.4

Cr3+ 0.5

Cu2+ 0.8

35

3.2 Kinetic Results

3.2.1 Effect of varying peroxodisulphate and amino acid concentrations on

the reaction rate

To determine the effect of peroxodisulphate concentration on the reaction rate,

kinetic runs were performed by varying peroxodisulphate concentrations for

five different concentrations ranging from 0.0025 to 0.0125M by 0.0025M each

time, at constant concentration of amino acid 0.003M, temperature at 600C,

constant pH, and ionic strength at 0.25M using potassium sulfate 0.25M.

To determine the effect of amino acid, kinetic runs were carried out by varying

concentration of individual amino acid for five different concentrations ranging

from 0.001 to 0.005M, keeping the concentration of peroxodisulphate constant

at 0.005M, temperature at 600C, constant pH, and constant ionic strength at

0.25M. The residual peroxodisulphate concentration was calculated.

Logarithms of residual concentration of peroxodisulphate were plotted against

time and the rate constant was calculated from the slope of the graphs using

first order reaction equation

kobs = 2.303 x log [S2O82-]o

t [S2O82-]

and kobs = 2.303 x slope.

The rate of the reaction (R) was calculated using the following equation

R = ∆M/∆t mol.L-1.min-1.

Where, ∆M is the change in peroxodisulphate concentration in mol.L-1 and ∆t is

the change in time per minutes.

Tables containing of results logarithm of peroxodisulphate concentration were

presented in appendices, but figures from these tables were presented with each

amino acid concerned.

36

(3.2.1.1) Alanine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at alanine concentration =

0.003M, µ = 0.25M, pH = 2.00 and temperature 600C (for more details see

appendix A.1), is linear fig (3.7), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and showed in fig (3.7).

The reaction rate and average values were calculated (for more details see

appendix A.2). Table (3.6) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear with a positive slope fig (3.9), indicating that the

observed rate constant increased with increasing initial peroxodisulphate

concentration, The plots of average rate values versus initial peroxodisulphate

concentration is linear with positive slope fig (3.10), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed first order dependence of the reaction rate on peroxodisulphate

concentration.

The oxidation of alanine in the concentration ranging from 0.001 to 0.005M

was studied at constant peroxodisulphate concentration = 0.005M, µ = 0.25M,

pH =2.00 and temperature 600C (for more details see appendix A.2), the linear

plots of log peroxodisulphate concentration versus time was used to obtain the

observed rate constants fig (3.8).

Table (3.7) includes values of initial alanine concentration and their correspond-

ding average rate and observed rate constants.

The plot of observed rate constant kobs versus initial alanine concentration is

linear fig (3.11) and parallel to the concentration axis, and the plot of average

rate versus initial alanine concentration fig (3.12), is linear and parallel to the

concentration axis indicating zero order dependence in alanine concentration.

37

◊-Flask (1) kobs = 0.0004606■-Flask (2) kobs = 0.0009212∆-Flask (3) kobs = 0.0004606□-Flask (4) kobs = 0.0011515▲-Flask (5) kobs = 0.0013818

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+Lo

g[K 2

S 2O

8]

Fig (3.7) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of alanine at constant alanine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Alanine] = 0.003M, pH = 2.00, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0034545■-Flask (7) kobs = 0.0027636∆-Flask (8) kobs = 0.002303□-Flask (9) kobs = 0.0025333▲-Flask (10) kobs = 0.0027636

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.8) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of alanine at constant peroxodisulphate concentration. [Alanine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 2.00, µ = 0.25M, Temperature = 333K

38

Table (3.6) average rate and rate constant of varying peroxodisulphate concentrations on the oxidation of alanine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1. 104 x kobs min-1 0.0025 1.530 4.606 0.0050 3.912 9.212 0.0075 4.020 4.606 0.0100 10.02 11.515 0.0125 13.12 13.818

0

4

8

12

16

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s m

in-1

0

3

6

9

12

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.10) plot of average Rate x 106 m.l-1min-1against initial peroxodisulphate concentration M.

Fig (3.9) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.7) average rate and rate constant of varying alanine concentra-tions on the oxidation of alanine:

[Alanine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 8.571429 34.545 0.002 7.619048 27.636 0.003 6.904762 23.03 0.004 7.380952 25.333 0.005 7.857143 27.636

0

10

20

30

40

50

60

0 0.002 0.004 0.006

[Alanine]o m.l-1

104 x

Kob

s min

-1

0

2

4

6

8

10

12

14

16

18

0 0.002 0.004 0.006

[Alanine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.12) plot of average Rate x 106 m.l-1.min-1 against initial alanine concentration M.

Fig (3.11) plot of 104 x kobs min-1 against initial alanine concentration M.

39

(3.2.1.2) Argnine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at argnine concentration =

0.003M, µ = 0.25M, pH =1.94 and temperature 600C (for more details see

appendix B.1), is linear fig (3.13), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.13).

The reaction rate and their average values were calculated (for more details see

appendix B.2). Table (3.8) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant.The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.15),

indicating that the observed rate constant is independent to the initial peroxodi-

sulphate concentration. The plots of average rate values versus initial

peroxodisulphate concentration is linear with positive slope fig (3.16),

indicating that the rate increased with increasing initial peroxodisulphate

concentration. This confirmed first order dependence of the reaction rate on

peroxodisulphate concentration. The oxidation of argnine in the concentration

ranging from 0.001 to 0.005M was studied at constant peroxodisulphate

concentration = 0.005M, µ = 0.25M, pH =1.94 and temperature = 600C (for

more detailed see appendix B.2), the linear plots of log peroxodisulphate

concentration versus time was used to obtain the observed rate constant fig

(3.14). Table (3.9) includes values of initial argnine concentration and their

corresp-onding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial argnine concentration is

linear with a positive slope fig (3.17) indicating that the observed rate constant

is increased with increasing initial argnine concentration, and the plot of

average rate versus initial argnine concentration fig (3.18), is linear with

positive slope indicating that the average rate is increased with increasing initial

argnine concentration.

40

◊-Flask(1) kobs = 0.0020727■-Flak(2) kobs = 0.0025333∆-Flas(3) kobs = 0.0029939□-Flsk(4) kobs = 0.0029939▲-Flask(5) kobs = 0.0025333

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.13) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of argnine at constant argnine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Argnine] = 0.003M, pH = 1.94, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0020727■-Flask (7) kobs = 0.0027636∆-Flask (8) kobs = 0.0036848□-Flask (9) kobs = 0.0050666▲-Flask (10) kobs = 0.0057575

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.14) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of argnine at constant peroxodisulphate concentration. [Argnine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.94, µ = 0.25M, Temperature = 333K

41

Table (3.8) average rate and rate constant of varying peroxodisulphate concentra-tions on the oxidation of argnine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 3.334524 20.727 0.0050 9.261429 25.333 0.0075 13.69524 29.939 0.0100 19.2619 29.939 0.0125 20.91429 25.333

0

5

10

15

20

25

30

35

40

45

50

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.16) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.15) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.9) average rate and rate constant of varying argnine concentra-tions on the oxidation of argnine:

[Argnine]o M Average Rate x 106 m.l-1min-1. 104 x kobs min-1 0.001 5.952381 20.727 0.002 7.857143 27.636 0.003 9.52381 36.848 0.004 11.42857 50.666 0.005 13.09524 57.575

0

20

40

60

80

100

0 0.002 0.004 0.006

[Argnine]o m.l-1.

104 x

Kob

s min

-1

0

4

8

12

16

20

0 0.002 0.004 0.006

[Argnine]o m.l-1.

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.18) plot of average Rate x 106 m.l-1.min-1 against initial argnine concentration M.

Fig (3.18) plot of 104 x kobs min-1 against initial argnine concentration M.

42

(3.2.1.3) Asparagine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at asparagine concentration =

0.003M, µ = 0.25M, pH =1.91 and temperature = 600C (for more details see

appendix C.1), is linear fig (3.19), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.19).

The reaction rate and their average values were calculated (for more details see

appendix C.2). Table (3.10) includes values of initial peroxodisulphate concen-

tration and the corresponding values of average rate and observed rate constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.21),

indicating that the observed rate constant is independent to the initial peroxodi-

sulphate concentration, the plot of average rate versus initial peroxodisulphate

concentration is linear with positive slope fig (3.22), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed first order dependence of the reaction rate on peroxodisulphate

concentration. The oxidation of asparagine in the concentration ranging from

0.001 to 0.005M was studied at constant peroxodisulphate concentration =

0.005M, µ = 0.25 M, pH =1.91 and temperature = 600C (for more details see

appendix C.2), the linear plots of log peroxodisulphate concentration versus

time was used to obtain the observed rate constant fig (3.20). Table (3.11)

includes values of initial asparagine concentration and their corresponding

average rate and observed rate constant.

The plot of observed rate constant kobs versus initial asparagine concentration is

linear fig (3.23) and parallel to the concentration axis, and the plot of average

rate versus initial asparagine concentration fig (3.24), is linear and parallel to

the concentration axis indicating zero order dependence in asparagine

concentration.

43

◊-Flask (1)kobs = 0.0020727■-Flask(2)kobs = 0.0011515∆-Flask(3)kobs = 0.0016121□-Flask(4)kobs = 0.0011515▲-Flask(5)kobs = 0.0013818

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+Lo

g[K

2S2O

8]

Fig (3.19) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of asparagine at constant asparagine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Asparagine] = 0.003M, pH = 1.91 , µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0018424■-Flask (7) kobs = 0.0020727∆-Flask (8) kobs = 0.0025333□-Flask (9) kobs = 0.0020727▲-Flask (10) kobs = 0.0020727

0

0.1

0.2

0.3

0.4

0.5

0.6

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.20) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of asparagine at constant peroxodisulphate concentration. [Asparagine]= flask(1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.91, µ = 0.25M, Temperature = 333K

44

Table (3.10) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of asparagine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 3.095238 20.727 0.0050 4.285714 11.515 0.0075 7.857143 16.121 0.0100 9.047619 11.515 0.0125 11.90476 13.818

0

10

20

30

40

0 0.002 0.004 0.006 0.008 0.01 0.012 0.014

[K2S2O8]o m.l-1

104 x

kob

s m

in-1

0

2

4

6

8

10

12

14

0 0.002 0.004 0.006 0.008 0.01 0.012 0.014

[K2S2O8]o m.l-1Av

erag

e Ra

te x

10

6 m.l-1

.min

-1.

Fig (3.22) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.21) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.11) average rate and rate constant of varying asparagine concent-rations on the oxidation of asparagine:

[Asparagine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 4.761905 18.424 0.002 5.238095 20.727 0.003 6.190476 25.333 0.004 5.47619 20.727 0.005 5.47619 20.727

05

101520253035404550

0 0.002 0.004 0.006

[Asparagine]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

12

14

0 0.002 0.004 0.006

[Asparagine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.24) plot of average Rate x 106 m.l-1.min-1 against initial asparagine concentration M.

Fig (3.23) plot of 104 x kobs min-1 against initial asparagine concentration M

45

(3.2.1.4) Aspartic acid-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at aspartic acid concentration =

0.003M, µ = 0.25M, pH =1.84 and temperature = 600C (for more details see

appendix D.1), is linear fig (3.25), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.25).

The rate of reaction and their average values were calculated (for more details

see appendix D.2). Table (3.12) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.27), the plot

of average rate versus initial peroxodisulphate concentration is linear with

positive slope fig (3.28), indicating that the rate increased with increasing in

initial peroxodisulphate concentration. This confirmed the first order

dependence of the reaction rate on peroxodisulphate concentration.

The oxidation of aspartic acid in the concentration ranging from 0.001 to

0.005M was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M,

pH = 1.84 and temperature = 600C (for more details see appendix D.2), the

linear plots of log peroxodisulphate concentration versus time was used to

obtain the observed rate constant fig (3.26). Table (3.13) includes values of

initial aspartic acid concentration and their corresponding average rate and

observed rate constant. The plot of observed rate constant kobs versus initial

aspartic acid concentration is linear and parallel to the concentration axis fig

(3.29), and the plot of average rate versus initial aspartic acid concentration fig

(3.30), is linear and parallel to the concentration axis indicating that the average

rate and the observed rate constant are independent to the initial aspartic acid

concentration.

46

◊-Flask (1) kobs = 0.000921■-Flask(2) kobs = 0.0004606∆-Flask(3) kobs = 0.0009212□-Flask(4) kobs = 0.000690 ▲-Flask(5) kobs = 0.0011515

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+Lo

g[K

2S2O

8]

Fig (3.25) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of aspartic acid at constant aspartic acid concentr-ation. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M flask (4) = 0.010M, flask (5) = 0.0125M.[Aspartic acid] = 0.003M, pH = 1.84, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0018424■-Flask (7) kobs = 0.0020727∆-Flask (8) kobs = 0.0025333□-Flask (9) kobs = 0.0020727▲-Flask (10) kobs = 0.0020727

0

0.1

0.2

0.3

0.4

0.5

0.6

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.26) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of aspartic acid at constant peroxodisulphate concentr-ation.[Aspartic acid] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003 M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.84, µ = 0.25M, Temperature = 333K

47

Table (3.12) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of aspartic acid:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 3.333333 9.21 0.005 1.904762 4.606

0.0075 5.714286 9.212 0.01 5.238095 6.9

0.0125 9.761905 11.515

0

4

8

12

16

20

24

28

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

12

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.28) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.27) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.13) average rate and rate constant of varying aspartic acid concen-trations on the oxidation of aspartic acid:

[Aspartic acid]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 3.33 11.515 0.002 3.81 13.818 0.003 4.76 16.121 0.004 5.00 18.424 0.005 5.00 18.424

0

5

10

15

20

25

30

35

40

0 0.002 0.004 0.006

[Aspartic acid]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

12

0 0.002 0.004 0.006

[Aspartic acid]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.30) plot of average Rate x 106 m.l-1.min-1 against initial aspartic acid concentration M.

Fig (3.29) plot of 104 x kobs min-1 against initial aspartic acid concentration M.

48

(3.2.1.5) Cysteine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at cysteine concentration =

0.003M, µ = 0.25M at pH =1.00, and temperature = 600C and (for more details

see appendix E.1), is linear fig (3.31), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.31). The

reaction rate and their average values were calculated (for more details see

appendix E.2). Table (3.14) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.33), the plot

of average rate versus initial peroxodisulphate concentration is linear with

positive slope fig (3.34), indicating that the rate increased with increasing initial

peroxodisulphate concentration. This confirmed first order dependence of the

reaction rate on peroxodisulphate concentration. The oxidation of cysteine in

the concentration ranging from 0.001 to 0.005M was studied at constant

peroxodisulphate concentration = 0.005M, µ = 0.25M, pH =1.00 and

temperature = 600C (for more details see appendix E.2), the linear plots of log

peroxodisulphate concentration versus time was used to obtain the observed

rate constant fig (3.32). Table (3.15) includes values of initial cysteine

concentration and their correspo-nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial cysteine concentration is

linear and parallel to the concentration axis fig (3.35), and the plot of average

rate versus initial cysteine concentration fig (3.36), is linear and parallel to the

concentration axis indicating that the average rate and observed rate consrtant

are independent to the initial cysteine concentration.

49

◊-Flask (1) kobs = 0.0041454■-Flask(2) kobs =0.0039151∆-Flask(3) kobs =0.002303□-Flask(4) kobs = 0.0025333▲-Flask(5) kobs = 0.0034545

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.31) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of cysteine at constant cysteine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Cysteine] = 0.003M, pH = 1.00, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0018424■-Flask (7) kobs = 0.002303∆-Flask (8) kobs = 0.002303□-Flask (9) kobs = 0.0025333▲-Flask (10) kobs = 0.0032242

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.32) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of cysteine at constant peroxodisulphate concentration. [Cysteine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.00, µ = 0.25M, Temperature = 333K

50

Table (3.14) average rate and rate constant of varying peroxodisulphate concent-rations on the oxidation of cysteine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 5.00 41.454 0.0050 9.52381 39.151 0.0075 7.45 23.03 0.0100 12.8571 25.333 0.0125 19.5238 34.545

-5

515

2535

4555

6575

85

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.34) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.33) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.15) average rate and rate constant of varying cysteine concentra-tions on the oxidation of cysteine:

[Cysteine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 3.571429 18.424 0.002 4.047619 23.03 0.003 4.285714 23.03 0.004 4.52381 25.333 0.005 5.47619 32.242

0

10

20

30

40

50

0 0.002 0.004 0.006

[Cystine]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

12

0 0.002 0.004 0.006

[Cystine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.36) plot of average Rate x 106 m.l-1.min-1 against initial cysteine concentration M.

Fig (3.35) plot of 104 x kobs min-1 against initial cysteine concentration M

51

(3.2.1.6) Glutamic acid-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at glutamic acid concentration =

0.003M, µ = 0.25M, pH =2.02 and temperature = 600C (for more details see

appendix F.1), is linear fig (3.37), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.37).

The reaction rate and their average values were calculated (for more details see

appendix F.2). Table (3.16) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.39),

indicating that the observed rate constant independent to the initial

peroxodisulphate concentration; the plot of average rate versus initial

peroxodisulphate concentration is linear with positive slope fig (3.40),

indicating that the rate increased with increasing initial peroxodisulphate

concentration. This confirmed first order dependence of the reaction rate on

peroxodisulphate concentration. The oxidation of glutamic acid in the

concentration ranging from 0.001 to 0.005M was studied at peroxodisulphate

concentration = 0.005M, µ = 0.25M, pH =2.08 and temperature = 600C (for

more details see appendix F.2), the linear plots of log peroxodisulphate concen-

tration versus time was used to obtain the observed rate constant fig (3.38).

Table (3.17) includes values of initial glutamic acid concentration and their

corresponding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial glutamic acid concentration

is linear fig (3.41) and parallel to the concentration axis, and the plot of average

rate versus initial glutamic acid concentration fig (3.42), is linear and parallel to

the concentration axis indicating zero order dependence on glutamic acid

concentration.

52

◊-Flask (1) kobs = 0.0009212■-Flask(2) kobs = 0.0043757∆-Flask(3) kobs = 0.0052969□-Flask(4) kobs = 0.0018424▲-Flask(5) kobs = 0.0018424

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+Lo

g[K

2S2O

8]

Fig (3.37) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glutamic acid at constant glutamic acid concentra-tion. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Glutamic acid] = 0.003M, pH = 2.02, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0029939■-Flask (7)kobs = 0.0034545∆-Flask (8) kobs =0.0027636□-Flask (9) kobs = 0.0032242▲-Flask (10) kobs = 0.0027636

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.38) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glutamic acid at constant peroxodisulphate concentration. [Glutamic acid] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 2.02, µ = 0.25M, Temperature = 333K

53

Table (3.16) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of glutamic acid:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 2.1000 09.212 0.0050 13.550 43.757 0.0075 22.3262 52.969 0.0100 19.8860 18.424 0.0125 24.3191 18.424

0

30

60

90

120

0 0.005 0.01 0.015[K2S2O8]o m.l-1

104 x

kob

s m

in-1

0

5

10

15

20

25

30

0 0.005 0.01 0.015

[K2S2O8]o m.l-1Av

erag

e Ra

te x

10

6 m.l-1

.min

-1.

Fig (3.40) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.39) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.17) average rate and rate constant of varying glutamic acid conc-entrations on the oxidation of glutamic acid:

[Glutamic acid]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 7.857143 29.939 0.002 8.571429 34.545 0.003 7.380952 27.636 0.004 8.571429 32.242 0.005 7.142857 27.636

0

10

20

30

40

0 0.002 0.004 0.006

[Glutamic acid]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

0 0.002 0.004 0.006

[Glutamic acid]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.42) plot of average Rate x 106 m.l-1.min-1 against initial glutamic acid concentration M.

Fig (3.41) plot of 104 x kobs min-1 against initial glutamic acid concentration M

54

(3.2.1.7) Glutamine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at glutamine concentration =

0.003M, µ = 0.25m/, pH =3.51 and temperature = 600C (for more details see

appendix G.1), is linear fig (3.43), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.43).

The rate of reaction and their average values were calculated (for more details

see appendix G.2). Table (3.18) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.45), the plot

of average rate versus initial peroxodisulphate concentration is linear with

positive slope fig (3.46), indicating that the rate increased with increasing initial

peroxodisulphate concentration. This confirmed first order dependence of the

reaction rate on peroxodisulphate concentration. The oxidation of glutamine in

the concentration ranging from 0.001 to 0.005M was studied at peroxodisul-

phate concentration = 0.005M,µ = 0.25M, pH =3.51 and temperature = 600C

(for more details see appendix G.2), the linear plots of log peroxodisulphate

concentration versus time was used to obtain the observed rate constant fig

(3.44). Table (3.19) includes values of initial glutamine concentration and their

corresponding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial glutamine concentration is

linear fig (3.47) and parallel to the concentration axis, and the plot of average

rate versus initial glutamine concentration fig (3.48), is linear and parallel to the

concentration axis, indicating zero order dependence on glutamine concen-

tration.

55

◊-Flask (1) kobs = 0.0016121■-Flask (2) kobs = 0.0006909∆-Flask (3) kobs = 0.0029939□-Flask (4) kobs = 0.0016121▲-Flask (5) kobs = 0.0020727

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+Lo

g[K

2S2O

8]

Fig (3.43) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glutamine at constant glutamine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Glutamine] = 0.003M, pH = 3.51, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0020727■-Flask (7) kobs = 0.0011515∆-Flask (8) kobs = 0.0011515□-Flask (9) kobs = 0.0013818▲-Flask (10) kobs = 0.0009212

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250Time min.

3+Lo

g[K

2S2O

8]

Fig (3.44) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glutamine at constant peroxodisulphate concentration. [Glutamine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 3.51, µ = 0.25M, Temperature = 333K

56

Table (3.18) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of glutamine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 2.857143 16.121 0.0050 3.095238 6.909 0.0075 17.14286 29.939 0.0100 14.7619 16.121 0.0125 19.04762 20.727

0

10

20

30

40

50

60

70

0 0.005 0.01

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.005 0.01

[K2S2O8]o m.l-1A

vera

ge R

atex

106

m.l-1

.min

-1.

Fig (3.46) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.45) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.19) average rate and rate constant of varying glutamine concent-rations on the oxidation of glutamine:

[Glutamine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 11.42857 20.727 0.002 5.714286 11.515 0.003 5.47619 11.515 0.004 5.238095 13.818 0.005 4.047619 9.212

0

10

20

30

40

50

0 0.002 0.004 0.006

[Glutamine]o m.l-1

104 x

kob

s min

-1

0

4

8

12

16

20

0 0.002 0.004 0.006

[Glutamine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1m

in-1

.

Fig (3.48) plot of average Rate x 106 m.l-1.min-1 against initial glutamine concentration M.

Fig (3.47) plot of 104 x kobs min-1 against initial glutamine concentration M.

57

(3.2.1.8) Glycine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at glycine concentration =

0.003M, µ = 0.25M, pH =2.42 and temperature = 600C (for more details see

appendix H.1), was linear fig (3.49), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.49).

The rate of reaction and their average values were calculated (for more details

see appendix H.2). Table (3.20) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration was linear and parallel to the concentration axis fig (3.51), the

plot of average rate versus initial peroxodisulphate concentration was linear

with positive slope fig (3.52), indicating that the rate increased with increasing

initial peroxodisulphate concentration. This confirmed first order dependence of

the reaction rate on peroxodisulphate concentration.

The oxidation of glycine in the concentration ranging from 0.001 to 0.005M

was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M, pH =2.42

and temperature = 600C (for more details see appendix H.2), the linear plots of

log peroxodisulphate concentration versus time was used to obtain the observed

rate constant fig (3.50).

Table (3.21) includes values of initial glycine concentration and their correspo-

nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial glycine concentration was

linear fig (3.53) and parallel to the concentration axis, and the plot of average

rate versus initial glycine concentration fig (3.54), was linear and parallel to the

concentration axis indicating zero order dependence in glycine concentration.

58

◊-Flask (1) kobs = 0.0006909■-Flask(2) kobs = 0.0011515∆-Flask(3) kobs = 0.0011515□-Flask(4) kobs = 0.0016121▲-Flask(5) kobs = 0.0006909

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+L

og[K

2S2O

8]

Fig (3.49) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glycine at constant glycine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Glycine] = 0.003M, pH = 2.42, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0029939■-Flask (7) kobs = 0.0034545∆-Flask (8) kobs =0.0027636□-Flask (9) kobs = 0.0032242▲-Flask (10) kobs = 0.0027636

0

0.1

0.2

0.3

0.4

0.5

0.6

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.50) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of glycine at constant peroxodisulphate concentration. [Glycine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 2.42, µ = 0.25M, Temperature = 333K

59

Table (3.20) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of glycine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 1.65000 6.909 0.0050 6.847619 11.515 0.0075 6.559524 11.515 0.0100 9.338095 16.121 0.0125 8.797619 6.909

0

5

10

15

20

25

30

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

2

4

6

8

10

12

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.52) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.51) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.21) average rate and rate constant of varying glycine concentra-tions on the oxidation of glycine:

[Glycine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 7.857143 29.939 0.002 7.142857 25.333 0.003 5.714286 18.424 0.004 7.380952 27.636 0.005 5.238095 18.424

0

20

40

60

80

0 0.002 0.004 0.006

[Glycine]o m.l-1

104 x

kob

s min

-1

0

4

8

12

16

20

0 0.002 0.004 0.006

[Glycine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.54) plot of average Rate x 106 m.l-1.min-1 against initial glycine conc-entration M.

Fig (3.53) plot of 104 x kobs min-1 against initial glycine concentration M.

60

(3.2.1.9) Histidine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at histidine concentration =

0.003M, µ = 0.25M, pH =3.01 and temperature = 600C (for more details see

appendix I.1), is linear fig (3.55), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.55).

The rate of reaction and their average values were calculated (for more details

see appendix I.2). Table (3.22) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.57),

indicating that the observed rate constant independent to increasing with initial

peroxodisulphate concentration, the plot of average rate versus initial

peroxodisulphate concentration is linear with positive slope fig (3.58),

indicating that the rate increased with increasing initial peroxodisulphate

concentration. This confirmed first order dependence of the reaction rate on

peroxodisulphate concentration.

The oxidation of histidine in the concentration ranging from 0.001 to 0.005M

was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M, pH = 3.01

and temperature = 600C (for more details see appendix I.2), the linear plots of

log peroxodisulphate concentration versus time was used to obtain the observed

rate constant fig (3.56).

Table (3.23) includes values of initial histidine concentration and their corresp-

onding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial histidine concentration is

linear fig (3.59) and parallel to the concentration axis, and the plot of average

rate versus initial histidine concentration fig (3.60), is linear and almost parallel

to the concentration axis.

61

◊-Flask (1) kobs = 0.0059878■-Flask(2) kobs = 0.0062181∆-Flask(3) kobs = 0.0048363□-Flask(4) kobs = 0.0055272▲-Flask(5) kobs = 0.0034545

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Time min.

3+L

og[K

2S2O

8]

Fig (3.55) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of histidine at constant histidine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Histidine] = 0.003M, pH = 3.01, µ = 0.25M, Temperature = 333 K

◊-Flask (1) kobs = 0.0034545■-Flask(2) kobs = 0.0050666∆-Flask(3) kobs =0.0055272□-Flask(4) kobs = 0.0057575▲-Flask(5) kobs = 0.0052969

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.56) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of histidine at constant peroxodisulphate concentration. [Histidine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 3.01, µ = 0.25M, Temperature = 333 K

62

Table (3.22) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of histidine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 10.00 59.88 0.0050 14.49 62.18 0.0075 17.98 48.36 0.0100 25.04 55.27 0.0125 23.25 34.55

0

20

40

60

80

100

120

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104

x k o

bs m

in-1

0

5

10

15

20

25

30

0 0.005 0.01 0.015

[K2S2O8]o m.l-1Av

erag

e R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.58) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.57) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.23) average rate and rate constant of varying histidine concentra-tions on the oxidation of histidine:

[Histidine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 10.91667 34.545 0.002 13.79976 50.666 0.003 14.28048 55.272 0.004 14.10000 57.575 0.005 14.12833 52.969

0

20

40

60

80

100

0 0.002 0.004 0.006

[Histidine]o m.l-1

104 x

kob

s min

-1

-2

6

14

22

30

0 0.002 0.004 0.006[Histidine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1m

in-1

.

Fig (3.60) plot of average Rate x 106 m.l-1.min-1 against initial histidine concentration M.

Fig (3.59) plot of 104 x kobs min-1 against initial histidine concentration.

63

(3.2.1.10) Leucine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at leucine concentration =

0.003M, µ = 0.25M, pH =1.98 and temperature = 600C (for more details see

appendix J.1), is linear fig (3.61), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.61).

The rate of reaction and their average values were calculated (for more details

see appendix J.2). Table (3.24) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.63), the plot

of average rate versus initial peroxodisulphate concentration is linear with

positive slope fig (3.64), indicating that the rate increased with increasing initial

peroxodisulphate concentration. This confirmed first order dependence of the

reaction rate on peroxodisulphate concentration. The oxidation of leucine in the

concentration ranging from 0.001 to 0.005M was studied at peroxodisulphate

concentration = 0.005M, µ = 0.25M, pH =1.98 and temperature = 600C (for

more details see appendix J.2), the linear plots of log peroxodisulphate

concentration versus time was used to obtain the observed rate constant fig

(3.62). Table (3.25) includes values of initial leucine concentration and their

correspo-nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial leucine concentration is

linear fig (3.65) with small positive slope value, and the plot of average rate

versus initial leucine concentration fig (3.66), is linear with small positive slope

value indicating that the observed rate constant kobs and average rate are

increased somewhat with increasing initial leucine concentration

64

◊-Flask (1) kobs =0.0029939■-Flask(2) kobs = 0.0043757∆-Flask(3) kobs = 0.0041454□-Flask(4) kobs = 0.0048363▲-Flask(5) kobs = 0.0036848

0

0.25

0.5

0.75

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.61) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of leucine at constant leucine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Leucine] = 0.003M, pH = 1.98, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0009212■-Flask (7) kobs = 0.0011515∆-Flask (8) kobs = 0.0013818□-Flask (9) kobs = 0.0020727▲-Flask (10) kobs = 0.0029939

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250Time min.

3+Lo

g[K 2

S 2O

8]

Fig (3.62) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of leucine at constant peroxodisulphate concentration. [Leucine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.98, µ = 0.25M, Temperature = 333K

65

Table (3.24) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of leucine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 4.047619 29.939 0.0050 9.047619 43.757 0.0075 14.04762 41.454 0.0100 18.09524 48.363 0.0125 19.28571 36.848

0

20

40

60

80

100

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

atex

106 m

.l-1.m

in-1

.

Fig (3.64) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.63) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3. 25) average rate and rate constant of varying leucine concentra-tions on the oxidation of leucine:

[Leucine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 2.7 9.212 0.002 3.7 11.515 0.003 4.35 13.818 0.004 6.2 20.727 0.005 7.65 29.939

0

10

20

30

40

50

60

70

0 0.002 0.004 0.006

[Leucine]o m.l-1

104 xk

obsm

in-1

0

4

8

12

16

20

0 0.002 0.004 0.006

[Leucine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.66) plot of average Rate x 106 m.l-1.min-1 against initial leucine concentration M.

Fig (3.65) plot of 104 x kobs min-1 against initial leucine concentration M.

66

(3.2.1.11) Lysine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration rangingom 0.0025 to 0.0125M at lysine concentration = 0.003M,

µ = 0.25M , pH =1.92 and temperature = 600C (for more details see appendix

K.1), is linear fig (3.67), indicating first order dependence of the reaction rate

on peroxodisulphate concentration, the rate constants kobs were calculated from

the slope of the above plots and represented in fig (3.67).

The rate of reaction and their average values were calculated (for more details

see appendix K.2). Table (3.26) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.69),

indicating that the observed rate constant independent to increasing with initial

peroxodisulphate conce-ntration, the plot of average rate versus initial

peroxodisulphate concentration is linear with positive slope fig (3.70),

indicating that the rate increased with increasing initial peroxodisulphate

concentration. This confirmed first order dependence of the reaction rate on

peroxodisulphate concentration.

The oxidation of lysine in the concentration ranging from 0.001 to 0.005M was

studied at peroxodisulphate concentration = 0.005M, µ = 0.25M, pH =1.92 and

temperature = 600C (for more details see appendix K.2), the linear plots of log

peroxodisulphate concentration versus time was used to obtained the observed

rate constant fig (3.68).

Table (3.27) includes values of initial lysine concentration and their correspon-

ding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial lysine concentration is

linear fig (3.71) and parallel to the concentration axis, and the plot of average

rate versus initial lysine concentration fig (3.72), is linear and parallel to the

concentration axis, indicating zero order dependence on lysine concentration.

67

◊-Flask (1) kobs = 0.0059878■-Flask(2) kobs = 0.0050666∆-Flask(3) kobs = 0.0050666□-Flask(4) kobs = 0.004606▲-Flask(5) kobs = 0.0036848

00.10.20.30.40.50.60.70.80.9

1

0 50 100 150 200

Time min.

3+L

og[K

2S2O

8]

Fig (3.67) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of lysine at constant lysine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Lysine] = 0.003M, pH = 1.92, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0018424■-Flask (7) kobs = 0.0029939∆-Flask (8) kobs = 0.0052969□-Flask (9) kobs = 0.0025333▲-Flask (10) kobs = 0.0013818

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.68) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of lysine at constant lysine concentration. [Lysine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.92, µ = 0.25M, Temperature = 333K

68

Table (3.26) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of lysine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 6.934286 59.878 0.0050 10.94286 50.666 0.0075 18.32 50.666 0.0100 22.00476 46.06 0.0125 25.78095 36.848

0

20

40

60

80

100

120

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

30

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.70) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.69) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.27) average rate and rate constant of varying lysine concentra-tions on the oxidation of lysine:

[Lysine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 6.15 18.424 0.002 8.35 29.939 0.003 12.1 52.969 0.004 7.15 25.333 0.005 4.35 13.818

0

20

40

60

80

100

0 0.002 0.004 0.006

[Lysine]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

30

0 0.002 0.004 0.006

[Lysine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.72) plot of average Rate x 106 m.l-1.min-1 against initial lysine concentration M.

Fig (3.71) plot of 104 x kobs min-1 against initial lycine concentration M.

69

(3.2.1.12) Metheonine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at metheonine concentration =

0.003M, µ = 0.25M, pH =1.90 and temperature = 600C (for more details see

appendix L.1), is linear fig (3.73), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.73).

The rate of reaction and their average values were calculated (for more details

see appendix L.2). Table (3.28) includes values of initial peroxodisulphate

concentration and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear with positive slope fig (3.75), indicating that the

observed rate constant increased with increasing initial peroxodisulphate

concentration, the plot of average rate versus initial peroxodisulphate

concentration is linear with positive slope fig (3.76), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed the first order dependence of the reaction rate on peroxodisulphate

concentration.

The oxidation of metheonine in the concentration ranging from 0.001 to

0.005M was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M,

pH =1.90 and temperature = 600C (for more details see appendix L.2), the linear

plots of log peroxodisulphate concentration versus time was used to obtain the

observed rate constant fig (3.74).

Table (3.29) includes values of initial metheonine concentration and their

corresponding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial metheonine concentration

is linear fig (3.77) and almost parallel to the concentration axis, and the plot of

average rate versus initial metheonine concentration fig (3.78), is linear and

almost parallel to the concentration axis.

70

◊-Flask (1) kobs = 0.0006909■-Flask(2) kobs = 0.0016121∆-Flask(3) kobs = 0.0025333□-Flask(4) kobs = 0.0027636▲-Flask(5) kobs = 0.0029939

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.73) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of metheonine at constant metheonine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Metheonine] = 0.003M, pH = 1.90, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0006909■-Flask (7) kobs = 0.0016121∆-Flask (8) kobs = 0.0029939□-Flask (9) kobs = 0.0009212▲-Flask (10) kobs = 0.0016121

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.74) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of metheonine at constant peroxodisulphate concentration. [Metheonine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003 M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.90, µ = 0.25M, Temperature = 333K

71

Table (3.28) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of metheonine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 1.5 6.909 0.0050 4.65 16.121 0.0075 8.85 25.333 0.0100 13.05 27.636 0.0125 17.25 29.939

0

5

10

15

20

25

30

35

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

4

8

12

16

20

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.76) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.75) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.29) average rate and rate constant of varying metheonine concen-trations on the oxidation of metheonine:

[Metheonine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 1.95 6.909 0.002 5.05 16.121 0.003 8.1 29.939 0.004 2.8 9.212 0.005 4.85 16.121

0

15

30

45

60

75

0 0.002 0.004 0.006

[Metheonine]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.002 0.004 0.006

[Metheonine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.78) plot of average Rate x 106 m.l-1.min-1 against initial metheonine concentration M.

Fig (3.77) plot of 104 x kobs min-1 against initial metheonine concentration M.

72

(3.2.1.13) Phenylalanine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at phenylalanine concentration

= 0.003M and µ = 0.25M at constant pH =1.88 and temperature (for more

details see appendix M.1), is linear fig (3.79), indicating first order dependence

of the reaction rate on peroxodisulphate concentration, the rate constants kobs

were calculated from the slope of the above plots and represented in fig (3.79).

The rate of reaction and their average values were calculated (for more details

see appendix M.2). Table (3.30) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.81),

indicating that the observed rate constant independent to initial peroxodisul-

phate concentration, the plot of average rate versus initial peroxodisulphate

concentration is linear with positive slope fig (3.82), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed first order dependence of the reaction rate on peroxodisulphate

concentration.

The oxidation of phenylalanine in the concentration ranging from 0.001 to

0.005M was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M,

pH =1.88 and temperature = 600C (for more details see appendix M.2), the

linear plots of log peroxodisulphate concentration versus time was used to

obtain the observed rate constant fig (3.80).

Table (3.31) includes values of initial phenylalanine concentration and their

corresponding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial phenylalanine concen-

tration is linear fig (3.83) and parallel to the concentration axis, and the plot of

average rate versus initial phenylalanine concentration fig (3.84), is linear and

parallel to the concentration axis, indicating zero order dependence in

phenylalanine concentration.

73

◊-Flask (1) kobs = 0.002303■-Flask(2) kobs = 0.0029939∆-Flask(3) kobs = 0.0034545□-Flask(4)kobs = 0.0029939▲-Flask(5) kobs = 0.0029939

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Fig (3.79) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of phenyl alanine at constant phenylalanine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Phenyl alanine] = 0.003M, pH = 1.88, µ = 0.25M, Temperature = 333 K

◊-Flask (6) kobs = 0.0064484■-Flask (7) kobs = 0.0018424∆-Flask (8) kobs = 0.0018424□-Flask (9) kobs = 0.0018424▲-Flask (10) kobs = 0.0016121

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.80) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of phenyl alanine at constant peroxodisulphate concentration. [Phenyl alanine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M.[K2S2O8] = 0.005M, pH = 1.88, µ = 0.25M, Temperature = 333 K

74

Table (3.30) average rate and rate constant of varying peroxodisulphate concentrations on the oxidation of phenyl alanine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 4.05 23.03 0.0050 8.65 29.939 0.0075 14.8 34.545 0.0100 18.4 29.939 0.0125 22.6 29.939

0

10

20

30

40

50

60

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.82) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.81) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.31) average rate and rate constant of varying alanine concentra-tions on the oxidation of phenyl alanine:

[Phenyl alanine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 14.04762 64.484 0.002 6.190476 18.424 0.003 5.952381 18.424 0.004 6.190476 18.424 0.005 5.714286 16.121

0

10

20

30

40

50

60

70

0 0.002 0.004 0.006

[Phenylalanine]o m.l-1

104 x

kob

s min

-1

024

68

1012

1416

0 0.002 0.004 0.006

[Phenylalanine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.84) plot of average Rate x 106 m.l-1.min-1 against initial phenyl-alanine concentration M.

Fig (3.83) plot of 104 x kobs min-1 against initial phenyl alanine concentration M.

75

(3.2.1.14) Proline-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at proline concentration =

0.003M, µ = 0.25 M, pH =1.94 and temperature = 600C (for more details see

appendix N.1), is linear fig (3.85), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.85).

The rate of reaction and their average values were calculated (for more details

see appendix N.2). Table (3.32) includes values of initial peroxodisulphate

concen-tration and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.87), the plot

of average rate versus initial peroxodisulphate concentration is linear with

positive slope fig (3.88), indicating that the rate increased with increasing initial

peroxodisulphate concentration. This confirmed first order dependence of the

reaction rate on peroxodisulphate concentration. The oxidation of proline in the

concentration ranging from 0.001 to 0.005M was studied at peroxodisulphate

concentration = 0.005M, µ = 0.25M, pH =1.94 and temperature = 600C (for

more details see appendix N.2), the linear plots of log peroxodisulphate

concentration versus time was used to obtain the observed rate constant fig

(3.86).

Table (3.33) includes values of initial proline concentration and their correspo-

nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial proline concentration is

linear and parallel to the concentration axis fig (3.89), and the plot of average

rate versus initial proline concentration fig (3.90), is linear and parallel to the

concentration axis, indicating that the average rate and observed rate constant

are increased with initial proline concentration.

76

◊-Flask (1) kobs = 0.006909■-Flask(2) kobs = 0.0078302∆-Flask(3) kobs = 0.009212□-Flask(4) kobs = 0.009212▲-Flask(5) Kobs = 0.0085211

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.85) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of proline at constant proline concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Proline] = 0.003M, pH = 1.94, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs =0.0013818■-Flask (7)kobs = 0.0020727∆-Flask (8) kobs = 0.0029939□-Flask (9) kobs = 0.0023030▲-Flask (10) kobs = 0.0018424

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.86) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of proline at constant peroxodisulphate concentration. [Proline] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.94, µ = 0.25M, Temperature = 333K

77

Table (3.32) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of proline:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 12.5 69.09 0.0050 19.04762 78.302 0.0075 31.66667 92.12 0.0100 41.66667 92.12 0.0125 44.7619 85.211

0

30

60

90

120

150

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

15

30

45

60

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.88) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.87) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.33) average rate and rate constant of varying proline concentra-tions on the oxidation of proline:

[Proline]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 4.285714 13.818 0.002 6.190476 20.727 0.003 9.047619 29.939 0.004 6.904762 23.03 0.005 6.190476 18.424

0

10

20

30

40

50

60

0 0.001 0.002 0.003 0.004 0.005 0.006

[Proline]o m.l-1

104 x

kob

s m

in-1

0

4

8

12

16

20

0 0.001 0.002 0.003 0.004 0.005 0.006

[Proline]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.90) plot of average Rate x 106 m.l-1.min-1 against initial proline concentration M.

Fig (3.89) plot of 104 x kobs min-1 against initial proline concentration M.

78

(3.2.1.15) Serine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at serine concentration =

0.003M, µ = 0.25 M, pH =2.08 and temperature = 60 0C (for more details see

appendix O.1), is linear fig (3.91), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

calculated from the slope of the above plots and represented in fig (3.91).

The rate of reaction and their average values were calculated (for more details

see appendix O.2). Table (3.34) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.93),

indicating that the observed rate constant is independent to initial peroxodisul-

phate concentration, the plot of average rate versus initial peroxodisulphate

concentration is linear with positive slope fig (3.94), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed first order dependence of the reaction rate on peroxodisulphate

concentration. The oxidation of serine in the concentration ranging from 0.001

to 0.005M was studied at peroxodisulphate concentration = 0.005M, µ =

0.25M, pH =2.08 and temperature = 600C (for more details see appendix O.2),

the linear plots of log peroxodisulphate concentration versus time was used to

obtain the observed rate constant fig (3.92). Table (3.35) includes values of

initial serine concentration and their correspo-nding average rate and observed

rate constant.

The plot of observed rate constant kobs versus initial serine concentration is

linear and parallel to the concentration axis fig (3.95), and the plot of average

rate versus initial serine concentration fig (3.96), is linear and parallel to the

concentration axis, indicating that the average rate and observed rate constant

are increased with increasing serine concentration.

79

◊-Flask (1)kobs = 0.0025333 ■-Flask(2)kobs = 0.0011515∆-Flask(3)kobs = 0.0009212 □-Flask(4)kobs = 0.0006909 ▲-Flask(5)kobs = 0.000691

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250Time min.

3+L

og[K

2S2O

8]

Fig (3.91) plot of logarithm of peroxodisulphate concentration against time per

minute for the oxidation of serine at constant serine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Serine] = 0.003M, pH = 2.08, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs =0.0034545■-Flask (7) kobs = 0.0029939∆-Flask (8) kobs = 0.002303□-Flask (9) kobs = 0.0039151▲-Flask (10) kobs = 0.0059878

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.92) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of serine at constant peroxodisulphate concentration. [Serine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 2.08, µ = 0.25M, Temperature = 333K

80

Table (3.34) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of serine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 4.4000 25.333 0.0050 5.30381 11.515 0.0075 5.335714 9.212 0.0100 6.511905 6.909 0.0125 6.5000 6.91

0

10

20

30

40

50

60

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

1

2

34

5

6

7

8

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.94) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.93) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.35) average rate and rate constant of varying serine concentra-tions on the oxidation of serine:

[Serine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 9.285714 34.545 0.002 8.809524 29.939 0.003 6.904762 23.03 0.004 10.71429 39.151 0.005 13.80952 59.878

-10

10

30

50

70

90

110

130

0 0.002 0.004 0.006

[serine]o m.l-1

104 x

kob

s min

-1

0

6

12

18

24

30

0 0.002 0.004 0.006

[serine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.96) plot of average Rate x 106 m.l-1.min-1 against initial serine concentration M.

Fig (3.95) plot of 104 x kobs min-1 against initial serine concentration M.

81

(3.2.1.16) Threonine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at threonine concentration =

0.003M, µ = 0.25M, pH =1.99 and temperature = 600C (for more details see

appendix P.1), is linear fig (3.97), indicating first order dependence of the

reaction rate on peroxodisulphate concentration, the rate constants kobs were

obtained from the slope of the above plots and represented in fig (3.97).

The rate of reaction and their average values were calculated (for more details

see appendix P.2). Table (3.36) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.99),

indicating that the observed rate constant is independent to increasing with

initial peroxodisulphate concentration, the plot of average rate versus initial

peroxodisulphate concentration is linear with positive slope fig (3.100),

indicating that the rate increased with initial peroxodisulphate concentration.

This confirmed first order dependence of the reaction rate on peroxodisulphate

concentration. The oxidation of threonine in the concentration ranging from

0.001 to 0.005M was studied at peroxodisulphate concentration = 0.005M, µ =

0.25M, pH =1.99 and temperature = 600C (for more details see appendix P.2),

the linear plots of log peroxodisulphate concentration versus time was used to

obtain the observed rate constant fig (3.98). Table (3.37) includes values of

initial threonine concentration and their corresponding average rate and

observed rate constant.

The plot of observed rate constant kobs versus initial threonine concentration is

linear and parallel to the concentration axis fig (3.101), and the plot of average

rate versus initial threonine concentration fig (3.102), is linear and parallel to

the concentration axis indicating that the average rate and observed rate

constant are increased with initial threonine concentration.

82

◊-Flask (1)kobs = 0.004606■-Flask(2)kobs = 0.002303∆-Flask(3)kobs = 0.0050666□-Flask(4)kobs =0.0029939▲-Flask(5)kobs = 0.0039151

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.97) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of threonine at constant threonine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Threonine] = 0.003M, pH = 1.99, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0013818■-Flask (7) kobs = 0.0016121∆-Flask (8) kobs = 0.0013818□-Flask (9) kobs = 0.002303

▲-Flask (10) kobs = 0.0027636

0

0.3

0.6

0.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.98) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of threonine at constant peroxodisulphate concentration. [Threonine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.99, µ = 0.25M, Temperature = 333K

83

Table (3.36) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of threonine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 7.857143 46.06 0.0050 8.809524 23.03 0.0075 21.42857 50.666 0.0100 20.71429 29.939 0.0125 27.38095 39.151

0

20

40

60

80

100

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

25

30

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1

Fig (3.100) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.99) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.37) average rate and rate constant of varying threonine concentra-tions on the oxidation of threonine:

[Threonine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 5.714286 16.121 0.002 5.714286 16.121 0.003 4.52381 13.818 0.004 7.142857 23.03 0.005 8.095238 27.636

0

15

30

45

60

75

0 0.002 0.004 0.006

[Threonine]o m.l-1

104 x

kob

s min

-1

0

4

8

12

16

20

0 0.002 0.004 0.006

[Threonine]o m.l-1

Ave

rage

Rat

e x

106

m.l-1

.min

-1.

Fig (3.102) plot of average Rate x 106 m.l-1.min-1 against initial threonine concentration M.

Fig (3.101) plot of 104 x kobs min-1 against initial threonine concentration M.

84

(3.2.1.17) Tyrosine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranged from 0.0025 to 0.0125M at tyrosine concentration =

0.003M, µ = 0.25 M at constant pH =1.46 and temperature = 600C (for more

details see appendix Q.1), is linear fig (3.103), indicating first order dependence

of the reaction rate on peroxodisulphate concentration, the rate constants kobs

were obtained from the slope of the above plots and represented in fig (3.103).

The rate of reaction and their average values were calculated (for more details

see appendix Q.2). Table (3.38) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant. The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.105),

indicating that the observed rate constant is independent to initial peroxodisul-

phate concentration, the plot of average rate versus initial peroxodisulphate

concentration was linear with positive slope fig (3.106), indicating that the rate

increased with increasing initial peroxodisulphate concentration. This

confirmed the first order dependence of the reaction rate on peroxodisulphate

concentration.

The oxidation of tyrosine in the concentration ranged from 0.001 to 0.005M

was studied at peroxodisulphate concentration = 0.005M, µ = 0.25M, pH =1.46

and temperature = 600C (for more details see appendix Q.2), the linear plots of

log peroxodisulphate concentration versus time was used to obtain the observed

rate constant fig (3.104).

Table (3.39) includes values of initial tyrosine concentration and their correspo-

nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial tyrosine concentration is

linear fig (3.107) and parallel to the concentration axis, and the plot of average

rate versus initial tyrosine concentration fig (3.108), is linear and parallel to the

concentration axis indicating zero order dependence on tyrosine concent-ration.

85

◊-Flask (1) kobs = 0.0041454■-Flask(2) kobs = 0.0062181∆-Flask(3) kobs = 0.0062181□-Flask(4) kobs = 0.004606▲-Flask(5) kobs = 0.0041454

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.103) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of tyrosine at constant tyrosine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Tyrosine] = 0.003M, pH = 1.46, µ = 0.25M, Temperature = 333K

◊-Flask (6) kobs = 0.0018424■-Flask (7) kobs = 0.0062181∆-Flask (8) kobs = 0.0016121□-Flask (9) kobs = 0.0020727▲-Flask (10) kobs = 0.002303

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.104) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of tyrosine at constant peroxodisulphate concentration. [Tyrosine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M.[K2S2O8] = 0.005M, pH = 1.46, µ = 0.25M, Temperature = 333K

86

Table (3.38) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of tyrosine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 7.619048 41.454 0.0050 15.47619 62.181 0.0075 23.09524 62.181 0.0100 26.19048 46.06 0.0125 28.09524 41.454

0

20

40

60

80

100

120

0 0.005 0.01 0.015

[K2S2O8]o m.l-1

104

x k o

bs m

in-1

0

5

10

15

20

25

30

35

0 0.005 0.01 0.015

[K2S2O8]o m.l-1A

vera

ge R

ate

x 10

6 m.l-1

.min

-1.

Fig (3.106) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.105) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.39) average rate and rate constant of varying alanine concentra-tions on the oxidation of tyrosine:

[Tyrosine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 6.428571 18.424 0.002 6.190476 62.181 0.003 5.714286 16.121 0.004 6.666667 20.727 0.005 7.619048 23.03

0

20

40

60

80

100

0 0.001 0.002 0.003 0.004 0.005 0.006

[Tyrosine]o m.l-1

104 x

kob

s min

-1

0

5

10

15

20

0 0.001 0.002 0.003 0.004 0.005 0.006

[Tyrosine]o m.l-1

Ave

rage

Rat

e x

106 m

.l-1.m

in-1

.

Fig (3.108) plot of average Rate x 106 m.l-1.min-1 against initial tyrosine concentration M.

Fig (3.107) plot of 104 x kobs min-1 against initial tyrosine concentration M.

87

(3.2.1.18) Valine-Peroxodisulphate system

The plot of log peroxodisulphate concentration versus time at peroxodisulphate

concentration ranging from 0.0025 to 0.0125M at valine concentration =

0.003M, µ = 0.25 M at constant pH =1.88 and temperature = 600C (for more

details see appendix R.1), is linear fig (3.109), indicating first order dependence

of the reaction rate on peroxodisulphate concentration, the rate constants kobs

were obtained from the slope of the above plots and represented in fig (3.109).

The rate of reaction and their average values were calculated (for more details

see appendix R.2). Table (3.40) includes values of initial peroxodisulphate

concentr-ation and the corresponding values of average rate and observed rate

constant.

The plot of observed rate constant kobs versus initial peroxodisulphate

concentration is linear and parallel to the concentration axis fig (3.111),

indicating that the observed rate constant independent to increasing in initial

peroxodisulphate concentration, the plot of average rate versus initial

peroxodisulphate concentration is linear with positive slope fig (3.112),

indicating that the rate increased with increasing initial peroxodisulphate

concentration. This confirmed first order dependence of the reaction rate on

peroxodisulphate concentration.

The oxidation of valine in the concentration ranging from 0.001 to 0.005M was

studied at peroxodisulphate concentration = 0.005M, µ = 0.25M, pH =1.88 and

temperature = 600C (for more details see appendix R.2), the linear plots of log

peroxodisulphate concentration versus time was used to obtain the observed

rate constant fig (3.110).

Table (3.41) includes values of initial valine concentration and their correspo-

nding average rate and observed rate constant.

The plot of observed rate constant kobs versus initial valine concentration is

linear fig (3.113) and parallel to the concentration axis, and the plot of average

rate versus initial valine concentration fig (3.114), is linear and parallel to the

concentration axis indicating zero order dependence in valine concentration.

88

◊-Flask (1) kobs = 0.0066787■-Flask(2) kobs = 0.0062181∆-Flask(3) kobs = 0.0013818□-Flask(4) kobs = 0.0066787▲-Flask(5) kobs = 0.004606

0

0.2

0.4

0.6

0.8

1

1.2

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.109) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of valine at constant valine concentration. [K2S2O8] = flask (1) = 0.0025M, flask (2) = 0.005M, flask (3) = 0.0075M, flask (4) = 0.010M, flask (5) = 0.0125M. [Valine] = 0.003M, pH = 1.88, µ = 0.25M, Temperature = 333 K

◊-Flask (6) kobs = 0.0036848■-Flask (7) kobs = 0.0062181∆-Flask (8) kobs = 0.0039151□-Flask (9) kobs = 0.0032242▲-Flask (10) kobs = 0.0011515

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.110) plot of logarithm of peroxodisulphate concentration against time per minute for the oxidation of valine at constant peroxodisulphate concentration. [Valine] = flask (1) = 0.001M, flask (2) = 0.002M, flask (3) = 0.003M, flask (4) = 0.004M, flask (5) = 0.005M. [K2S2O8] = 0.005M, pH = 1.88, µ = 0.25M, Temperature = 333 K

89

Table (3.40) average rate and rate constant of varying peroxodisulphate concentr-ations on the oxidation of valine:

[K2S2O8]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.0025 11.8 66.787 0.0050 9.1 62.181 0.0075 6.35 13.818 0.0100 33.65 66.787 0.0125 31.95 46.06

0102030405060708090

100

0 0.005 0.01 0.015[K2S2O8]o m/l

104 x

Kob

s min

-1

0102030405060708090

100

0 0.005 0.01 0.015

[K2S2O8]o m/lA

vera

ge R

ate

x 10

6 m/l/

min

.

Fig (3.112) plot of average Rate x 106 m.l-1.min-1 against initial peroxodisulphate concentration M.

Fig (3.111) plot of 104 x kobs min-1 against initial peroxodisulphate concentration M.

Table (3.41) average rate and rate constant of varying valine concentra-tions on the oxidation of valine:

[Valine]o M Average Rate x 106 m.l-1min-1 104 x kobs min-1 0.001 12.2 36.848 0.002 12.5 62.181 0.003 12.0 39.151 0.004 11.0 32.242 0.005 4.25 11.515

0

30

60

90

120

0 0.001 0.002 0.003 0.004 0.005 0.006

[Valine]o m.l-1

104 x

kob

s m

in-1

0

5

10

15

20

25

30

0 0.001 0.002 0.003 0.004 0.005 0.006

[Valine]o m.l-1

Ave

rage

Rat

e x

106

m.l-1

.min

-1.

Fig (3.114) plot of average Rate x 106 m.l-1.min-1 against initial valine concentration M.

Fig (3.113) plot of 104 x kobs min-1 against initial valine concentration M.

90

3.2.2 Effect of Temperature

The effect of temperature on the reaction was studied in the temperature

range (60-800C), applying the following conditions:

[K2S2O8] = 0.0075M, [Amino acid] = 0.003M, µ = 0.25M with individual

pH value for each amino acid solution. The logarithm of peroxodisulphate

concentration was calculated with each amino acid and these results were

represent in the appendices from appendix (S.1) to appendix (S.9). Figure

(3.115) to figure (3.133) show the observed rate constants kobs which were

obtained from the slope of the plot of log peroxodisulphate concentration

versus time.

◊-Flask at 600C kobs = 0.0016121■-Flask at 650C kobs = 0.0064484∆-lask at 700C kobs = 0.011515

□-Flask at 750C kobs = 0.0135877

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊- Flak at 600C kobs = 0.0018424■-Flak at 650C kobs = 0.0059878∆-Flas at 700C kobs = 0.0124362□-Flsk at 750C kobs = 0.0131271

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.116) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Argnine.

Fig (3.115) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Alanine.

91

◊-Flask at 600C kobs = 0.0016121■-Flask at 650C kobs = 0.0032242∆-lask at 700C kobs = 0.0041454□-Flask at 750C kobs = 0.0082908▲-Flask at 800C kobs = 0.0103635

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.0020727∆-Flaskat 700C kobs = 0.0064484□-Flaskat 750C kobs = 0.0089817▲-Flaskat 800C kobs = 0.013818

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250Time min.

3+L

og[K

2 S2 O

8 ]

Fig (3.118) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Aspartic acid.

Fig (3.117) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Asparagine

◊-Flask at 600C kobs = 0.0025333■-Flaskat 650C kobs = 0.0064484∆-Flaskat 700C kobs = 0.0112847

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.0032242∆-Flaskat 700C kobs = 0.0041454□-Flaskat 750C kobs = 0.0080605▲-Flaskat 800C kobs = 0.0147392

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.120) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Glutamic acid.

Fig (3.119) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Cysteine.

92

◊-Flask at 600C kobs = 0.0029939■-Flask at 650C kobs = 0.0029939∆-lask at 700C kobs = 0.0013818□-Flask at 750C kobs = 0.0085211▲-Flask at 800C kobs = 0.0145089

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K2

S2O

8 ]

◊-Flask at 600C kobs = 0.0018424■-Flaskat 650C kobs = 0.0018424∆-Flaskat 700C kobs = 0.0154301□-Flaskat 750C kobs = 0.0251027

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250

Time min. 3

+Log

[K2 S

2 O8 ]

Fig (3.122) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Glycine.

Fig (3.121) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Glutamine.

◊-Flask at 600C kobs = 0.0004606■-Flaskat 650C kobs = 0.0020727∆-Flaskat 700C kobs = 0.0055272□-Flaskat 750C kobs = 0.0089817▲-Flaskat 800C kobs = 0.0112847

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250Time min.

3+L

og[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0004606■-Flaskat 650C kobs = 0.0043757∆-Flaskat 700C kobs = 0.009212□-Flaskat 750C kobs = 0.0147392

0

0.2

0.4

0.6

0.8

0 50 100 150 200 250Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.124) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Leucine.

Fig (3.123) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Histidine.

93

◊-Flask at 600C kobs = 0.0041454■-Flaskat 650C kobs = 0.0094423∆-Flaskat 700C kobs = 0.0177331

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0020727■-Flaskat 650C kobs = 0.004606∆-Flaskat 700C kobs = 0.0066787□-Flaskat 750C kobs = 0.0094423▲-Flaskat 800C kobs = 0.0099029

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.126) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Methionine.

Fig (3.125) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Lysine.

◊-Flask at 600C kobs = 0.002303■-Flaskat 650C kobs = 0.0085211∆-Flaskat 700C kobs = 0.018424

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0013818■-Flaskat 650C kobs = 0.0034545∆-Flaskat 700C kobs = 0.0089817□-Flaskat 750C kobs = 0.0124362

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.128) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Proline.

Fig (3.127) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Phenylalanine.

94

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.0013818∆-Flaskat 700C kobs = 0.0016121□-Flaskat 750C kobs = 0.0039151▲-Flaskat 800C kobs = 0.0154301

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.0048363∆-Flaskat 700C kobs = 0.0105938□-Flaskat 750C kobs = 0.0131271

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]Fig (3.130) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Threonine.

Fig (3.129) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Serine.

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.0013818∆-Flaskat 700C kobs = 0.0087514□-Flaskat 750C kobs = 0.0128968

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

◊-Flask at 600C kobs = 0.0009212■-Flaskat 650C kobs = 0.002303∆-Flaskat 700C kobs = 0.0087514□-Flaskat 750C kobs = 0.0119756

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.132) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Valine.

Fig (3.131) plot of logarithm of peroxodisulphate concentration against time per minutes at different temperatures for the oxidation of Tyrosine.

95

Table (3.42) shows the variation of values of Ln kobs with temperature for the

oxidation of Alanine (kobs values were taken from figure (3.115)).

Temperature K 1/T x 103 K-1 kobs min-1 Ln kobs 333 3.003 0.001612 -6.430218 338 2.9586 0.006448 -5.043923

343 2.9155 0.011515 -4.464105 348 2.8736 0.013588 -4.29859

Table (3.43) shows the variation of values of Ln kobs with temperature for the

oxidation of Argnine (kobs values were taken from figure (3.116)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.001842 -6.2966862 338 2.9586 0.005988 -5.1180312 343 2.9155 0.012436 -4.3871437 348 2.8736 0.013127 -4.3330765

Table (3.44) shows the variation of values of Ln kobs with temperature for the

oxidation of Aspargine (kobs values were taken from figure (3.117)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.001612 -6.43022 338 2.9586 0.003224 -5.73707 343 2.9155 0.004145 -5.48576 348 2.8736 0.008291 -4.79261 353 2.8329 0.010364 -4.56947

Table (3.45) shows the variation of values of Ln kobs with temperature for the

oxidation of Aspartic acid (kobs values were taken from figure (3.118)).

Temperature K 1/T x 103 K-1 kobs min-1 Ln kobs 333 3.003 0.000921 -6.98983 338 2.9586 0.002073 -6.1789 343 2.9155 0.006448 -5.04392 348 2.8736 0.008982 -4.71257 353 2.8329 0.013818 -4.28178

96

Table (3.46) shows the variation of values of Ln kobs with temperature for the

oxidation of Cysteine (kobs values were taken from figure (3.119)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.002533 -5.97835 338 2.9586 0.006448 -5.04399 343 2.9155 0.011285 -4.48428

Table (3.47) shows the variation of values of Ln kobs with temperature for the

oxidation of Glutamic acid (kobs values were taken from figure (3.120)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000921 -6.98983 338 2.9586 0.003224 -5.73707 343 2.9155 0.004145 -5.48576 348 2.8736 0.008061 -4.82078 353 2.8329 0.014739 -4.21724

Table (3.48) shows the variation of values of Ln kobs with temperature for the

oxidation of Glutamine (kobs values were taken from figure (3.121)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.002994 -5.81118 338 2.9586 0.002994 -5.81118 343 2.9155 0.001382 -6.58437 348 2.8736 0.008521 -4.76521 353 2.8329 0.014509 -4.23299

Table (3.49) shows the variation of values of Ln kobs with temperature for the

oxidation of Glycine (kobs values taken were from figure (3.122)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.001842 -6.29669 338 2.9586 0.001842 -6.29669 343 2.9155 0.01543 -4.17144 348 2.8736 0.025103 -3.68478

97

Table (3.50) shows the variation of values of Ln kobs with temperature for the

oxidation of Hestidine (kobs values were taken from figure (3.123)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000461 -7.68298 338 2.9586 0.002073 -6.1789 343 2.9155 0.005527 -5.19807 348 2.8736 0.008982 -4.71257 353 2.8329 0.011285 -4.48431

Table (3.51) shows the variation of values of Ln kobs with temperature for the

oxidation of Lecine (kobs values were taken from figure (3.124)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000461 -7.68298 338 2.9586 0.004376 -5.43169 343 2.9155 0.009212 -4.68725 348 2.8736 0.014739 -4.21724

Table (3.52) shows the variation of Ln kobs with temperature for the oxidation of

Lysine (kobs values were taken from figure (3.125)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.004145 -5.48576 338 2.9586 0.009442 -4.66256 343 2.9155 0.017733 -4.03232

Table (3.53) shows the variation of values of Ln kobs with temperature for the

oxidation of Metheonine (kobs values were taken from figure (3.126)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.002073 -6.1789 338 2.9586 0.004606 -5.3804 343 2.9155 0.006679 -5.00883 348 2.8736 0.009442 -4.66256 353 2.8329 0.009903 -4.61493

98

Table (3.54) shows the variation of values of Ln kobs with temperature for the

oxidation of Phenyl alanine (kobs values were taken from figure (3.127)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.002303 -6.07354 338 2.9586 0.008521 -4.76521 343 2.9155 0.018424 -3.9941

Table (3.55) shows the variation of values of Ln kobs with temperature for the

oxidation of Proline (kobs values were taken from figure (3.128)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.001382 -6.58437 338 2.9586 0.003455 -5.66808 343 2.9155 0.008982 -4.71257 348 2.8736 0.012436 -4.38714

Table (3.56) shows the variation of values of Ln kobs with temperature for the

oxidation of Serine (kobs values were taken from figure (3.129)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000921 -6.98983 338 2.9586 0.001382 -6.58437 343 2.9155 0.001612 -6.43022 348 2.8736 0.003915 -5.54291 353 2.8329 0.01543 -4.17144

Table (3.57) shows the variation of values of Ln kobs with temperature for the

oxidation of Threonine (kobs values were taken from figure (3.130)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000921 -6.98983 338 2.9586 0.004836 -5.33161 343 2.9155 0.010594 -4.54749 348 2.8736 0.013127 -4.33308

99

Table (3.58) shows the variation of values of Ln kobs with temperature for the

oxidation of Tyrosine (kobs values were taken from figure (3.131)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000921 -6.98983 338 2.9586 0.001382 -6.58437 343 2.9155 0.008751 -4.73854 348 2.8736 0.012897 -4.35078

Table (3.59) shows the variation of values of Ln kobs with temperature for the

oxidation of Valine (kobs values were taken from figure (3.132)).

Temperature K 1/T x 103 K-1 kobs min-1 Lnkobs 333 3.003 0.000921 -6.98983 338 2.9586 0.002303 -6.07354 343 2.9155 0.008751 -4.73854 348 2.8736 0.011976 -4.42488

Figure (3.133) to figure (3.151) show the Arrhenius plot of Ln kobs against 1/T K-1

for the reaction of peroxodisulphate with each amino acid. A considerable

extrapolation to 1/T = 0 would be very necessary to determine the constant Ln A

from the intercept of this line, and this values were stated in the equation of the

straight line.

y = -16.251x + 42.681

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 3 3.11/Tx103 K-1

Lnk o

bs

y = -15.426x + 40.282

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 3 3.11/T x 103 K-1

Lnko

bs

Fig (3.133) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Alanine.

Fig (3.134) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Argnine.

100

y = -10.987x + 26.644

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 3 3.1

1/T x 103 K-1

Lnk o

bs

y = -16.234x + 41.91

-8

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 3

1/T x 103 K-1

Lnk o

bs

Fig (3.135) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Aspargine.

Fig (3.136) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Aspartic acid

y = -17.095x + 45.416

-8-7-6-5-4-3-2-10

2.9 2.95 31/T x 103 K-1

Lnko

bs

y = -15.223x + 38.95

-8

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 31/T x 103 K-1

Lnk o

bs

Fig (3.137) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Cystine.

Fig (3.138) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Glutamic acid.

y = -9.7564x + 23.016

-7

-6

-5

-4

-3

-2

-1

02.8 2.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnk o

bs

y = -23.057x + 62.62

-7

-6

-5

-4

-3

-2

-1

02.85 2.95 3.05

1/T x 103 K-1

Lnk o

bs

Fig (3.139) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Glutamine.

Fig (3.140) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Glycine.

101

y = -18.592x + 48.577

-9-8-7-6-5-4-3-2-10

2.8 2.9 31/T x 103 K-1

Lnk o

bs

y = -25.949x + 70.724

-9

-8

-7

-6

-5

-4

-3

-2

-1

02.85 2.95 3.05

1/T x 103 K-1

Lnk o

bs

Fig (3.141) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Hestidine.

Fig (3.142) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of leucine.

y = -16.62x + 44.453

-6

-5

-4

-3

-2

-1

02.85 2.95 3.051/T x 103 K-1

Lnk o

bs

y = -9.094x + 21.356

-7

-6

-5

-4

-3

-2

-1

02.7 2.8 2.9 3 3.1

1/T x 103 K-1

Lnk o

bs

Fig (3.143) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Lysine.

Fig (3.144) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Metheonine.

y = -23.794x + 65.462

-7

-6

-5

-4

-3

-2

-1

02.9 2.95 3

1/T x 103 K-1

Lnko

bs

y = -17.535x + 46.174

-7

-6

-5

-4

-3

-2

-1

02.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnko

bs

Fig (3.145) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Phenylalanine.

Fig (3.146) Arrhenius plot of Ln

kobs against 1/T K-1 for the oxida-

tion of Proline.

102

y = -15.612x + 39.591

-8

-7

-6

-5

-4

-3

-2

-1

02.8 2.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnko

bs

y = -20.391x + 54.603

-8

-7

-6

-5

-4

-3

-2

-1

02.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnko

bs

Fig (3.147) Arrhenius plot of Ln kobs

against 1/T K-1 for the oxidation of

Serine.

Fig (3.148) Arrhenius plot of Ln kobs

against 1/T K-1 for the oxidation of

Threonine.

y = -22.632x + 60.82

-8

-7-6

-5-4

-3

-2-1

02.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnko

bs

y = -20.973x + 56.054

-8

-7

-6

-5

-4

-3

-2

-1

02.85 2.9 2.95 3 3.05

1/T x 103 K-1

Lnko

bs

Fig (3.149) Arrhenius plot of Ln

kobs against 1/T K-1 for the

oxidation of Tyrosine.

Fig (3.150) Arrhenius plot of Ln

kobs against 1/T K-1 for the

oxidation of Valine.

103

3.2.3 Test for free radicals

Induced polymerization with acrylonitrile was used to detect free radicals in

the oxidation of amino acids by peroxodisulphate in dark and inert

conditions.

The reaction mixtures each of 0.005M of peroxodisulphate, 0.005M of

individual amino acid, 5.0ml of acrylonitrile and deionized water to

complete the total volume 100ml, were kept in the dark for 3 hours, the

milky appearance were observed in all amino acids studied, indicating an in

situ formation of free radicals; proper control experiment was performed as

above except amino acid was omitted and the milky appearance was

observed after 10 hours indicating that the oxidation of water itself was

taking place.

3.2.4 Effect of solvent composition

All twenty essential amino acids were grouped according to the

characteristics of the side chains as follows:

Aliphatic amino acids:

These include Alanine, Glycine, Isoleucine, Leucine, Proline, and Valine.

Aromatic amino acids:

These include, Phnylalanine, Tryptophan, and Tyrosine.

Acidic amino acids:

These include Aspartic acid, and Glutamic acid.

Basic amino acids:

These include Argnine, Histidine, and Lysine.

Hydroxylic amino acids:

These include Serine, and Threonine.

Sulphur-cotaining amino acids:

These include Cysteine, and Methionine.

Amidic (containing amide group) amino acids:

These include Asparagine, Glutamine.

104

To study the effect of solvent composition on the reaction rate, effect of

ionic strength, and effect of catalyst and to avoid the repetition, one amino

acid was selected to stand for its group, the following seven amino acids

have been selected to stand for the seven groups; Alanine, Phenylalanine,

Glutamic acid, Lysine, Serine, Cysteine, and Asparagine.

The dielectric constant (or permittivity, D) of solvent medium was altered by

the addition of acetic acid to the reaction mixtures (5- 15 % v/v).

This effect on the reaction rate was studied applying the following

conditions: Temperature = 600C, [K2S2O8] = 0.005M, [Amino acid] =

0.003M, µ = 0.25M.

The observed rate constants of the reactions were obtained from plots of

logarithm of peroxodisulphate concentration against time, fig (3.151) to fig

(3.157) for more details see appendix (T.1) to appendix (T.4).

Logarithm of observed rate constant was plotted against the inverse of

dielectric constant, fig (3.158) to fig. (3.164).

◊-Flask (1) kobs = 0.0029939■-Flask (2) kobs =0.0025333∆-Flask (3) kobs = 0.0027636

00.10.20.30.40.50.60.70.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0036848■-Flask (2) kobs = 0.0029939∆-Flask (3) kobs = 0.002303

00.10.20.30.40.50.60.70.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.152) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Asparagine.

Fig (3.151) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Alanine.

105

◊-Flask (1) kobs = 0.0055272■-Flask (2) kobs = 0.0052969∆-Flask (3) kobs = 0.0052969

00.1

0.20.3

0.40.5

0.60.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0032242■-Flask (2) kobs = 0.0032242∆-Flask (3) kobs = 0.0025333

0

0.2

0.4

0.6

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.154) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Glutamic acid.

Fig (3.153) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Cystine.

◊-Flask (1) kobs = 0.0055272■-Flask (2) kobs = 0.0043757∆-Flask (3) kobs = 0.0027636

0

0.2

0.4

0.6

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0145089■-Flask (2) kobs = 0.0145089∆-Flask (3) kobs = 0.0112847

0

0.2

0.4

0.6

0.8

1

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.156) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Phenylalanine.

Fig (3.155) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Lysine.

◊-Flask (1) kobs = 0.0087514■-Flask (2) kobs = 0.0087514∆-Flask (3) kobs = 0.0075999

00.10.20.30.40.50.60.70.80.9

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.157) Plot of logarithm of peroxodisulphate concentration against time at different solvent compositions for the oxidation of Serine.

106

Table (3.60) Values of inverse of dielectric constant of the solvent and

observed rate constant for each amino acid.

Log kobs 1/D Ala. Asp. Cys. Glu. Lys. Phe. Ser.

0.0160 -2.52 -2.43 -2.26 -2.49 -2.26 -1.84 -2.06

0.0167 -2.60 -2.52 -2.28 -2.49 -2.36 -1.84 -2.06

0.0175 -2.56 -2.64 -2.28 -2.60 -2.56 -1.95 -2.12

y = -21.761x - 2.195

-2.7

-2.6

-2.5

-2.4

-2.30.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

y = -134.18x - 0.2838

-2.7

-2.6

-2.5

-2.4

-2.3

-2.20.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

Fig (3.159) Effect of solvent composition on the oxidation of Asparagine.

Fig (3.158) Effect of solvent composition on the oxidation of Alanine.

107

y = -11.953x - 2.0696

-2.3

-2.25

-2.20.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

y = -69.731x - 1.3583

-2.7

-2.6

-2.5

-2.4

-2.3

-2.20.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

Fig (3.161) Effect of solvent composition on the oxidation of Glutamic acid.

Fig (3.160) Effect of solvent composition on the oxidation of Cysteine.

y = -198.49x + 0.9336

-2.6

-2.4

-2.2

-20.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

y = -72.666x - 0.6574

-2

-1.9

-1.8

-1.7

-1.6

-1.50.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

kobs

Fig (3.163) Effect of solvent composition on the oxidation of Phenylalanine.

Fig (3.162) Effect of solvent composition on the oxidation of Lysine.

y = -40.792x - 1.395

-2.16

-2.12

-2.08

-2.04

-20.0155 0.016 0.0165 0.017 0.0175 0.018

1/D

Log

k obs

Fig (3.164) Effect of solvent composition on the oxidation of Serine.

108

3.2.5 Effect of Catalysts

The effect of catalyst was studied by the addition of diluted concentrations of

Ag+, Cu2+, and Both ions to the three different reaction mixtures and the

fourth reaction mixture was free of catalyst, and the following conditions

were applied:

Temperature = 60 0C, [K2S2O8] = 0.005M, [Amino acid] = 0.003M,

µ = 0.25M.

The rate of reaction and their average values were calculated, for more

details see appendix (U.1) to appendix (U.4).

Table (3.61) Values of each catalyst initial concentration and the

corresponding values of average rate.

106 x Rate m.l-1min-1 [Catalyst]o M

Ala. Asp. Cys. Glu. Lys. Phe. Ser.

0.00 3.35 3.95 9.55 13.55 10.15 8.65 4.7

Ag+ 0.005 32.7 39.4 25.95 42.3 25.0 41.95 51.6

Cu2+ 0.005 17.3 14.4 21.15 24.05 19.25 35.5 45.2

Ag+ 0.0025 + Cu2+ 0.0025 35.6 41.35 37.5 61.55 32.7 51.6 90.3

109

3.2.6 Effect of ionic strength

The ionic strength effect was studied by the addition of potassium sulfate

0.025M to the reaction mixtures, and the following conditions were applied:

Temperature = 60 0C, [K2S2O8] = 0.005M, [Amino acid] = 0.003M.

The observed rate constant of the reactions were calculated from the graph of

logarithm of peroxodisulphate concentration against time, fig (3.165) to fig

(3.171) for more details see appendix (V.1) to appendix (V.4).

The logarithm of observed rate constant was plotted against square root of

ionic strength and presented in fig (3.172) to fig (3.178).

◊-Flask (1) kobs = 0.0006909■-Flask (2) kobs = 0.002303∆-Flask (3) kobs = 0.0020727

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0009212■-Flask (2) kobs = 0.0020727∆-Flask (3) kobs = 0.0027636

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.166) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Asparagine.

Fig (3.165) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Alanine.

110

◊-Flask (1) kobs = 0.0025333■-Flask (2) kobs = 0.0039151∆-Flask (3) kobs = 0.0071393

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0039151■-Flask (2) kobs = 0.0039151∆-Flask (3) kobs = 0.004606

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.168) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Glutamic acid.

Fig (3.167) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Cystine.

◊-Flask (1) kobs = 0.0025333■-Flask (2) kobs = 0.0027636∆-Flask (3) kobs = 0.0029939

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

◊-Flask (1) kobs = 0.0020727■-Flask (2) kobs = 0.0013818∆-Flask (3) kobs = 0.0022101

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2 S2 O

8 ]

Fig (3.170) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Phenylalanine.

Fig (3.169) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Lysine.

◊-Flask (1) kobs = 0.0009212■-Flask (2) kobs = 0.0020727∆-Flask (3) kobs = 0.0011515

0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0 50 100 150 200 250

Time min.

3+Lo

g[K

2S2O

8]

Fig (3.171) Plot of logarithm of peroxodisulphate concentration against time at different values of ionic strength for the oxidation of Serine.

111

Table (3.62) Values of square root of ionic strength and logarithm of

observed rate constant.

Log kobs (µ)1/2

Ala. Asp. Cys. Glu. Lys. Phe. Ser.

0.500 -3.16 -3.04 -2.60 -2.41 -2.60 -2.68 -3.06

0.5477 -2.64 -2.68 -2.41 -2.41 -2.56 -2.86 -2.68

0.5916 -2.68 -2.56 -2.15 -2.34 -2.52 -2.66 -2.94

y = 5.2919x - 5.719

-3.5

-3

-2.5

-2

-1.5

-1

-0.5

00.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

y = 5.2399x - 5.6225

-3.5

-3

-2.5

-2

-1.5

-1

-0.5

00.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

Fig (3.173) Effect of Ionic Strength on the oxidation of Asparagine.

Fig (3.172) Effect of Ionic Strength on the oxidation of Alanine

y = 4.8981x - 5.0598

-3

-2.5

-2

-1.5

-1

-0.5

00.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

y = 0.7593x - 2.7986

-2.5

-2.4

-2.3

-2.2

-2.1

-20.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

Fig (3.175) Effect of Ionic Strength on the oxidation of Glutamic acid.

Fig (3.174) Effect of Ionic Strength on the oxidation of Cysteine.

112

y = 0.792x - 2.9923

-2.7

-2.6

-2.5

-2.4

-2.3

-2.2

-2.1

-20.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

y = 0.2462x - 2.8674

-3.5

-3

-2.5

-2

-1.5

-1

-0.5

00.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

Fig (3.177) Effect of Ionic Strength on the oxidation of Phenylalanine.

Fig (3.176) Effect of Ionic Strength on the oxidation of Lysine.

y = 1.1498x - 3.5142

-3.5

-3

-2.5

-2

-1.5

-1

-0.5

00.45 0.5 0.55 0.6

(Ionic Strength)1/2

Log

kobs

Fig (3.178) Effect of Ionic Strength on the oxidation of Serine.

113

3.3 Identification of reaction products

The products analyses were carried out under kinetic conditions. A mixture of

5.0ml of each amino acid, 10ml of 0.1M K2S2O8 solution and completed the

total volume to 50ml with deionized water was kept stoppered in a thermostatic

water bath at 600C till completion of the reaction (approximately five hours)

and the following tests were performed:-

(1) Test for Ammonia

5.0ml of the final solution was diluted to 50ml and a few drops of Nessler's

reagent were added; brownish color was observed in all amino acids mixtures

indicating that ammonia liberated in reaction mixtures.

(2) Test for Carbon dioxide

10ml of each final solution was re-evaporated and the gaseous product formed

was passed through a freshly prepared lime water solution; all the solutions

studied were turned milky, this confirmed the liberation of carbon dioxide.

(3) Test for Aldehydes

To confirm the presence of aldehydes, the reaction mixtures were treated with

solution of 2,4-Dinitrophenylhydrazine in methanol for products of aromatic

amino acids, and with solution of 2,4-Dinitrophenylhydrazine in 2.0M HCl for

products of aliphatic amino acids. The hydrazone derivatives were recrystaly-

zed twice from ethanol. Furthermore the recorded infrared spectra of these

hydrazones were obtained and showed that the absorption of characteristic

functional group C=N of hydrazone derivatives were observed in all cases at

certain wave length region (1610-1620 cm-1), as shown in the following spectra.

114

450600750900105012001350150016501800195021002400270030003300360039001/cm

45

50

55

60

65

70

75

80

85

90

%T

3444

.63

3292

.26

3112

.89

2358

.78

1618

.17

1515

.94

1417

.58

1303

.79

1263

.29

1218

.93

1135

.99

1072

.35

921.

91 887.

19

829.

33

742.

5472

1.33

599.

82

1 Fig (3.179) IR spectrum of the hydrazone derivative from reaction mixture of the Alanine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

35

40

45

50

55

60

65

70

75

80

%T

3296

.12

3091

.68

2362

.64

1708

.81

1618

.17

1593

.09

1517

.87

1423

.37

1278

.72

1220

.86

1135

.99

1074

.28

831.

26

742.

54

607.

54

2 Fig (3.180) IR spectrum of the hydrazone derivative from reaction mixture of Argnine and peroxodisulphate.

115

450600750900105012001350150016501800195021002400270030003300360039001/cm

30

35

40

45

50

55

60

65

70

75

80

85

%T

3425

.34

3290

.33

3103

.25

2929

.67

1726

.17

1620

.09

1593

.09

1517

.87

1421

.44

1305

.72

1263

.29

1217

.00

1135

.99

1072

.35

921.

91 887.

19

829.

33

742.

5472

1.33

592.

11

511.

10

3 Fig (3.181) IR spectrum of the hydrazone derivative from reaction mixture of Asparagine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

22.5

30

37.5

45

52.5

60

67.5

75

82.5

90

%T

3288

.40

3087

.82

1618

.17

1591

.16

1539

.09

1517

.87

1421

.44

1305

.72

1263

.29

1215

.07

1143

.71

1072

.35

921.

91

838.

98 721.

33

607.

54

511.

10

4 Fig (3.182) IR spectrum of the hydrazone derivative from reaction mixture of Aspartic acid and peroxodisulphate.

116

450600750900105012001350150016501800195021002400270030003300360039001/cm

47.5

50

52.5

55

57.5

60

62.5

65

67.5

70

72.5

75

77.5

80

82.5

85%T

2341

.42

5 Fig (3.183) IR spectrum of the hydrazone derivative from reaction mixture of Cysteine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

70

72.5

75

77.5

80

82.5

85

87.5

90

92.5

95

97.5

100

102.5

105

107.5%T

3440

.77

3305

.76

3093

.61

2360

.71

1591

.16

1517

.87

1423

.37

1315

.36

1278

.72

1184

.21

1130

.21

1068

.49

923.

84

831.

26

752.

19

7 Fig (3.184) IR spectrum of the hydrazone derivative from reaction mixture of Glycine and peroxodisulphate.

117

450600750900105012001350150016501800195021002400270030003300360039001/cm

60

62.5

65

67.5

70

72.5

75

77.5

80

82.5

85

87.5

90

92.5

95

97.5%T

3292

.26

3112

.89

2964

.39

2360

.71

1712

.67

1618

.17

1591

.16

1512

.09

1415

.65

1305

.72

1263

.29

1218

.93

1137

.92 10

70.4

2

919.

98

831.

26

744.

4772

3.26

613.

32

9 Fig (3.185) IR spectrum of the hydrazone derivative from reaction mixture of Leucine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

40

45

50

55

60

65

70

75

80

85

90

95

%T

2362

.64

1614

.31

1500

.52

1332

.72

10 Fig (3.186) IR spectrum of the hydrazone derivative from reaction mixture of Lysine and peroxodisulphate.

118

450600750900105012001350150016501800195021002400270030003300360039001/cm

22.5

30

37.5

45

52.5

60

67.5

75

82.5

90

97.5

%T

3446

.56

3288

.40

3103

.25

2360

.71

1618

.17

1591

.16

1500

.52

1421

.44

1305

.72

1263

.29

1215

.07

1143

.71

1072

.35

921.

91

838.

98

744.

47

605.

61

511.

10

14 Fig (3.187) IR spectrum of the hydrazone derivative from reaction mixture of Serine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

65

70

75

80

85

90

95

100

105

110

%T

3434

.98

3296

.12

3091

.68

1612

.38

1591

.16

1502

.44

1427

.23

1313

.43

1218

.93

1141

.78

1087

.78

1058

.85

833.

19

742.

54

15 Fig (3.188) IR spectrum of the hydrazone derivative from reaction mixture of Threonine and peroxodisulphate.

119

450600750900105012001350150016501800195021002400270030003300360039001/cm

35

40

45

50

55

60

65

70

75

%T

2908

.45

1614

.31

17 Fig (3.189) IR spectrum of the hydrazone derivative from reaction mixture of Tyrosine and peroxodisulphate.

450600750900105012001350150016501800195021002400270030003300360039001/cm

17.5

20

22.5

25

27.5

30

32.5

35

37.5

40

42.5

45

%T

2902

.67 23

60.7

1

1614

.31

1429

.15

1340

.43

18 Fig (3.190) IR spectrum of the hydrazone derivative from reaction mixture of Valine and peroxodisulphate.

120

4. DISCUSSION The progress of the reaction was followed by examining the concentration of

peroxodisulphate in the reaction mixture at different time intervals

iodometrically. The rate equation for the reaction peroxodisulphate-amino acid

can be expressed as:

S2O8ddt kobs

m n[AA]

[ ]2-

2-[ ]S2O8

-=

At fixed concentration of amino acid (0.003M) and peroxodisulphate

concentrations were ranged from 2.5x10-3 to 12.5x10-3 m.l-1. the plot of Log

–d[S2O82-]/dt against log [S2O8

2-]0 was linear and the slope of this line give the

value of m (the order with respect to peroxodisulphate) which is almost equal to

one with each amino acid studied. On the other hand, at fixed concentration of

peroxodisulphate (0.005M) and amino acid concentration were ranged from

1.0x10-3 to 5.0x10-3 the plot of Log –d[S2O82-]/dt against log [amino acid]0 was

linear and the slope of this line give the value of n (the order with respect to

amino acid), values of n almost equal to zero.

The rate equation can be written as the following form:

S2O8ddt kobs

n[AA]

[ ]2-

2-[ ]S2O8

-=

n = zero for each amino acid studied.

The values of kobs at fixed concentration of amino acid, H+ concentration, ionic

strength, and temperature, were found to vary linearly with initial

peroxodisulphate concentration with lines parallel to the concentration axis

except in case of alanine fig (3.9) and metheonine Fig (3.75).

Furthermore the values of kobs at fixed concentration of peroxodisulphate, H+

concentration, ionic strength, and temperature were found to vary linearly with

initial concentration of amino acid with lines parallel to the concentration axis

121

except in the case of Argnine Fig (3.18), leucine Fig (3.65) and Threonine Fig

(3.101).

Increasing in kobs values with increase in initial concentration of amino acid or

initial concentration of peroxodisulphate is explained by the fact that the

resultant product (aldehyde) in the oxidation of amino acid by peroxodisulphate

forms an additional compound with the parent amino acid in the reaction

mixture. This new compound is Schiff base, and increasing kobs values with

increasing peroxodisulphate concentration at fixed concentration of amino acid

is attributed to autocatalytic effect of this Schiff base , this effect formed from

the initial power of peroxodisulphate concentration, in which a considerable

amount of amino acid is converted to aldehyde at the beginning of the reaction

and this resultant product (aldehyde) will combined with the parent amino acid

to form Schiff base and the later is highly oxidizable than the parent amino acid.

Increasing kobs values with increasing concentration of amino acid at fixed

concentration of peroxodisulphate is attributed to the autocatalytic effect, in this

case it formed from the initial power of amino acid concentration which is give

a higher concentration of resultant product at the beginning of the reaction and

this is resulting in formation of Schiff base.

In general the autocatalytic effect of the resulting Schiff base is due to the fact

that Schiff base is highly oxidizable than the parent amino acids.

The autocatalytic effect is absence in some of amino acids may be due to the

following two assumptions:

(1) The resulting Schiff base may be formed in all cases between the formed

aldehyde and parent amino acid, the autocatalytic effect is absence in this group

of amino acids because the rate of oxidation of the formed Schiff base by

peroxodisulphate is quite near to the rate of oxidation of amino acid itself.

(2) The resulting Schiff base may not be possible to form due to steric

hindrance between the resulting aldehyde and parent amino acid.

122

There are a large number of references of the reaction of aldehyde with amino

acid to form a Schiff base(46) in fact the catalytic effect of pyridoxal phosphate

in the amino acid metabolism by enzymes is attributed to the Schiff base

formed between the amino acid and pyridoxal phosphate(62, 63, 64)

The reaction amino acid-peroxodisulphate studied over temperatures range

(60-800C), for each amino acid studied. It quite clear from the logarithm of

peroxodisulphate concentration against time in minutes, fig (3.115) to fig

(3.132), in all cases studied the observed rate constants increased with increase

in temperature; these results were tabulated in table (3.42) to table (3.59). By

applying arrhinius equation (eq 2.5), the natural logarithm of observed rate

constant were plotted against 1/T in K-1, and straight lines were found, fig

(3.133) to fig (3.150). From arrhinius equation

Lnk = LnA – E/ŘT

The slope of the graph is equal to – E/Ř, from these graphs the activation

energy for oxidation of each amino acid by peroxodisulphate was calculated

and tabulated below.

Fig (3.133) to Fig (3.150), straight lines cut the Ln kobs axis in values different

to that in the equation of the straight line attached, the actual values of Ln A is

that when the straight line cut the Ln k axis at 1/T is equal to zero, for example

as the case of alanine,

(A)

y = -16.251x + 42.681

-7

-6

-5

-4

-3

-2

-1

02.8 2.9 3 3.1

1/Tx103 K-1

LnK

obs

(B)

y = -16.251x + 42.681

-7

3

13

23

33

43

0 1 2 3 41/Tx103 K-1

Lnk o

bs

Fig (4.1) an arrhinius plot of Ln kobs against 1/T K-1 for the oxidation of alanine (A) extrapolation to Ln kobs axis at 1/T equal to 2.8, (B) extrapolation to Ln kobs axis at 1/T is equal to zero ( Actual value of Ln A).

123

Table (4.1) the thermodynamic parameters for the oxidation of amino acids by

peroxodisulphate.

Amino acid Ln A Ea kJ.mol-1 ∆S

J.K-1 at 333K

∆H kJ.mol-1 at 333K

∆G kJ. mol-1at 333K

Alanine 42.681 135.11 74.97 135.11 110.14 Arginine 40.282 128.25 55.02 128.25 109.93 Asparagine 26.644 91.35 -58.36 91.346 110.78 Aspartic acid 41.91 134.97 68.56 134.97 112.14 Cysteine 45.416 142.13 97.709 142.13 109.59 Glutamine 23.016 81.115 -88.52 81.111 110.59 Glutamic acid 38.95 126.56 43.95 126.56 111.92 Glycine 62.62 191.7 240.74 191.7 111.53 Histidine 48.577 154.57 123.97 154.58 113.29 Leucine 70.724 215.74 308.12 215.74 113.14 Lysine 44.453 138.18 89.7 138.18 108.31 Methionine 21.356 75.61 -102.3 75.605 109.68 Phenylalanine 65.462 197.82 264.37 197.82 109.78 Proline 46.174 145.79 104.01 145.79 111.15 Serine 39.591 129.8 49.23 129.82 113.41 Threonine 54.603 169.53 174.09 169.53 111.56 Tyrosine 60.82 188.16 225.77 188.16 112.98 Valine 56.054 174.37 186.15 174.37 112.38

It is clear from the above table that the free energy change ∆G# for all cases

under study is approximately the same, which is found to vary between 108.31

KJ.Mol-1 as the minimum value to 113.41 KJ.Mol-1 as the maximum value.

These values of free energy change are the same as Chandraju’s values(37) for

the oxidation of Serine by Manganese (III) in three different media, acetic acid,

pyrophosphate, and sulphuric acid respectively.

This equality in the values of free energy change, to some extent, suggests that

the dependence of ∆G# on solvent or substituents is much smaller than that of

124

E and ∆H#. The values of ∆H# are approximately equal to the values of E. The

values of enthalpy change were calculated from the equation

∆G# = ∆H# - T∆S# While the values of E were calculated from the Arrhinius equation .

As the 20 essential amino acids were grouped in 7 groups, there will be

a relationship can be noticed in each individual group concerning the

differences in the values of activation energy with increasing in inductive effect

of the R group of the amino acid structure (for more details about the amino

acid structure see appendix (W)).

In the aliphatic amino acid group, Glycine, Alanine, Valine, Leucine, and

Proline the inductive effect will explain clearly the increase in activation

energy, and according to inductive effect this group can be ordered as increase

in activation energy as follow

Alanine <Proline <Valine <Leucine

In the hydroxylic amino acids group which include Serine and Threonine, the

differences in activation energy is attributed to inductive effect of the R group,

in which Serine < Threonine in activation energy.

There is a great difference in activation energy between Alanine and

Phenylalanine, and according to the inductive effect of the R group these two

amino acids can be ordered as Phenylalanine< Alanine in activation energy.

In the basic amino acids group, Lysine and Argnine, Lysine due to the inductive

effect possess a higher value of activation energy.

Generally if there is a way to perform structural relationship between two or

more amino acids, the greater in value of activation energy is the greater in

inductive effect of the R group of amino acid structure.

The reaction is highly catalyzed by catalysts such as Ag+ and Cu2+. Solutions of

low concentration of Ag+, Cu2+ and mixture of Ag+ and Cu2+ were used and the

order of the reaction with respect to peroxodisulphate is always remaining to be

125

one. Table (3.61) gives values of initial concentration of catalysts and their

corresponding values of average rate of reaction, the rate of oxidation of amino

acids by peroxodisulphate is much faster in the presence of mixture of (Ag+ and

Cu2+) compared to the rate of oxidation in the presence of only Ag+ or Cu2+ ions

under the same conditions.

The rate of oxidation of amino acids by peroxodisulphate can be ordered as the

presence of the above catalysts at the same concentration as follow, rate of

uncatalyzed reaction < rate in presence of 0.005M Cu2+< rate in presence of

0.005M Ag+ < rate in presence of 0.0025M of Ag+ and 0.0025M of Cu2+.

The enhanced catalysis by Cu2+ observed in the oxidation of amino acids by

peroxodisulphate in the presence of Ag+ could be attributed to the facile

oxidation of radicals by Cu2+ (produced by the reaction of amino acid with

radicals of Ag2+ and sulphate ion free radicals) and the rapid reoxidation of Cu+

produced by peroxodisulphate ion could be the reason for enhanced rate of

decomposition of peroxodisulphate.

Acrylonitrile (radical trabing agent) was used to detect the free radical in the

reaction mixture, and give a milky appearance with the reaction mixture in dark

conditions indicating an in situ formation of free radical, therefore the oxidation

reaction of amino acids by peroxodisulphate is free radical in mechanism, and

the reaction of acrylonitrile with sulfate ion free radical can be written as

SO4 CH2 CH

CN

SO4 CH2 CH CH2 CH CH2 CH SO4

CN CN nCN

.+

According to BrÖnsted and Bjerrum theory for activated complex applied to the

charged particles, the reaction between peroxodisulphate and amino acid will

considered to proceed through an activated complex.

The complex is considered to be in equilibrium with reactants, and the

equilibrium constant is k‡ in activities is expressed as

126

k‡ = a‡ /aA. aS = (C‡ /CA. CS) .(γ‡ / γA. γS)

Where a‡ is the activity of the activated complex, aA is the activity of amino

acid in acidic form, aS is the activity of peroxodisulphate ion, γ‡ is the activity

coefficient of the activated complex, γA is the activity coefficient of amino acid

in acidic form, γS is the activity coefficient of peroxodisulphate ion, C‡ is the

concentration of the activated complex, CA is the concentration of amino acid

in acidic form, CS is the concentration of peroxodisulphate ion, and k‡ is the

equilibrium constant. The concentration of the activated complex is

C‡ = k‡ CA. CS (γA. γS) / γ‡

The reaction rate is

(-dCS / dt) = kobs . CA. CS

The reaction rate in activated complex concentration term is

(-dCS / dt) = (KT/h) C‡

K is Boltzman constant, and h is Planck constant.

kobs = (KT/h) C‡ / CA. CS

Introduce value of C‡from above equation

kobs = (KT/h) k‡ .(γA. γS / γ‡)

Taking the logarithm

Log kobs = Log (KT/h) + Log k‡ +Log γA + Log γS - Log γ‡

In dilute aqueous solution the activity coefficient term can be estimated from

the Debye-Hϋckel theory

Log γi = -0.509 Zi2 √ µ

Where µ is the ionic strength of the aqueous solution.

Substituting the Debye-Hϋckel expression in the above equation

Log kobs = Log (kTk‡ / h) + [-0.509 ZA2 -0.509 ZS

2-0.509 (ZA + ZS) 2 √ µ ]

And then,

127

Log kobs = Log (kTk‡ / h) + [1.018 ZA ZS] √ µ ]

This equation predicts that the plot of Log kobs against the square root of ionic

strength expressed straight lines relationship with all cases studied, and fig

(3.172) to fig (3.178) express a primary kinetic salt effect because the slope of

each case is positive value.

The rate decreased with increase in dielectric constant (D) of the medium using

variety of acetic acid- water percentages fig (3.158) to fig (3.164).

The relative dielectric constant of the medium acetic acid- water system can be

calculated using the following equation(65)

D = (C1D1 + C2D2) / 100

C1 and C2 are the percentages of water and acetic acid in the mixture, D1 is the

dielectric constant of pure water (approximately 65.2 at temperature of the

reaction), D2 is the dielectric constant of pure acetic acid (approximately 10.9 at

temperature of the reaction).

The plot of logarithm of kobs versus 1/D is almost linear in all cases studied with

a negative slope, the effect of dielectric constant on the rate for a reaction

involving two ions was given by the standard relationship, as the following

equation(23)

Log kobs = Log ko – ZAZSe2 / DKTdAS

Where ko is the rate constant in a medium of infinite dielectric constant, ZAe ,

ZSe are the charges of the two ions, dAS is the activated complex size, K is the

Boltzman constant, and T is absolute temperature. The slope is equal to

-ZAZSe2 / KTdAS. The values of dAS, the activated complex size, were computed

as 43.0 x102, 6.9 x 102, 77.8 x102, 13.3 x102, 4.7 x102, 12.7 x102, and 22.7 x102

Å for Alanine, Asparagine, Cysteine, Glutamic acid, Lysine, Phenylalanine,

and Serine respectively, the electron charge was taken as (4.8029 x 10-10 esu). Amis(66) shown that in a straight line plot of logarithm kobs versus 1/D a positive

slope indicates a positive ion-dipole reaction, while a negative slope indicates

128

the involvement of two dipoles or a negative ion-dipole reaction, in this

investigation a plot of logarithm kobs versus 1/D, fig (3.158) to fig (3.164) give

straight lines with negative slopes, these clearly supporting the involvement of

two dipoles, a negative ion-dipole, or a negative ion dissociation in the rate

determining step.

The stoichiometry of peroxodisulphate amino acid reaction was studied

extensively by several workers(57.58.59.60.61) and found that the reaction follow a

1:1 mole ratio, in this investigation a modification was introduced in

determination of amino acids by ninhydrin and chromotropic acid using

spectrophotometric methods, according to this modification the reaction

glycine-peroxodisulphate reaction follow a 1:1 mole ratio, this modification can

be used in stoichiometric determination of any amino acid peroxodisulphate

reaction, but the most important condition is availability of the expected

aldehyde formed from the oxidation reaction.

Stoichiometric study of the reaction peroxodisulphate amino acid by keeping

excess of peroxodisulphate concentration over amino acid concentration and

determining the residual peroxodisulphate iodometrically were performed for

all amino acids under study and no significant results were obtained, probably

because peroxodisulphate will oxidizes the product aldehyde after or within

oxidation of parent amino acid and a little bit concentration of peroxodisulphate

is consumed in oxidation of water.

The products analyses were carried out under the kinetic conditions, all

mixtures of amino acids and peroxodisulphate were tested with Nessler’s

reagent, a brownish color was observed in all amino acid mixtures, indicating

the liberation of ammonia, and therefore the deamination reaction was occurred.

All reaction mixtures after re-evaporation of small amount of solution and

passes the gaseous products into freshly prepared solution of lime water were

turned milky, indicating the presence of carbon dioxide; therefore the

129

decarboxylation was also occurred. Treatment of the reaction mixtures with

2,4-dinitrophenylhydrazine result in precipitation of hydrazone derivatives,

after recrystallization the IR measurements were recorded for these derivatives,

and the characteristic functional group C=N was clearly observed at wave

length between (1610- 1620 cm-1) fig (3.179) to fig (3.190), this functional

group was produced from the condensation reaction of aldehyde and hydrazine.

The data in table (4.1) showed that the energy of activation is highest for the

slowest reaction, indicating that the reaction is enthalpy controlled, and the

highly negative values of ∆S# in the case of Asparagine, Glutamine, and

Methionine indicated the formation of more rigid transition state and this

transition state is extensively solvated than the reactants, while the relatively

higher values of ∆S# in the other amino acids indicated a moderately rigid

transition state. Values of ∆S# and ∆H# from the above table, were linearly

interrelated fig (4.2) with R2 = 0.9987,

y = 0.3375x + 110.77R2 = 0.9987

0

50

100

150

200

250

-200 -100 0 100 200 300 400

∆S J.K-1 at 333K

∆H

kJ.

mol

-1

Fig (4.2) plot of enthalpy of activation against entropy of activation for the

oxidation of amino acids by peroxodisulphate.

130

The iso-kinetic temperature computed from the plot of ∆H# versus ∆S# is

337.5K which is quite near to the actual experimental temperature 333K.

The linear correlation implies that all the amino acids under study are oxidized

by the same mechanism and the changes in the rate are governed by the changes

in enthalpy and entropy.

The following reaction scheme explain all the observed experimental results.

S2O82-

SO4.-

2k1 ............ ︵1 ︶

Step (1) is rate determining step.

The sulfate free radicals produced attacks the molecules of water resulting in

formation of hydroxyl free radicals in fast step,

-.SO4 + H2O

k2 HSO4- + OH

. ........... ︵2 ︶ The hydroxyl free radicals attacks the molecules of amino acid,

.OH R CH CO2H

NH3+

k3+

NH2

R CH CO2H.

H2O+ ....... ︵3 ︶

Bergel and Bolz(67) who studied the oxidative decarboxylation of amino acids

by oxidants, found that α-dimethylaminoisobutyric acid was smoothly oxidized

to acetone, dimethylamine, and carbon dioxide, and the reaction thus taking

place in a compound lacking in replaceable hydrogen atom on the α-carbon

atom, and the fact that trialkylbetaines(68) resist oxidation suggest that

availability of an ionizable proton on the amino group is a prerequisite for the

reaction. According to the above facts the hydroxyl radicals attack the amino

acid molecule on the amino group resulting in formation of amino acid radical

which attacks the peroxodisulphate molecule to form sulfate ion and sulfate ion

free radical,

131

2-S2O8

.R CH CO2H

NH2

+ k4

NH2

R CH CO2H+

-.SO4+ + SO4

2- ....... ︵4 ︶

Step (5) resulting in formation of aldehyde,

H2O+ +++

R CH CO2HNH2

k5 R C H NH4 CO2 H+O

+ + ........ ︵5 ︶

The resulting aldehyde with the parent amino acid forming Schiff base,

OR C H +

+NH3

R CH CO2H k6 R C N CH CO2HH

+

RH

.......... ︵6 ︶

Hydroxyl free radical undergoes the oxidation by attacking the Schiff base

resulting in step (6) to form Shiff base free radical as fast as in step (3),

.+OH.

H R

+

HR C N CH CO2H k7 R C N CH CO2H+

RH

+ H2O ......... ︵7 ︶

Peroxodisulphate molecule in step (8) is a subject to attack by Schiff base free

radical to form sulfate ion and sulfate free radical,

H R

+R C N CH CO2H.+S2O82-

R C N CH CO2H+

RH

k8 ....... ︵8 ︶2-SO4++ SO4

.-

Two molecules of aldehyde were formed in step (9),

H R

+R C N CH CO2H+H2O2 k9 ........ ︵9 ︶++O

H+CO2NH4R C H + +2

The reaction was terminated by step (10) and (11) in which all remaining amino

acid free radicals were consumed by molecules of Shiff base,

132

.R CH CO2H

NH2

+

H R+

HR C N CH CO2H k10 R C N CH CO2H

+RH

+

NH3

R CH CO2H. ........ ︵10 ︶

And in step (11) the sulfate ion free radicals were consumed by the Sciff base

free radicals.

.

H R+

R C N CH CO2H-.

SO4 + k11 R C N CH CO2H+

RH

SO42- ....... ︵11 ︶+

Applying the steady state hypotheses for the above scheme lead to, [AA] stand

for amino acid concentration, and [Sch] stand for Schiff base concentration,

=-d - --2-

[S2O8]k1[S2O8]2-

/dt k4.

[AA]2-

[S2O8] k8.

[SCh][S2O8]2- ........ ︵1 ︶

And the rate of all the free radicals will be constant at the end of the reaction,

--/dt-d =.

[OH]-.

[SO4][H2O]k2 k3.

[OH][AA] k7.

[OH][SCh] ........ ︵2 ︶

--=-d /dt

-.[SO4] 2k1[S2O8]

2-k2

-.[SO4][H2O]+k4

.[AA]

2-[S2O8]

2-[S2O8][SCh]

.k8+ k11

.[SCh]

-.[SO4] ........ ︵3 ︶

--/dt-d.

[AA] [AA][OH].

k3= [S2O8]2-

[AA].

k4 k10.

[AA][SCh] ........ ︵4 ︶

- -=-d /dt.

[SCh] [SCh][OH].

k7 k8[SCh][S2O8]2-.

[SCh][AA].

k10+ [SO4].-

[SCh].

k11 ........ ︵5 ︶ From equation (2),

........ ︵6 ︶k3[AA] ︶+ [SCh]k7︵/-.

[SO4][H2O]k2=[OH].

︶︵ From equation (3),

........ ︵7 ︶︶+ [SCh].

k11/ ︵k8[SCh][S2O8]2-.

+k4.

[AA]2-

[S2O8]+[S2O8]2-

2k1=[SO4].-

[H2O]k2︶︵ Substituting equation (6) into equation (4),

-.[SO4]= ︵k3k4[AA]

.[AA]

2-[S2O8]+k4k7[SCh]

2-[S2O8]

.[AA]+k3k10 [AA]

.[AA] [SCh]

+k7 k10.

[AA][SCh]2

︶/ k3k2 [AA][H2O] ........ ︵8 ︶

133

Substituting equation (6) into equation (5),

k3︵=[SO4].-

k8[AA].

[SCh]2-

[S2O8]+ k8k7.

[SCh][SCh]2-

[S2O8]-k3k10 [AA].

[AA] [SCh]

︶

2[SCh][AA]

.k10k7- / ︵k2k7[SCh][H2O]-k3k11[AA]

.[SCh]-k7k11

.[SCh][SCh] ︶........ ︵9 ︶

Subtracting equation (3) from equation (5),

+-.

[SO4]= ︵2k1k3[AA]2-

[S2O8] k1k7[SCh]2-

[S2O8]+k3k4 [AA].

[AA]2-

[S2O8]

+k4k7.

[AA][SCh]2-

[S2O8]+2k3k8[AA].

[SCh]2-

[S2O8]+ k8k72 [SCh].

[SCh]2-

[S2O8] ︶/

︵k2k3[AA][H2O]+2k2k7[H2O][SCh] ︶........ ︵10 ︶

Taking equation (8) equal to equation (10),

+-.

[SO4]= ︵2k1k3[AA]2-

[S2O8] k1 k7[SCh]2-

[S2O8] + k3 [AA]2-

[S2O8]

+

222 k3 [AA] 2

22 k8.

[SCh]

2k3k7k8.

[SCh][SCh]2-

[S2O8][AA]-k3 k102 2

[AA].

[AA][SCh]- k3k7k10[AA].

[AA][SCh]2

-2k3k4k7[SCh].

[AA][AA]2-

[S2O8]-2k4k72 [SCh]

2 2-[S2O8]

.[AA]

3

- k7k1022 [SCh]3

.[AA]= Zero ........ ︵11 ︶

When equation (7) is equal to equation (9),

........ ︵12 ︶[H2O][SCh].

[AA]2-

[S2O8]k7k4k2-

[SCh].

[AA][S2O8]2-

[AA].

k11k4k3+[SCh][SCh].

[S2O8]2-

[AA].

k11k7k4+[H2O][SCh][S2O8]2-

k7k2k12-

[SCh].

[AA][S2O8]2-

k11k3k12+[SCh].

[SCh][S2O8]2-

k11k7k12+[SCh]2[SCh].

[AA].

k11k10k7-

[SCh][SCh].

[AA][AA].

k11k10k3-[SCh].

[SCh]2 2-[S2O8]k11k8k72+2 [S2O8]

2-[AA]2[SCh]

.k11k8k3+

2[SCh][AA].

[H2O]k10k7k2-[H2O][SCh][AA].

[AA]k10k3k2-[S2O8]2-

[SCh].

[AA][H2O]k8k3k2

= Zero Taking equation (7) equal to equation (11),

+ k3k8k11.

[SCh]2 [AA]2-

[S2O8]2 +2 k7k8k11 [S2O8]2-2

[SCh].

[SCh]

+2k1k7k112-

[S2O8] [SCh].

[SCh]+2k1k3k112-

[S2O8][AA].

[SCh]

-2k1k2k72-

[S2O8][SCh][H2O]

+k4k7k11.

[AA]2-

[S2O8].

[SCh][SCh]+k3k4k11.

[AA]2-

[S2O8][AA].

[SCh]

-k2k4k7 [S2O8]2-

[AA].

[SCh][H2O]

[S2O8]2-

[SCh].

[AA][H2O]k8k3k2

Zero= ........ ︵13 ︶ Subtracting equation (12) from equation (13),

134

k2k7k10 [H2O].

[AA][SCh]2 k3k10k11.

[AA][AA].

[SCh][SCh]

k7k10k11.

[AA].

[SCh] 2[SCh] ........ ︵14 ︶

[H2O][SCh][AA][AA]k10k3k2.

+ +

+ Zero= Rearrangement equal (14) lead to,

-.

[SCh]= k2 [H2O] / k11 ........ ︵15 ︶ Substituting equation (15) into equation (11),

.[AA] = 2k3[AA]

2-[S2O8] ︵k1 k3k11[AA]+k1k7 k11[SCh]-k2k3k8[H2O][AA]-k2k7k8[H2O][AA] ︶/

︵k11[SCh] k32 k10 [AA]

2+3k3k7k10[AA][SCh]+ k32 k4k7[AA]

2-[S2O8]

+2k42k7 [SCh]

2-[S2O8]+2k7

2 k10[SCh]2

︶........ ︵16 ︶ Substituting equations (15) and (16) into equation (1)

[S2O8]

2-[S2O8]

-

k4

k8

2k3[S2O8]2-

A k11 ︶

-k2 [H2O] k11

/ [SCh]B

2-

[S2O8]2-

[AA]2-

=/dt-d[S2O8] k1

︵ / ︶

- ︵

[S2O8]-d dt/ = -k1 k4k8 [SCh]︵ 2k3 [S2O8]

2-AB k11 ︶[SCh] -k2 [H2O] k11B / [SCh]B

2-[S2O8]

2-[AA]

And this is give the overall rate equation,

[S2O8]-d dt/ = [S2O8]2-

[AA]kobsn

2-

And the values of the constants kobs, A, and B can be obtained from the following equations,

[AA] B[SCh]/B k11[H2O]k2 -[SCh] ︶k11B A2-

[S2O8]k32[SCh]k8 k4k1-=kobs

k3 [AA] k1 [AA]k7 k11[SCh] k2k7k8 [H2O] ︶k3+ - k82k3A k11 [SCh]-k2 [H2O][AA]2-

= [S2O8] k1︵ and

k42-

[S2O8]

k3

+

k11

k10

[AA]2+3k3k7k10[AA][SCh] + k32 k4k7[AA]

2-[S2O8]

+2 2k7 [SCh] 2k72

k10

[SCh]2

︶

= [SCh] ︵B 2

The rate law is consistent with experimental results, first order in [S2O4

2-], kobs

is highly affected by concentration of amino acid, peroxodisulphate, and Schiff

base, in some cases of amino acids the effect of amino acid concentration is

similar to effect of peroxodisulphate.

135

APPENDICES

(1) Alanine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and alanine on reaction rate. APPENDIX (A.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M, and [Alanine]0 = 0.003 M, For effect of Alanine, [Alanine]0 = flask (6) =0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005M. [K2S2O8]0=0.005M, µ=0.25M by potassium sulfate, pH =2.00,Temp.= 333K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Alanine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.371 0.663 0.82 0.922 1.011 0.544 0.556 0.568 0.574 0.562 30 0.342 0.639 0.80 0.881 0.982 0.512 0.525 0.544 0.544 0.531 60 0.348 0.628 0.80 0.873 0.963 0.477 0.498 0.519 0.519 0.505 90 0.332 0.608 0.778 0.845 0.932 0.439 0.462 0.491 0.477 0.462 120 0.336 0.604 0.784 0.843 0.927 0.398 0.415 0.447 0.455 0.439 150 0.33 0.592 0.778 0.828 0.909 0.371 0.389 0.431 0.423 0.398 180 0.324 0.58 0.772 0.813 0.891 0.290 0.332 0.389 0.380 0.362 210 0.318 0.568 0.766 0.798 0.873 0.230 0.301 0.352 0.342 0.301 APPENDIX (A.2) Detailed results showing the average rate of reaction in each flask for peroxodisu-lphate alanine system according to the conditions applied in appendix (A.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Alanine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 3.24 8.33 8.93 20.05 21.67 8.33 8.33 6.67 8.33 8.33 30 1.25 2.30 3.2 8.35 11.7 8.33 6.67 6.67 6.67 6.67 60 2.18 3.35 3.2 8.35 11.7 8.33 8.33 6.67 10.00 10.00 90 0.33 3.35 3.2 8.35 11.7 8.33 10.00 10.00 5.00 5.00 120 1.25 3.35 3.2 8.35 11.7 5.000 5.00 3.33 6.67 8.33 150 1.25 3.35 3.2 8.35 11.7 13.33 10.00 8.33 8.33 6.67 180 1.25 3.35 3.2 8.35 11.7 8.33 5.00 6.67 6.67 10.00 210 / / / / / / / / / /

Aver. 1.53 3.91 4.02 10.02 13.12 8.57 7.62 6.90 7.38 7.857

136

(2) Argnine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and argnine on reaction rate. APPENDIX (B.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M [Argnine]0 = 0.003 M For effect of Argnine, [Argnine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. [K2S2O8]0 = 0.005 M, t µ = 0.25 M by potassium sulfate, pH = 1.94 , Temp. =333K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Argnine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.29 0.613 0.782 0.914 0.991 0.574 0.58 0.568 0.556 0.585 30 0.267 0.544 0.76 0.863 0.942 0.550 0.550 0.531 0.512 0.538 60 0.235 0.486 0.699 0.823 0.91 0.531 0.525 0.498 0.462 0.484 90 0.217 0.462 0.628 0.758 0.852 0.512 0.491 0.455 0.407 0.423 120 0.161 0.431 0.602 0.732 0.816 0.477 0.455 0.407 0.352 0.352 150 0.154 0.422 0.591 0.715 0.82 0.455 0.423 0.362 0.279 0.278 180 0.127 0.401 0.555 0.679 0.79 0.415 0.380 0.301 0.176 0.176 210 0.1 0.367 0.519 0.643 0.76 0.398 0.332 0.230 0.079 0.041 APPENDIX (B.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate- argnine system according to the conditions applied in appendix (B.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Argnine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 3.33 16.83 10.00 30.00 34.80 6.67 8.33 10.00 11.67 13.33 30 3.76 8.00 21.37 17.83 18.60 5.00 6.67 8.33 11.67 13.33 60 2.91 15.17 12.90 17.40 18.60 5.00 8.33 10.00 11.67 13.33 90 6.67 6.67 12.90 28.10 18.60 8.33 8.33 10.00 10.00 13.33 120 0.18 2.16 12.90 6.70 18.60 5.00 6.67 8.33 11.67 11.67 150 3.25 8.00 12.90 17.40 18.60 8.33 8.33 10.00 13.33 13.33 180 3.25 8.00 12.90 17.40 18.60 3.33 8.33 10.00 10.00 13.33 210 / / / / / / / / / /

Aver. 3.33 9.26 13.70 19.26 20.91 5.95 7.86 9.52 11.43 13.10

137

(3) Asparagine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and asparagine on reaction rate. APPENDIX (C.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M, [Asparagine]0 = 0.003 M, For effect of Asparagine, [Asparagine]0 = flask (6) = 0.001M, flask (7) = 0.002 M,flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. [K2S2O8]0 = 0.005 M, µ = 0.25 M potassium sulfate, pH = 1.91and Temp. = 333K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Asparagine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.290 0.608 0.775 0.898 1.029 0.505 0.512 0.505 0.505 0.512 30 0.279 0.585 0.748 0.886 1.011 0.484 0.491 0.477 0.477 0.491 60 0.230 0.574 0.740 0.826 0.996 0.4624 0.462 0.447 0.455 0.462 90 0.190 0.568 0.708 0.823 0.978 0.431 0.447 0.423 0.431 0.439 120 0.176 0.562 0.686 0.813 0.957 0.423 0.415 0.389 0.398 0.415 150 0.161 0.538 0.677 0.806 0.932 0.398 0.380 0.352 0.362 0.380 180 0.130 0.519 0.653 0.796 0.927 0.362 0.362 0.322 0.342 0.352 210 0.114 0.498 0.633 0.778 0.914 0.342 0.332 0.279 0.312 0.322 APPENDIX (C.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-asparagine system according to the conditions applied in appendix (C.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Asparagine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 1.67 6.67 11.67 6.67 15.00 5.00 5.00 6.67 6.67 5.00 30 6.67 3.33 3.33 33.33 11.67 5.00 6.67 6.67 5.00 6.67 60 5.00 1.67 13.33 1.67 13.33 6.67 3.33 5.00 5.00 5.00 90 1.67 1.67 8.33 5.00 15.00 1.67 6.67 6.67 6.67 5.00 120 1.67 6.67 3.33 3.33 16.67 5.00 6.67 6.67 6.67 6.667 150 3.33 5.00 8.33 5.00 3.33 6.67 3.33 5.00 3.33 5.00 180 1.67 5.00 6.67 8.33 8.33 3.33 5.00 6.67 5.00 5.00 210 / / / / / / / / / /

Aver. 3.10 4.29 7.86 9.05 11.90 4.76 5.24 6.19 5.48 5.48

138

(4) Aspartic acid-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and aspartic acid on the rate. APPENDIX (D.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Aspartic acid]0 = 0.003 M For effect of Aspartic acid, [Aspartic acid]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.84 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Aspartic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.389 0.591 0.803 0.898 0.985 0.512 0.498 0.519 0.512 0.512 30 0.290 0.585 0.789 0.884 0.966 0.498 0.477 0.498 0.491 0.491 60 0.279 0.58 0.752 0.878 0.959 0.477 0.462 0.477 0.470 0.470 90 0.279 0.574 0.740 0.854 0.937 0.47 0.447 0.455 0.439 0.447 120 0.267 0.58 0.720 0.851 0.919 0.462 0.431 0.431 0.423 0.415 150 0.290 0.568 0.732 0.845 0.900 0.447 0.407 0.415 0.398 0.389 180 0.267 0.556 0.724 0.842 0.889 0.431 0.389 0.398 0.380 0.362 210 0.243 0.544 0.712 0.833 0.881 0.407 0.371 0.362 0.342 0.342 APPENDIX (D.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-aspartic acid system according to the conditions applied in appendix (D.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Aspartic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 5.00 3.33 8.33 5.00 1.67 3.33 5.00 5.00 5.00 5.00 30 5.00 5.00 3.33 6.67 6.67 5.00 3.33 5.00 5.00 5.00 60 6.67 8.33 11.67 6.67 10.00 1.67 3.33 5.00 6.67 5.00 90 5.00 0.00 3.33 6.67 5.00 1.67 3.33 5.00 3.33 6.67 120 3.33 15.00 6.67 6.67 5.00 3.33 5.00 3.33 5.00 5.00 150 1.67 1.67 5.00 8.33 10.00 3.33 3.33 3.33 3.33 5.00 180 3.33 3.33 5.00 3.33 5.00 5.00 3.33 6.67 6.67 3.33 210 / / / / / / / / / /

Aver. 4.29 5.24 6.19 6.19 6.19 3.33 3.81 4.76 5.00 5.00

139

(5) Cysteine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and cysteine on reaction rate. APPENDIX (E.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Cysteine]0 = 0.003 M For effect of Cysteine, [Cysteine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M.at µ = 0.25 M by potassium sulfate, pH = 3.30 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Cysteine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.253 0.544 0.623 0.816 0.903 0.352 0.362 0.371 0.342 0.352 30 0.218 0.506 0.6 0.79 0.87 0.332 0.342 0.342 0.312 0.322 60 0.178 0.465 0.574 0.762 0.835 0.312 0.312 0.322 0.290 0.290 90 0.133 0.419 0.548 0.732 0.796 0.279 0.290 0.290 0.255 0.243 120 0.084 0.368 0.52 0.7 0.753 0.255 0.255 0.267 0.217 0.204 150 0.027 0.311 0.489 0.665 0.706 0.230 0.230 0.230 0.190 0.161 180 -0.037 0.244 0.456 0.627 0.653 0.204 0.190 0.204 0.146 0.114 210 -0.113 0.165 0.421 0.587 0.593 0.176 0.161 0.161 0.097 0.041

APPENDIX (E.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-Cysteine system according to the conditions applied in appendix (E.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Cysteine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 4.9 9.7 7.45 12.85 19.45 3.33 3.33 5.00 5.00 5.00 30 4.9 9.7 7.45 12.85 19.45 3.33 5.00 3.33 3.33 5.00 60 4.9 9.7 7.45 12.85 19.45 5.00 3.33 5.00 5.00 6.67 90 4.9 9.7 7.45 12.85 19.45 3.33 5.00 3.33 5.00 5.00 120 4.9 9.7 7.45 12.85 19.45 3.33 3.33 5.00 3.33 5.00 150 4.9 9.7 7.45 12.85 19.45 3.33 5.00 3.33 5.00 5.00 180 4.9 9.7 7.45 12.85 19.45 3.33 3.33 5.00 5.00 6.67 210 / / / / / / / / / /

Aver. 4.9 9.7 7.45 12.85 19.45 3.57 4.05 4.29 4.52 5.48

140

(6) Glutamic acid-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and glutamic acid on the rate. APPENDIX (F.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Glutamic acid]0 = 0.003 M For effect of glutamic acid, [Glutamic acid]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M (potassium sulfate), pH = 2.02 ,Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Glutamic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.380 0.672 0.842 1.021 1.097 0.544 0.544 0.544 0.562 0.544 30 0.368 0.648 0.796 0.920 1.070 0.498 0.512 0.519 0.531 0.519 60 0.356 0.607 0.763 0.900 1.021 0.484 0.477 0.484 0.502 0.484 90 0.344 0.512 0.677 0.892 0.978 0.447 0.439 0.455 0.462 0.45 5 120 0.332 0.487 0.628 0.875 0.975 0.407 0.389 0.415 0.423 0.423 150 0.319 0.425 0.554 0.867 0.929 0.377 0.352 0.371 0.371 0.389 180 0.305 0.380 0.466 0.854 0.906 0.322 0.301 0.332 0.322 0.342 210 0.291 0.267 0.354 0.801 0.869 0.267 0.230 0.290 0.267 0.301 APPENDIX (F.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-glutamic acid system according to the conditions applied in appendix (F.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Glutamic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 2.10 8.175 23.33 73.33 25.00 11.67 8.33 6.67 8.33 6.67 30 2.10 13.33 15.00 11.67 41.67 3.33 8.33 8.33 7.33 8.33 60 2.10 26.67 35.00 5.00 33.33 8.33 8.33 6.67 9.33 6.67 90 2.10 6.025 16.8 10.00 1.67 8.33 10.00 8.33 8.33 6.67 120 2.10 13.55 22.05 5.00 31.67 5.5 6.67 8.33 10.00 6.67 150 2.10 8.758 22.05 6.67 15.00 9.5 8.33 6.67 8.33 8.33 180 2.10 18.34 22.05 27.53 21.90 8.33 10 6.67 8.33 6.67 210 / / / / / / / / / /

Aver. 2.10 13.55 22.33 19.89 24.32 7.86 8.579 7.38 8.57 7.14

141

(7) Glutamine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and glutamine on reaction rate. APPENDIX (G.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Glutamine]0 = 0.003 M For effect of Glutamine, [Glutamine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 3.51 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Glutamine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.332 0.690 0.900 0.993 1.041 0.740 0.667 0.658 0.653 0.628 30 0.312 0.681 0.860 0.952 0.980 0.574 0.613 0.585 0.638 0.560 60 0.290 0.672 0.796 0.903 0.966 0.544 0.574 0.574 0.602 0.556 90 0.255 0.663 0.760 0.884 0.959 0.538 0.568 0.568 0.560 0.550 120 0.243 0.648 0.720 0.866 0.884 0.531 0.562 0.562 0.568 0.556 150 0.230 0.643 0.695 0.854 0.875 0.525 0.556 0.556 0.556 0.544 180 0.217 0.638 0.653 0.845 0.860 0.498 0.544 0.544 0.544 0.538 210 0.190 0.628 0.638 0.829 0.845 0.491 0.538 0.531 0.531 0.531 APPENDIX (G.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-glutamine system according to the conditions applied in appendix (G.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Glutamine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 3.33 3.33 23.33 30.00 48.33 58.33 18.33 23.33 5 15 30 3.33 3.33 33.33 31.67 10.00 8.33 11.67 3.33 11.67 6.67 60 5.00 3.33 16.67 11.67 5.00 1.67 1.67 1.67 6.67 1.67 90 1.67 5.00 16.67 10.00 48.33 1.67 1.67 1.67 3.33 -1.67 120 1.67 1.67 10.00 6.67 5.00 1.67 1.67 1.67 3.33 3.33 150 1.67 1.67 15.00 5.00 8.33 6.67 3.33 3.33 3.33 1.67 180 3.33 3.33 5.00 8.33 8.33 1.67 1.67 3.33 3.33 1.67 210 / / / / / / / / / /

Aver. 2.86 3.10 17.14 14.76 19.05 11.43 5.71 5.48 5.24 4.05

142

(8) Glycine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and glycine on reaction rate. APPENDIX (H.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Glycine]0 = 0.003 M For effect of Glycine, [Glycine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 2.42 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Glycine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.380 0.792 0.829 0.878 1.025 0.544 0.562 0.556 0.544 0.556 30 0.371 0.740 0.806 0.8729 0.998 0.512 0.538 0.531 0.519 0.538 60 0.3614 0.748 0.798 0.853 0.993 0.477 0.512 0.512 0.484 0.512 90 0.352 0.735 0.786 0.834 0.989 0.455 0.477 0.491 0.455 0.491 120 0.342 0.721 0.778 0.815 0.977 0.407 0.447 0.455 0.415 0.47 150 0.332 0.716 0.756 0.785 0.975 0.371 0.415 0.439 0.371 0.447 180 0.322 0.693 0.744 0.752 0.954 0.322 0.380 0.407 0.332 0.415 210 0.312 0.678 0.730 0.747 0.942 0.267 0.332 0.380 0.290 0.398 APPENDIX (H.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-glycine system according to the conditions applied in appendix (H.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Glycine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 1.65 23.33 11.67 3.57 21.67 8.33 6.67 6.67 6.67 5.00 30 1.65 -3.15 3.75 10.3 3.33 8.33 6.67 5.00 8.33 6.67 60 1.65 5.55 6.10 10.3 3.58 5.00 8.33 5.00 6.67 5.00 90 1.65 5.55 3.48 10.3 8.25 10.00 6.67 8.33 8.33 5.00 120 1.63 2.05 10.00 13.87 1.5 6.67 6.67 3.33 8.33 5.00 150 1.67 9.05 5.00 15.00 15 8.33 6.67 6.67 6.67 6.67 180 1.65 5.55 5.92 2.03 8.25 8.33 8.33 5.00 6.67 3.33 210 / / / / / / / / / /

Aver. 1.65 6.848 6.56 9.34 8.798 7.86 7.143 5.71 7.381 5.24

143

(9) Histidine -Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and histidine on reaction rate. APPENDIX (I.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Histidine]0 = 0.003 M For effect of Histidine, [Histidine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 3.01 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Histidine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.332 0.585 0.767 0.878 0.977 0.633 0.633 0.633 0.614 0.633 30 0.267 0.533 0.718 0.828 0.944 0.589 0.573 0.569 0.566 0.579 60 0.190 0.473 0.672 0.777 0.908 0.552 0.525 0.519 0.514 0.529 90 0.097 0.407 0.633 0.728 0.869 0.512 0.447 0.463 0.407 0.447 120 / 0.324 0.562 0.654 0.826 0.469 0.410 0.398 0.383 0.409 150 / 0.224 0.495 0.576 0.778 0.420 0.339 0.323 0.30 0.333 180 / 0.094 0.415 0.482 0.724 0.365 0.253 0.231 0.196 0.241 210 / / 0.317 0.360 0.663 0.303 0.147 0.114 0.06 0.125 APPENDIX (I.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate- histidine system according to the conditions applied in appendix (I.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Histidine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 10.00 14.72 20.83 27.38 23.25 14.02 18.60 19.86 14.1 17.00 30 10.00 14.45 17.50 24.65 23.25 10.4 13.00 13.35 14.1 13.65 60 10.00 14.16 13.33 21.3 23.25 10.58 18.40 13.35 23.74 19.35 90 10.00 14.74 21.67 28.00 23.25 10.22 7.60 13.35 4.46 7.95 120 10.00 14.45 17.50 24.65 23.25 10.4 13.00 13.35 14.1 13.65 150 10.00 14.45 17.50 24.65 23.25 10.4 13.00 13.35 14.1 13.65 180 10.00 14.45 17.50 24.65 23.25 10.4 13.00 13.35 14.1 13.65 210 / / / / / / / / / /

Aver. 10.00 14.49 17.98 25.04 23.25 10.92 13.80 14.28 14.1 14.13

144

(10) Leucine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and leucine on reaction rate. APPENDIX (J.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [leucine]0 = 0.003 M For effect of Leucine, [Leucine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.98 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Leucine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.230 0.512 0.690 0.792 0.878 0.544 0.556 0.562 0.550 0.544 30 0.211 0.477 0.607 0.765 0.857 0.534 0.543 0.546 0.527 0.515 60 0.177 0.434 0.567 0.718 0.819 0.523 0.529 0.530 0.502 0.483 90 0.141 0.398 0.522 0.665 0.777 0.513 0.514 0.513 0.476 0.449 120 0.102 0.332 0.473 0.604 0.73 0.502 0.499 0.495 0.448 0.412 150 0.041 0.243 0.398 0.477 0.643 0.491 0.484 0.478 0.418 0.372 180 0.010 0.199 0.353 0.451 0.62 0.479 0.467 0.457 0.386 0.327 210 / 0.130 0.290 0.380 0.544 0.467 0.451 0.437 0.352 0.277

APPENDIX (J.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-leucine system according to the conditions applied in appendix (J.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Leucine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 2.53 8.40 28.46 12.75 11.83 2.7 3.7 4.35 6.2 7.65 30 4.00 9.45 11.95 19.95 20.2 2.7 3.7 4.35 6.2 7.65 60 4.00 7.15 11.95 19.95 20.2 2.7 3.7 4.35 6.2 7.65 90 4.00 11.75 11.95 19.95 20.2 2.7 3.7 4.35 6.2 7.65 120 5.47 13.25 15.69 34.07 32.57 2.7 3.7 4.35 6.2 7.65 150 2.53 5.65 8.21 5.83 7.83 2.7 3.7 4.35 6.2 7.65 180 5.80 7.68 10.12 14.17 22.17 2.7 3.7 4.35 6.2 7.65 210 / / / / / / / / / /

Aver. 4.05 9.05 14.05 18.10 19.29 2.7 3.7 4.35 6.2 7.65

145

(11) Lysine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and Lysine on reaction rate. APPENDIX (K.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Lysine]0 = 0.003 M For effect of Lysine, [Lysine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. At µ = 0.25 M by potassium sulfate, pH = 1.92 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Lysine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.267 0.512 0.724 0.833 0.952 0.585 0.562 0.574 0.574 0.574 30 0.161 0.462 0.628 0.744 0.845 0.564 0.531 0.53 0.548 0.559 60 0.061 0.322 0.498 0.648 0.789 0.542 0.498 0.481 0.521 0.543 90 0.044 0.301 0.484 0.628 0.778 0.518 0.462 0.425 0.492 0.526 120 / 0.271 0.459 0.593 0.740 0.493 0.423 0.361 0.461 0.509 150 / 0.176 0.352 0.505 0.686 0.466 0.38 0.287 0.428 0.491 180 / 0.099 0.301 0.447 0.622 0.438 0.332 0.196 0.391 0.472 210 / / 0.162 0.338 0.549 0.408 0.278 0.082 0.352 0.453

APPENDIX (K.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-Lysine system according to the conditions applied in appendix (K.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Lycine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 13.33 11.67 35 .00 41.67 65.00 6.15 8.35 12.1 7.15 4.35 30 10 .00 26.67 36.67 36.67 28.33 6.15 8.35 12.1 7.15 4.35 60 1.41 3.33 3.33 6.67 5.00 6.15 8.35 12.1 7.15 4.35 90 5.95 4.48 5.84 10.98 16.73 6.15 8.35 12.1 7.15 4.35 120 5.98 12.18 20.83 24.02 21.6 6.15 8.35 12.1 7.15 4.35 150 5.92 8.12 8.33 13.33 22.00 6.15 8.35 12.1 7.15 4.35 180 5.95 10.15 18.24 20.7 21.8 6.15 8.35 12.1 7.15 4.35 210 / / / / / / / / / /

Aver. 6.93 10.94 18.32 22.00 25.78 6.15 8.35 12.1 7.15 4.35

146

(12) Methionine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and Methionine on reaction rate. APPENDIX (L.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Methionine]0 = 0.003 M For effect of Methionine, [Methionine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.90 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Methionine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.371 0.544 0.638 0.796 0.898 0.004 0.550 0.544 0.562 0.556 30 0.363 0.526 0.611 0.7678 0.868 0.004 0.543 0.525 0.532 0.546 60 0.354 0.508 0.582 0.738 0.837 0.003 0.537 0.5047 0.500 0.536 90 0.345 0.489 0.551 0.705 0.803 0.003 0.528 0.484 0.466 0.525 120 0.336 0.469 0.517 0.671 0.766 0.003 0.521 0.461 0.428 0.514 150 0.327 0.448 0.480 0.633 0.725 0.003 0.513 0.438 0.386 0.502 180 0.318 0.425 0.440 0.591 0.681 0.003 0.505 0.413 0.341 0.491 210 0.309 0.402 0.396 0.545 0.631 0.003 0.497 0.387 0.29 0.479

APPENDIX (L.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-methionine system according to the conditions applied in appendix (L.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Methionine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 1.5 4.65 8.85 13.05 17.25 0.562 1.95 5.05 8.1 2.8 30 1.5 4.65 8.85 13.05 17.25 0.545 1.95 5.05 8.1 2.8 60 1.5 4.65 8.85 13.05 17.25 0.526 1.95 5.05 8.1 2.8 90 1.5 4.65 8.85 13.05 17.25 0.507 1.95 5.05 8.1 2.8 120 1.5 4.65 8.85 13.05 17.25 0.487 1.95 5.05 8.1 2.8 150 1.5 4.65 8.85 13.05 17.25 0.466 1.95 5.05 8.1 2.8 180 1.5 4.65 8.85 13.05 17.25 0.444 1.95 5.05 8.1 2.8 210 / / / / / / / / / /

Aver. 1.5 4.65 8.85 13.05 17.25 0.420 1.95 5.05 8.1 2.8

147

(13) Phenyl alanine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and phenylalanine on the rate. APPENDIX (M.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Phenylalanine]0 = 0.003 M For effect of Phenylalanine, [Phenylalanine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. µ = 0.25 M, potassium sulfate, pH = 1.88 and Temp = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Phenylalanine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.342 0.597 0.785 0.906 1.002 0.591 0.5851 0.591 0.597 0.597 30 0.318 0.567 0.753 0.875 0.972 0.544 0.562 0.568 0.574 0.578 60 0.292 0.535 0.717 0.842 0.939 0.484 0.544 0.550 0.550 0.556 90 0.264 0.501 0.678 0.806 0.904 0.423 0.519 0.525 0.591 0.538 120 0.234 0.464 0.636 0.767 0.866 0.342 0.491 0.505 0.505 0.519 150 0.202 0.424 0.589 0.723 0.823 0.255 0.462 0.477 0.477 0.498 180 0.168 0.379 0.536 0.676 0.777 0.146 0.439 0.455 0.455 0.477 210 0.130 0.329 0.476 0.622 0.725 -0.022 0.407 0.423 0.423 0.439 APPENDIX (M.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-phenylalanine system according to the conditions applied in appendix (M.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Phenylalanine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 4.05 8.65 14.8 18.4 22.6 13.33 6.67 6.67 6.67 5.00 30 4.05 8.65 14.8 18.4 22.6 15.00 5.00 5.00 6.67 6.67 60 4.05 8.65 14.8 18.4 22.6 13.33 6.67 6.67 6.67 5.00 90 4.05 8.65 14.8 18.4 22.6 15.00 6.67 5.oo 5.00 5.00 120 4.05 8.65 14.8 18.4 22.6 13.33 6.67 6.67 6.67 5.00 150 4.05 8.65 14.8 18.4 22.6 13.33 5.00 5.oo 5.00 5.00 180 4.05 8.65 14.8 18.4 22.6 15.00 6.67 6.67 6.67 8.33 210 / / / / / / / / / /

Aver. 4.05 8.65 14.8 18.4 22.6 14.05 6.19 5.95 6.19 5.71

148

(14) Proline-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and proline on reaction rate. APPENDIX (N.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Proline]0 = 0.003 M For effect of Proline, [Proline]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.94 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Proline Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.423 0.653 0.839 0.961 1.006 0.585 0.58 0.607 0.585 0.591 30 0.352 0.597 0.778 0.898 0.947 0.568 0.556 0.574 0.562 0.568 60 0.279 0.525 0.699 0.823 0.875 0.550 0.531 0.544 0.538 0.556 90 0.190 0.447 0.607 0.732 0.789 0.538 0.512 0.512 0.512 0.525 120 0.061 0.342 0.491 0.613 0.672 0.512 0.484 0.477 0.477 0.498 150 / 0.217 0.332 0.455 0.525 0.505 0.455 0.431 0.439 0.477 180 / 0.021 0.079 0.217 0.322 0.477 0.431 0.380 0.415 0.447 210 / / / / / 0.47 0.398 0.332 0.380 0.415

APPENDIX (N.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-proline system according to the conditions applied in appendix (N.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Proline Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 13.33 18.33 30.00 41.67 43.33 5.00 6.67 10.00 6.67 6.67 30 11.67 20.00 33.33 41.67 45.00 5.00 6.67 8.33 6.67 3.33 60 11.67 18.33 31.67 41.67 45.00 3.33 5.00 8.33 6.67 8.33 90 13.33 20.00 31.67 43.33 48.33 6.67 6.67 8.33 8.33 6.67 120 11.67 18.33 31.67 41.67 45.00 1.67 6.67 10.00 8.33 5.00 150 13.33 20.00 31.67 40.00 41.67 6.67 5.00 10.00 5.00 6.67 180 12.5 18.33 31.67 41.67 45.0 1.67 6.67 8.33 6.67 6.67 210 / / / / / / / / / /

Aver. 12.5 19.05 31.67 41.67 44.76 4.29 6.19 9.05 6.90 6.19

149

(15) Serine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and Serine on reaction rate. APPENDIX (O.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Serine]0 = 0.003 M For effect of Serine, [Serine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 2.08 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Serine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.352 0.648 0.799 0.916 1.043 0.591 0.585 0.585 0.597 0.602 30 0.326 0.613 0.7924 0.889 1.036 0.556 0.556 0.556 0.562 0.556 60 0.298 0.597 0.785 0.884 1.028 0.525 0.512 0.538 0.519 0.505 90 0.268 0.585 0.768 0.878 1.02 0.484 0.491 0.512 0.477 0.439 120 0.236 0.58 0.755 0.860 1.012 0.439 0.447 0.477 0.423 0.371 150 0.201 0.556 0.732 0.857 1.003 0.389 0.398 0.447 0.371 0.290 180 0.164 0.544 0.724 0.848 0.995 0.332 0.352 0.415 0.301 0.190 210 0.123 0.523 0.714 0.838 0.986 0.290 0.301 0.380 0.230 0.041

APPENDIX (O.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-serine system according to the conditions applied in appendix (O.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Serine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 4.4 11.67 3.33 16.67 6.5 10.00 8.33 8.33 10.00 13.33 30 4.4 5.00 3.33 3.33 6.5 8.33 11.6 5.00 11.67 13.33 60 4.4 3.33 8.08 3.33 6.5 10.00 5.00 6.67 10.00 15.00 90 4.4 1.67 5.65 10.00 6.5 10.00 10.00 8.33 11.67 13.33 120 4.4 6.67 9.6 1.67 6.5 10.00 10.00 6.67 10.00 13.33 150 4.4 3.33 3.33 5.03 6.5 10.00 8.33 6.67 11.67 13.33 180 4.4 5.46 4.02 5.55 6.5 6.67 8.33 6.67 10.00 15.00 210 / / / / / / / / / /

Aver. 4.4 5.30 5.34 6.51 6.5 9.29 8.81 6.90 10.71 13.81

150

(16) Threonine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and threonine on reaction rate. APPENDIX (P.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Threonine]0 = 0.003 M For effect of Threonine, [Threonine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.99 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Threonine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.371 0.667 0.836 0.957 1.011 0.628 0.613 0.607 0.602 0.602 30 0.322 0.643 0.792 0.927 0.975 0.613 0.597 0.585 0.58 0.574 60 0.267 0.618 0.744 0.892 0.934 0.591 0.574 0.574 0.556 0.544 90 0.217 0.58 0.690 0.857 0.892 0.574 0.556 0.556 0.525 0.519 120 0.146 0.556 0.633 0.82 0.842 0.556 0.538 0.544 0.505 0.47 150 0.061 0.525 0.562 0.778 0.785 0.531 0.512 0.519 0.47 0.439 180 / 0.491 0.477 0.728 0.724 0.512 0.491 0.505 0.439 0.389 210 / 0.447 0.371 0.672 0.653 0.484 0.462 0.491 0.398 0.362

APPENDIX (P.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-threonine system according to the conditions applied in appendix (P.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Threonine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 8.33 8.33 21.67 20.00 26.67 5.00 5.00 6.67 6.67 8.33 30 8.33 8.33 21.67 21.67 28.33 6.67 6.67 3.33 6.67 8.33 60 6.67 11.67 21.67 20.00 26.67 5.00 5.00 5.00 8.33 6.67 90 8.33 6.67 20.00 20.00 28.33 5.00 5.00 3.33 5.00 11.67 120 8.33 8.33 21.67 20.00 28.33 6.67 6.67 6.67 8.33 6.67 150 6.67 8.33 21.67 21.67 26.67 5.00 5.00 3.33 6.67 10.00 180 8.33 10.00 21.67 21.67 26.67 6.667 6.67 3.33 8.33 5.00 210 / / / / / / / / / /

Aver. 7.86 8.81 21.43 20.71 27.38 5.71 5.71 4.52 7.14 8.1

151

(17) Tyrosine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and tyrisine on reaction rate. APPENDIX (Q.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Tyrosine]0 = 0.003 M For effect of Tyrosine, [Tyrosine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.46 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Tyrosine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.371 0.638 0.823 0.942 1.006 0.628 0.618 0.618 0.607 0.618 30 0.332 0.591 0.775 0.903 0.968 0.607 0.597 0.60 0.585 0.591 60 0.267 0.531 0.720 0.857 0.929 0.585 0.5798 0.5798 0.562 0.568 90 0.230 0.47 0.653 0.806 0.881 0.568 0.550 0.562 0.538 0.538 120 0.176 0.407 0.58 0.748 0.833 0.544 0.531 0.544 0.505 0.505 150 0.097 0.301 0.498 0.686 0.775 0.519 0.505 0.519 0.477 0.477 180 / 0.204 0.380 0.602 0.708 0.491 0.484 0.491 0.447 0.439 210 / 0.041 0.255 0.512 0.628 0.462 0.455 0.47 0.423 0.407

APPENDIX (Q.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-Tyrosine system according to the conditions applied in appendix (Q.1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Tyrosine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 6.67 15.00 23.33 25.00 28.33 6.67 6.67 5.00 6.67 8.33 30 10.00 16.67 23.33 26.67 26.67 6.67 5.00 6.67 6.67 6.67 60 5.00 15.00 25.00 26.67 30.00 5.00 8.33 5.00 6.67 8.33 90 6.67 13.33 23.33 26.67 26.67 6.67 5.00 5.00 8.33 8.33 120 8.33 18.33 21.67 25.00 28.33 6. 67 6.67 6. 67 6.67 6.67 150 8.33 13.33 25.00 28.33 28.33 6.67 5.00 6.67 6.67 8.33 180 8.33 16.67 20.00 25.00 28.33 6.67 6. 67 5.00 5.00 6.67 210 / / / / / / / / / /

Aver. 7.62 15.48 23.10 26.19 28.1 6.43 6.19 5.71 6.67 7.61

152

(18) Valine-Peroxodisulphate System Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of peroxodisulphate and valine on reaction rate. APPENDIX (R.1) For effect of peroxodisulphate, [K2S2O8]0 = flask (1) = 0.0025 M, flask (2) = 0.005 M, flask (3) = 0.0075 M, flask (4) = 0.010 M, flask (5) = 0.0125 M and [Valine]0 = 0.003 M For effect of Valine, [Valine]0 = flask (6) = 0.001M, flask (7) = 0.002 M, flask (8) = 0.003M, flask (9) = 0.004 M, flask (10) = 0.005 M. and [K2S2O8]0 = 0.005 M. at µ = 0.25 M by potassium sulfate, pH = 1.88 and Temp. = 333 K

3+ Log [K2S2O8] Effect of Peroxodisulphate Effect of Valine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 0.407 0.663 0.736 0.961 1.027 0.677 0.658 0.653 0.663 0.653 30 0.342 0.636 0.721 0.911 0.986 0.642 0.621 0.617 0.630 0.641 60 0.265 0.608 0.705 0.853 0.941 0.604 0.58 0.577 0.595 0.628 90 0.1726 0.578 0.688 0.787 0.891 0.563 0.535 0.534 0.558 0.615 120 0.0546 0.545 0.671 0.709 0.834 0.517 0.484 0.486 0.516 0.601 150 / 0.51 0.653 0.613 0.768 0.465 0.427 0.431 0.467 0.587 180 / 0.472 0.634 0.490 0.690 0.407 0.362 0.369 0.418 0.572 210 / 0.43 0.615 0.319 0.596 0.340 0.284 0.297 0.36 0.557

APPENDIX (R.2) Detailed results showing the average rate of reaction in each flask for peroxodisul-phate-valnine system according to the conditions applied in appendix (R1).

106 x Rate m.l-1.min-1. Effect of Peroxodisulphate Effect of Valine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (5)

Flask (6)

Flask (7)

Flask (8)

Flask (9)

Flask (10)

0 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 30 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 60 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 90 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 120 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 150 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 180 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25 210 / / / / / / / / / /

Aver. 11.8 9.1 6.35 33.65 31.95 12.2 12.5 12 11 4.25

153

APPENDIX (S) Detailed results showing the values of logarithm of peroxodisulphate concentration with time to determine the effect of temperature on reaction rate. APPENDIX (S.1) Effect of temperature on the oxidation of Alanine and Argnine: [K2S2O8]0 = 0.0075M, [Alanine]0 = [Argnine]0 = 0.003 M µ = 0.25 M, pH (for Alanine = 2.00, for Argnine = 1.94).

3+ Log [K2S2O8]

Oxidation of Alanine at Oxidation of Argnine at Time min.

600C 650C 700C 750C 600C 650C 700C 750C

0 0.736 0.720 0.740 0.732 0.699 0.732 0.724 0.720 30 0.721 0.686 0.633 0.616 0.676 0.681 0.642 0.581 60 0.702 0.621 0.494 0.456 0.655 0.629 0.52 0.375 90 0.682 0.557 0.290 0.20 0.632 0.568 0.352 / 120 0.661 0.505 / / 0.608 0.497 0.062 / 150 0.648 0.392 / / 0.562 0.411 / / 180 0.616 0.277 / / 0.557 0.305 / / 210 0.592 0.121 / / 0.528 0.164 / /

APPENDIX (S.2) Effect of temperature on the oxidation of Asparagine and Aspartic acid: [K2S2O8]0 = 0.0075 M, [Asparagine] = [Aspartic acid]0 = 0.003 M µ = 0.25 M, pH (for Asparagine = 1.91, for Aspartic acid = 1.84).

3+ Log [K2S2O8]

Oxidation of Asparagine at Oxidation of Aspartic acid at Time min.

600C 650C 700C 750C 800C 600C 650C 700C 750C 800C

0 0.775 0.658 0.789 0.708 0.763 0.803 0.695 0.81 0.736 0.763 30 0.748 0.648 0.729 0.627 0.613 0.789 0.672 0.739 0.643 0.664 60 0.740 0.585 0.688 0.512 0.519 0.752 0.597 0.684 0.519 0.535 90 0.708 0.538 0.642 0.407 0.342 0.740 0.585 0.621 0.404 0.351 120 0.686 0.491 0.591 0.279 / 0.720 0.574 0.547 0.267 0.025 150 0.677 0.462 0.556 / / 0.732 0.562 0.484 / / 180 0.653 0.439 0.467 / / 0.724 0.526 0.346 / / 210 0.633 0.354 0.389 / / 0.712 0.495 0.194 / /

154

APPENDIX (S.3) Effect of temperature on the oxidation of Cysteine and Glutamic acid: [K2S2O8]0 = 0.0075M, [Cysteine]0 = [Glutamic acid]0 = 0.003 M µ = 0.25 M, pH (for Cysteine = 3.30, for Glutamic acid = 2.02).

3+ Log [K2S2O8]

Oxidation of Cysetine at Oxidation of Glutamic acid at Time min.

600C 650C 700C 600C 650C 700C 7 800C

0 0.574 0.562 0.556 0.796 0.724 0.803 0.699 0.833 30 0.547 0.508 0.434 0.778 0.681 0.746 0.618 0.623 60 0.519 0.447 0.263 0.763 0.667 0.706 0.491 0.477 90 0.488 0.375 / 0.76 0.602 0.661 0.4 0.24 120 0.456 0.289 / 0.752 0.574 0.623 0.290 / 150 0.420 0.182 / 0.744 0.538 0.58 / / 180 0.382 0.039 / 0.724 0.477 0.490 / / 210 0.34 / / 0.712 0.420 0.414 / /

APPENDIX (S.4) Effect of temperature on the oxidation of Glutamine and Glycine: [K2S2O8]0 = 0.0075 M, [Glutamine]0= [Glycine]0 = 0.003 M µ = 0.25 M, pH (for Glutamine = 3.51, for Glycine= 2.42).

3+ Log [K2S2O8]

Oxidation of Glutamine at Oxidation of Glycine at Time min.

600C 650C 700C 750C 800C 600C 650C 700C 750C

0 0.900 0.875 0.881 0.869 0.866 0.715 0.738 0.778 0.767 30 0.860 0.845 0.866 0.80 0.764 0.695 0.718 0.621 0.553 60 0.796 0.812 0.850 0.72 0.631 0.674 0.697 0.374 0.114 90 0.76 0.777 0.834 0.62 0.437 0.652 0.674 / / 120 0.720 0.739 0.818 0.489 0.079 0.63 0.650 / / 150 0.695 0.697 0.801 0.303 / 0.605 0.625 / / 180 0.653 0.650 0.783 / / 0.58 0.598 / / 210 0.638 0.598 0.764 / / 0.552 0.57 / /

155

APPENDIX (S.5) Effect of temperature on the oxidation of Histidine and Leucine: [K2S2O8]0 = 0.0075M, [Histidine]0 = 0.003 M , [Leucine]0 = 0.003 M µ = 0.25 M, pH (for Histidine = 3.01, for Leucine = 1.98).

3+ Log [K2S2O8]

Oxidation of Histidine at Oxidation of Leucine at Time min.

600C 650C 700C 750C 800C 600C 650C 700C 750C

0 0.708 0.728 0.740 0.712 0.695 0.712 0.700 0.721 0.720 30 0.701 0.706 0.694 0.641 0.591 0.707 0.660 0.640 0.597 60 0.694 0.682 0.643 0.557 0.454 0.701 0.616 0.541 0.425 90 0.686 0.656 0.585 0.452 0.254 0.696 0.567 0.412 0.134 120 0.679 0.629 0.517 0.314 / 0.691 0.512 0.228 / 150 0.672 0.600 0.438 0.11 / 0.685 0.449 / / 180 0.664 0.57 0.340 / / 0.66 0.374 / / 210 0.656 0.536 0.214 / / 0.674 0.285 / /

APPENDIX (S.6) Effect of temperature on the oxidation of Lysine and Methionine: [K2S2O8]0 = 0.0075 M, [Lysine]0 = [Methionine] = 0.003 M µ = 0.25 M, pH (for Lysine = 1.92, for Methionine = 1.90).

3+ Log [K2S2O8]

Oxidation of Lysine at Oxidation of Methionine at Time min.

600C 650C 700C 600C 650C 700C 750C 800C

0 0.681 0.7521 0.732 0.659 0.699 0.681 0.663 0.653 30 0.634 0.6703 0.585 0.637 0.658 0.631 0.591 0.558 60 0.592 0.569 0.390 0.614 0.612 0.574 0.504 0.436 90 0.546 0.438 0.028 0.59 0.561 0.508 0.395 0.266 120 0.398 0.248 / 0.564 0.503 0.430 0.25 / 150 0.435 / / 0.537 0.436 0.336 0.031 / 180 0.368 / / 0.508 0.357 0.215 / / 210 0.301 / / 0.476 0.26 0.047 / /

156

APPENDIX (S.7) Effect of temperature on the oxidation of Phenylalanine and Proline: [K2S2O8]0 = 0.0075M, [Phenylalanine]0 = 0.003 M , [Proline]0 = 0.003 M µ = 0.25 M, pH (for Phenylalanine = 1.88, for Proline = 1.94).

3+ Log [K2S2O8]

Oxidation of Phenylalanine at Oxidation of Proline at

Time min.

600C 650C 700C 600C 650C 700C 750C

0 0.720 0.731 0.751 0.728 0.714 0.675 0.699 30 0.678 0.669 0.613 0.711 0.681 0.596 0.588 60 0.654 0.596 0.41 0.693 0.645 0.501 0.437 90 0.628 0.508 0.015 0.675 0.606 0.377 0.206 120 0.601 0.398 / 0.656 0.563 0.204 / 150 0.572 0.251 / 0.636 0.515 / / 180 0.541 0.026 / 0.616 0.461 / / 210 0.508 / / 0.594 0.400 / /

APPENDIX (S.8) Effect of temperature on the oxidation of Serine and Threonine: [K2S2O8]0 = 0.0075 M, [Serine]0 = [Threonine] = 0.003 M µ = 0.25 M, pH (for Serine = 2.08, for Threonine = 1.99).

3+ Log [K2S2O8]

Oxidation of Serine at Oxidation of Threonine at Time min.

600C 650C 700C 750C 800C 600C 650C 700C 750C

0 0.799 0.799 0.78 0.793 0.78 0.712 0.756 0.727 0.728 30 0.792 0.782 0.762 0.756 0.624 0.702 0.699 0.650 0.613 60 0.785 0.764 0.744 0.717 0.378 0.691 0.659 0.557 0.457 90 0.768 0.746 0.725 0.673 / 0.681 0.608 0.439 0.209 120 0.755 0.727 0.705 0.624 / 0.67 0.549 0.275 / 150 0.732 0.707 0.684 0.569 / 0.659 0.481 0.009 / 180 0.724 0.686 0.662 0.506 / 0.647 0.400 / / 210 0.714 0.664 0.639 0.432 / 0.636 0.301 / /

157

APPENDIX (S.9) Effect of temperature on the oxidation of Tyrosine and Valine: [K2S2O8]0 = 0.0075M, [Tyrosine]0 = 0.003 M , [Valine]0 = 0.003 M µ = 0.25 M, pH (for Tyrosine = 1.46, for Valine = 1.88).

3+ Log [K2S2O8]

Oxidation of Tyrosine at Oxidation of Valine at Time min.

600C 650C 700C 750C 600C 650C 700C 750C

0 0.748 0.732 0.728 0.699 0.687 0.703 0.706 0.709 30 0.738 0.717 0.650 0.586 0.677 0.677 0.628 0.600 60 0.727 0.702 0.555 0.432 0.667 0.650 0.534 0.454 90 0.716 0.686 0.432 0.191 0.657 0.621 0.412 0.234 120 0.705 0.669 0.261 / 0.646 0.591 0.243 / 150 0.694 0.652 / / 0.635 0.557 / / 180 0.682 0.634 / / 0.623 0.5214 / / 210 0.667 0.615 / / 0.612 0.482 / /

158

APPENDIX (T) Detailed results showing the values of rate of reaction to determine the effect of solvent composition. APPENDIX (T.1) Effect of solvent composition on the oxidation of Alanine and Asparagine Acetic acid in Flask (1) 5%, Flask (2) 10%, and Flask (3) 15% v/v [K2S2O8]0 = 0.005M, [Alanine]0 = [Asparagine]0 = 0.003 M µ = 0.25 M, and Temperature = 333 K

3+Log[K2S2O8] Oxidation of Alanine Oxidation of Asparagine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.711807 0.694605 0.69897 0.698735 0.703291 0.69897 30 0.681693 0.667453 0.670153 0.669782 0.669596 0.678473 60 0.649335 0.638489 0.639287 0.633468 0.63809 0.65215 90 0.61437 0.607455 0.606059 0.59384 0.604118 0.624127

120 0.576341 0.574031 0.570076 0.550228 0.567262 0.594171 150 0.534661 0.537819 0.53084 0.501744 0.526985 0.561995 180 0.488551 0.498311 0.487704 0.447158 0.482588 0.527243 210 0.436957 0.454845 0.439806 0.384712 0.43313 0.489466

APPENDIX (T.2) Effect of solvent composition on the oxidation of Cysteine and Glutamic acid: Acetic acid in Flask (1) 5%, Flask (2) 10%, and Flask (3) 15% v/v [K2S2O8]0 = 0.0075M, [Cysteine]0 = [Glutamic acid]0 = 0.003 M µ = 0.25 M, and Temperature = 333 K

3+Log[K2S2O8] Oxidation of Cysteine Oxidation of Glutamic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.69897 0.703291 0.69897 0.69897 0.69897 0.694605 30 0.653695 0.658345 0.654417 0.666518 0.666518 0.667453 60 0.603144 0.608205 0.604766 0.631444 0.631444 0.638489 90 0.545925 0.551511 0.548696 0.593286 0.593286 0.607455

120 0.480007 0.486289 0.4843 0.55145 0.55145 0.574031 150 0.402261 0.40951 0.408664 0.50515 0.50515 0.537819 180 0.307496 0.31618 0.317018 0.453318 0.453318 0.498311 210 0.186108 0.197143 0.200714 0.394452 0.394452 0.454845

159

APPENDIX (T.3) Effect of solvent composition on the oxidation of Lysine and Phenylalanine: Acetic acid in Flask (1) 5%, Flask (2) 10%, and Flask (3) 15% v/v [K2S2O8]0 = 0.0075 M, [Lysine]0 = [Phenylalanine] = 0.003 M µ = 0.25 M, and Temperature = 333 K

3+Log[K2S2O8]

Oxidation of Lysine Oxidation of Phenylalanine Time min.

Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.69897 0.69897 0.69897 0.703291 0.703291 0.703291 30 0.653261 0.659155 0.670988 0.63829 0.63829 0.641425 60 0.602169 0.615319 0.641077 0.561817 0.561817 0.569257 90 0.544254 0.566555 0.608954 0.468938 0.468938 0.482659

120 0.477411 0.511616 0.574263 0.350636 0.350636 0.374382 150 0.398374 0.448706 0.536558 0.187521 0.187521 0.22981 180 0.301681 0.375115 0.495267 -0.07676 -0.07676 0.01157 210 0.177103 0.286456 0.449633 -0.86646 -0.86646 -0.44794

APPENDIX (T.4) Effect of solvent composition on the oxidation of Serine: Acetic acid = Flask (1) 5%, Flask (2) 10%, and Flask (3) 15% v/v [K2S2O8]0 = 0.0075 M, [Serine]0 = 0.003 M, µ = 0.25 M, and Temp. = 333 K

3+Log[K2S2O8]

Oxidation of Serine Time min.

Flask (1)

Flask (2)

Flask (3)

0 0.703291 0.703291 0.703291 30 0.646453 0.646453 0.649676 60 0.581039 0.581039 0.588496 90 0.503995 0.503995 0.517262

120 0.410271 0.410271 0.432007 150 0.290591 0.290591 0.325823 180 0.12483 0.12483 0.184975 210 -0.14661 -0.14661 -0.0248

160

APPENDIX (U) Detailed results showing the values of rate of reaction to determine the effect of catalysts. APPENDIX (U.1) Effect of catalyst on the oxidation of Alanine and Asparagine: Flask (1) = 0.005 M (Ag+), Flask (2) 0.005 M (Cu2+), and Flask (3) =0.0025 (Ag+) + 0.0025 (Cu2+) M, [K2S2O8]0 = 0.005M, [Alanine]0 = [Asparagine]0 = 0.003 M, µ = 0.25 M, pH (for Alanine = 2.00, for Asparagine = 1.91), and Temp. = 333 K

106 x Rate m.l-1.min-1. Oxidation of Alanine Oxidation of Asparagine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (1)

Flask (2)

Flask (3)

Flask (4)

0 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 30 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 60 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 90 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 120 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 150 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 180 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35 210 3.35 32.7 17.3 35.6 3.95 39.4 14.4 41.35

APPENDIX (U.2) E Effect of catalyst on the oxidation of Cysteine and Glutamic acid: Flask (1) = 0.005 M (Ag+), Flask (2) 0.005 M (Cu2+), and Flask (3) =0.0025 (Ag+) + 0.0025 (Cu2+) M. [K2S2O8]0 = 0.0075M, [Cysteine]0 = [Glutamic acid]0 = 0.003 M, µ = 0.25 M pH ( Cysteine = 3.30, Glutamic acid = 2.02) and Temp. = 333 K

106 x Rate m.l-1.min-1. Oxidation of Cysteine Oxidation of Glutamic acid

Time min.

Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (1)

Flask (2)

Flask (3)

Flask (4)

0 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 30 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 60 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 90 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 120 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 150 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 180 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55 210 9.55 25.95 21.15 37.5 13.55 42.3 24.05 61.55

161

APPENDIX (U.3) Effect of catalyst on the oxidation of Lysine and Phenylalanine: Flask (1) = 0.005 M (Ag+), Flask (2) 0.005 M (Cu2+), and Flask (3) =0.0025 (Ag+) + 0.0025 (Cu2+) M. [K2S2O8]0 = 0.0075 M, [Lysine]0 = [Phenylalanine] = 0.003 M µ = 0.25 M, pH (for Lysine = 1.92, for Phenylalanine = 1.88) and Temp. = 333 K

106 x Rate m.l-1.min-1.

Oxidation of Lysine Oxidation of Phenylalanine Time min.

Flask (1)

Flask (2)

Flask (3)

Flask (4)

Flask (1)

Flask (2)

Flask (3)

Flask (4)

0 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 30 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 60 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 90 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 120 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 150 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 180 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6 210 10.15 25 19.25 32.7 8.65 41.95 35.5 51.6

APPENDIX (U.4) Effect of catalyst on the oxidation of Serine: Flask (1) = 0.005 M (Ag+), Flask (2) 0.005 M (Cu2+), and Flask (3) =0.0025 (Ag+) + 0.0025 (Cu2+) M. [K2S2O8]0 = 0.0075 M, [Serine]0 = 0.003 M, µ = 0.25 M, pH (for Serine = 2.08) and Temp. = 333 K

106 x Rate m.l-1.min-1.

Oxidation of Serine Time min.

Flask (1)

Flask (2)

Flask (3)

Flask (4)

0 4.7 51.6 45.15 90.3 30 4.7 51.6 45.15 90.3 60 4.7 51.6 45.15 90.3 90 4.7 51.6 45.15 90.3

120 4.7 51.6 45.15 90.3 150 4.7 51.6 45.15 90.3 180 4.7 51.6 45.15 90.3 210 4.7 51.6 45.15 90.3

162

APPENDIX (V) Detailed results showing the values of logarithm of peroxodisulphate concentration to determine the effect of Ionic Strength. APPENDIX (V.1) Effect of Ionic strength on the oxidation of Alanine and Asparagine: Ionic Strength in Flask (1) 0.25, Flask (2) 0.3, and Flask (3) 0.35 M. [K2S2O8]0 = 0.005M, [Alanine]0 = [Asparagine]0 = 0.003 M, and Temperature = 333 K

3+ Log [K2S2O8] Oxidation of Alanine Oxidation of Asparagine Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.69741 0.6902 0.67669 0.694605 0.703291 0.711807 30 0.68856 0.66642 0.65345 0.684082 0.681467 0.683047 60 0.67952 0.64128 0.6289 0.673297 0.658488 0.652246 90 0.6703 0.61458 0.60287 0.662238 0.634225 0.619093

120 0.66087 0.58614 0.57519 0.65089 0.608526 0.583199 150 0.65123 0.5557 0.54562 0.639237 0.58121 0.544068 180 0.64138 0.52297 0.51388 0.627263 0.55206 0.501059 210 0.6313 0.48756 0.47965 0.61495 0.520811 0.453318

APPENDIX (V.2) Effect of Ionic strength on the oxidation of Cysteine and Glutamic acid: Ionic Strength in Flask (1) 0.25, Flask (2) 0.3, and Flask (3) 0.35 M. [K2S2O8]0 = 0.0075M, [Cysteine]0 = [Glutamic acid]0 = 0.003 M, and Temperature = 333 K

3+ Log [K2S2O8] Oxidation of Cysteine Oxidation of Glutamic acid Time

min. Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.69897 0.70757 0.69897 0.69897 0.716003 0.69897 30 0.673344 0.670941 0.646992 0.662144 0.680109 0.658011 60 0.64611 0.630936 0.587935 0.621903 0.640978 0.612784 90 0.617053 0.586868 0.519566 0.577549 0.597969 0.562293

120 0.585912 0.537819 0.438384 0.528145 0.550228 0.50515 150 0.552364 0.482516 0.338456 0.472391 0.496584 0.439333 180 0.516006 0.419129 0.208441 0.40841 0.435367 0.361728 210 0.476324 0.344883 0.022016 0.333346 0.364082 0.267172

163

APPENDIX (V.3) Effect of Ionic strength on the oxidation of Lysine and Phenylalanine: Ionic Strength in Flask (1) 0.25, Flask (2) 0.3, and Flask (3) 0.35 M. [K2S2O8]0 = 0.0075 M, [Lysine]0 = [Phenylalanine] = 0.003 M, and Temperature = 333 K

3+ Log [K2S2O8]

Oxidation of Lysine Oxidation of Phenyl alanine Time min.

Flask (1)

Flask (2)

Flask (3)

Flask (1)

Flask (2)

Flask (3)

0 0.703291 0.711807 0.711807 0.69897 0.69897 0.69897 30 0.676282 0.684262 0.681829 0.675824 0.683182 0.675687 60 0.647481 0.65485 0.649627 0.651375 0.666799 0.651084 90 0.616633 0.623301 0.614845 0.625467 0.649773 0.625004

120 0.583426 0.589279 0.577032 0.597914 0.632052 0.597256 150 0.547467 0.552364 0.535611 0.568495 0.613578 0.567614 180 0.50826 0.512017 0.489818 0.536937 0.594282 0.5358 210 0.46516 0.467534 0.438621 0.502905 0.574089 0.50147

APPENDIX (V.4) Effect of Ionic strength on the oxidation of Serine: Ionic Strength in Flask (1) 0.25, Flask (2) 0.3, and Flask (3) 0.35 M. [K2S2O8]0 = 0.0075 M, [Serine]0 = 0.003 M, and Temperature = 333 K

3+ Log [K2S2O8]

Oxidation of Serine Time min.

Flask (1)

Flask (2)

Flask (3)

0 0.703291 0.70757 0.70757 30 0.690993 0.684756 0.694474 60 0.678336 0.660676 0.68097 90 0.665299 0.635182 0.667032

120 0.651859 0.608098 0.652633 150 0.63799 0.579212 0.63774 180 0.623663 0.548267 0.622318 210 0.608847 0.514946 0.606328

164

APPENDIX (W) List of standard amino acids Structures and symbols of the 20 amino acids which are directly encoded for protein synthesis by the standard genetic code.

L-Alanine (Ala / A)

L-Arginine (Arg / R)

L-Asparagine (Asn / N)

L-Aspartic acid (Asp / D)

L-Cysteine (Cys / C)

L-Glutamic acid (Glu / E)

L-Glutamine (Gln / Q)

L-Glycine (Gly / G)

L-Histidine (His / H)

L-Isoleucine (Ile / I)

L-Leucine (Leu / L)

L-Lysine (Lys / K)

L-Methionine (Met / M)

L-Phenylalanine (Phe / F)

L-Proline (Pro / P)

L-Serine (Ser / S)

L-Threonine (Thr / T)

L-Tryptophan (Trp / W)

L-Tyrosine (Tyr / Y)

L-Valine (Val / V)

165

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