STUDY OF STABILITY CONSTANTS OF METAL COMPLEXES WITH

SULFA DRUGS

DISSERTATION SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS

FOR THE AWARD OF THE DEGREE OF

iWajftcr of $l)rtos(opl)|» IN

CHEMISTRY

BY

MOHD. NAZIM

UNDER THE SUPERVISION OF

PROF. K. S. SIDDIQI

DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY

ALIGARH (INDIA)

2008

V «.'"

, ^ ^ " - ' i %

^ 9 S y 20tt

D,S4II50

^lof. *?C< . Siddifi/ DEPARTMENT OF CHEMISTRY ALIGARH MUSLIM UNIVERSITY ALIGARH-202002 (INDIA)

Certificate

This is to certify that the dissertation entitled ''Study of Stability

Constants of Metal Complexes with Sulfa Drugs" is the original work

carried out by Mr. Mohd. Nazim under my supervision and is suitable for

submission for the award of M.Phil, degree in Chemistry of the Aligarh

Muslim University, Aligarh.

V A

Date a^^ (Prof. K.S. Siddi

In the name of ALLAH, the most beneficent and merciful who enabled me to complete this work,

I hereby wish to express my sincere thanks and deep sense of

gratitude to my supervisor Prof, K.S. Siddiqi, Department of Chemistry,

Aligarh Muslim University, Aligarh for his keen interest and help in this

work and interpretation of the results, I am particularly grateful to him for

his encouragement, removing the obstacles and patching my ignorance.

Thanks are also due to Prof, (Mrs,) A. Lai, Chairperson, Department

of Chemistry Aligarh Muslim University, Aligarh for providing necessary

research facilities,

I am also thankful to all of my lab colleagues specially

Ahmad Husain and friends specially Mohd. Shahid for their cooperation

and useful suggestions during this work.

Eventually I would like to express my gratitude to my parents for their

belief in me. They always encourage and guide me properly. No words can

express my indebtedness to my uncle Mr. Shafiq Ahmad, who always

inspired me to strive hard to achieve my goal.

Mohd. Nazim

ABSTRACT

The potentiometric studies have been carried out on divalent metal

ions with sulfamethazine, Pyridoxal 5-phosphate and DL-cystine. The

Calvin-Bjerrum pH-metric titration technique, as modified by Irving and

Rossotti, has been applied to determine the stepwise protonation constants

and stepwise formation constants of the metal complexes in aqueous

medium. The ionic strength was maintained constant at 0,1 M by adding

NaCl(lM) solution.

The results indicate that one proton of sulfamethazine and two protons

each of pyridoxal 5-phosphate and DL-Cystine were being released in the

pH range 4.0 to 10.0.

Stability of sulfamethazine, Pyridoxal-5-phosphate and DL-Cystine

upon complexation with divalent metal ions deceases with an increase in pH

following Irving William series:

Mn(II) < Fe(II) < Co(II) < Ni(II) < Cu(II)

The thermodynamic parameters that is, Gibbs free energy (AG),

Entropy (AS) and Enthalpy (AH) have also evaluated from the stability

constant.

CONTENTS

Title Page No.

1. Introduction 1-11

2. Experimental 12-15

3. Results and Discussion 16-59

4. References 60-65

INTRODUCTION

The sulfa drugs are synthetic antimicrobial agents with a wide

spectrum encompassing most gram positive and gram negative

organisms. These drugs are employed systematically for the

prevention and cure of bacterial infections.

The first sulfa drug, prontosil was discovered by the German

physician and chemist, Gerhard Domagk in 1935. He observed that

prontosil protects mice against streptococcal infections. The

sulfonamide is a generic name for derivatives of para-

aminobenzenesulfonamide (sulfanilamide). In the liver it exists as

follows:

H2N- /y -N=N-HfJ ) -S02-NH2 ^ ^ ^ > H^N-Z^V-NH^ 2i

Prontosil NH2 +

H 2 N - / V-SO2NH2

Sulfanilamide

Sulfa drugs are white crystalline powder and sparingly

soluble in water. They readily react with alkali and form salts

which are more soluble than the parent drug.

They kill bacteria and fungi by interfering with their

metabolism. They were the "wonder drugs" before penicillin was

discovered. They are used mainly for the treatment of urinary tract

infections (UTI).

Sulfa drugs are usually not allergic by themselves but when a

sulfonamide molecule is metabolized in the body, it becomes

capable of attaching to proteins. The allergy is, therefore, not

related to the original drug but a drug-protein complex. The

antibacterial activity of sulfonamides is related to its pKa value.

Maximum activity is found in those having pKa value between

6.0-7.5. Various types of complexes of sulfa drugs Schiff bases of

rare-earth metal ions were synthesized and characterized by

potentiometric studies earlier [1,2].

Due to the presence of amino group, sulfa drugs form schiff

bases with various types of aldehyde and keto compounds. The

voltammetric and potentiometric studies of some sulfa drug-Schiff-

base compounds and their metal complexes have been described

earlier [3]. Spectral and thermal studies of ternary complexes of

nickel with sulfasalazine and some amino acids were studied in

which sulfasalazine is used as primary ligand and amino acids as

secondary ligand [4].

N=^ *^M-H >-S02-NH-< )

The metal complexes of Schiff-base derived from sulfadiazine

and an amine display obvious inhibitor effects against gram

positive and gram negative bacteria. Transition metal complexes of

Schiff base derived from sulfadiazine and salicylaldehyde have

been extensity studied by UV-Vis, conductance, polarography, ESR

and TGA. They have also exhibited antibacterial and anti-fungal

properties in solution [5],

\ /

NHR'

^ B u ^ O ^ \ = /

NHR

The organotin (IV) form complexes with biologically active

sulfa drug Schiff base. They have also been found to be highly

active against bacteria and fungi [6]. Sulfamethazine has been used

to treat bacterial diseases in human and veterinary medicine and to

promote growth in cattle, sheep, pigs and poultry. It produces

thyroid-tumours in mice and rats by a non-genotoxic mechanism,

which involves inhibition of thyroid peroxidase resulting in

alterations in thyroid ,

(sulfamethazine)

Sulfamethazine is widely used in veterinary medicine in

combination with chlorotetracycline and penicillin in pigs for

maintenance of weight gain in the presence of atrophic rhinitis,

growth promotion and increased feed efficiency. The potentiometric

and thermodynamic studies of 2-acrylamidosulfamethazine and its

metal complexes in mono and polymeric forms have been

characterized by elemental analysis and IR studies in 40% (v/v)

ethanol water mixture [7].

Synthesis and equilibrium studies of metal complexes of

sulfamethazine have been investigated by potentiometric techniques

in acidic solution and also characterized by elemental analysis,

electrolytic conductance, IR, 'H NMR, mass spectra and

thermogravimetric studies [8].

It is water soluble and colourless in aqueous solution. Due to

the presence of free amino group, it forms Schiff-base with

aldehyde and keto compounds easily.

, , . . . „._// \ S O 2 — N H — ^

(Schiff base of sulfamethazine and PLP)

Pyridoxal-5-phosphate is metabolically active form of

vitamin B6. It works as a coenzyme in a large group of specific

enzyme catalyzed reactions of amino-group transfer,

decarboxylation and other metabolic transformations of individual

amino acids [9].

(Pyridoxal 5-phosphate)

Vitamin B6 is first Okidlzed to pyridoxal and rapidly

phosphorylated to pyridoxal 5-phosphate in liver. It is the main

circulating form of vitamin B6. It is a brown powder and dissolves

in water to give a pale yellow colour. Only this form can be used by

enzyme involved in the biological process of nitrogen, protein and

haeme synthesis [10].

Pyridoxal 5-phosphate is used as coenzyme in many

biological reactions. Coenzymes are low molecular weight

substances that provide unique chemical functionalities for certain

enzyme complexes.

All water soluble vitamins, except vitamin-C are coenzymes

or precursor of coenzymes. They can act as carriers of specific

functional groups i.e. methyl, acetyl groups. . The second order rate

constant for complex formation between PLP and cysteine has been

determined as function of pH [11]. Pyridoxal 5-phosphate participates

in a wide variety of reactions involving amino acids, as:

1. Transamination reaction.

2. a and p decarboxylation.

3. p and y elimination.

4. Racemization.

5. Aldol reactions

These reactions occur due to the ability of pyridoxal 5-phosphate

to form stable Schiff base with amino group of a amino acids [12].

The rate constant of hydrolysis and formation of Schiff bases

formed by pyridoxal 5-phosphate with L-tryptophan and their

methyl and n-butyl esters at variable pH and IS^C and at O.IM

ionic strength has been reported [13]. Pyridoxal 5-phosphate has

two absorption maxima, one at 415 nm in acidic medium and the

other at 340 nm in basic medium. PLP has been found to undergo a

rapid benzoin type condensation in oxygen free solution when

exposed to light [14, 15].

The screening of PLP dependent activities toward amino

acids were studied which confirms that the Schiff base is an

obligatory intermediate in all PLP dependent reactions of amino

acids [16]. The role of PLP in the structural stabilization has been

studied by fluorescence and ^ P NMR [17].

Equilibria and absorpation spectra of various Schiff bases of

PLP have been studied and the kinetic and thermodynamic

parameters have been evaluated in aqueous solution as a function of

pH and constant ionic strength [18-20].

Since cystine partially dissolves in water, 1-2 drops of acid

were added to make it water soluble.lt contains two amino groups

and two carboxyKc acid groups with a disulphide bond. It is also

found in the zwitter-ionic form in solution. As a component, cystine

is a significant determinant of the tertiary structure of most

proteins.

COOH ^COOH

CH-CH2-S-S-CH2-CI^

NH2^ NH2

(Cystine)

The excess of cystine in the body can cause cystinosis, a rare disease

that can cause cystine crystals to form in the body and produce bladder or

kidney stones. This side effect has not been associated with cysteine.

Cysteine is more easily absorbed by the body than cystine, and hence

most supplements contain cysteine rather than cystine.

However, cysteine is unstable, and is often converted to

cystine in the body. To avoid the conversion of cysteine to

potentially harmful amounts of cystine, intake of vitamin C

supplements is advised. It is particularly abundant in skeletal and

connective tissues, hair and digestive enzymes. Due to the presence

of sulphur, metal complexes of DL-cystine have an important role

in therapeutic action against toxic heavy metals.

Cysteine is oxidized to cystine at neutral pH. In acidic

medium, reaction does not take place while cystine is a non

essential amino acid and synthesized by the body itself [21].

However, food such as eggs, meat, dairy products, and whole grains

are also good sources of cystine. It functions as a powerful anti­

oxidant in detoxifying harmful toxins, protects the body from

radiation damage, liver and brain damage due to alcohol, drugs, and

toxic compounds. It has been also used in the treatment of

rheumatoid arthritis and hardening of the arteries, promotes the

recovery from severe burns and surgery, promotes the burning of

fat and the building of muscle and slows down the aging process.

Cysteine is used as a secondary ligand in the complex reactions

with transition metal ions reported earlier [22]. People suffering from

diabetes should be careful while taking supplementation, as it could

inactivate insulin.

H HC=S I

H2N—CH— COOH

L-cystine is used as precursor of p-lactum ring of penicillin.

For biosynthesis of penicillin, cystine related compounds are

necessary. Various methods of synthesis of a and p-methyl DL-

cystine from DL-allothreonine are described [24].

Many a-substituted cystines are synthesized by using a-

haloketone condensed with sodium salt of a-toluenethiol followed

by treatment with KCN and ammonium carbonate in aqueous

ethanol (50%) solution at 60-70 T , hydrolyzed then reduced [25].

Adsorption and desorption behaviour of cystine and cysteine

in neutral and basic media were studied with gold(III) in which

absorption of cystine takes place at different potentials. It confirms

that cysteine monolayer presents a higher packing density than

cystine [26].

The polarographic study of cystine in O.IN HCl with thymol

as suppresser has been done. The system was not reversible and

potentiometric titration curves were obtained for a-substituted

cystines [27]. Incorporation of sulphur from DL-cystine to

glutathione and protein in rat has indicated that cystine-sulphur was

incorporated much more rapidly in glutathione than in proteins[28].

Metabolic studies of DL-cystine in rat showed that cystine was

converted into pyruvate first, then it gives alanine, aspartic acid and

glutamic acid [29].

10

In this project, the stability constant and thermodynamic

parameters of sulfamethazine, pyridoxal 5-phosphate and DL-cystine

were determined by Calvin-Bjerrum pH-metric titration technique, as

modified by Irving and Rossotti at O.IM ionic strength at room

temperature. The strength of complexation between organic ligand and

metal ions is usually expressed in terms of stability constant (LogK).

The composition of the complex formed during the

interaction of transition metal ions with PLP, sulfamethazine and

DL-cystine were established by measuring the magnitude of proton

displacement during the titration of these molecules in absence and

presence of metal ions.

Metal ligand stoichiometry of all the ligands was also

determined by conductivity measurements.

11

EXPERIMENTAL

Apparatus and Reagents:

All chemicals were obtained from BDH or Loba Chemie.

Hydrated metal chlorides (BDH), sulfamethazine(sigma), pyridoxal

5-phosphate(Koch-Light) and DL-cystine were used as received. A

digital pH meter (Elico LI-120) fitted with combined glass

electrode was used.

Titration procedure:

The following solutions were prepared in double distilled water.

1. O.INHCIO4

2. 1 M NaCl

3. 0.005 M Ligand

4. 0.005 M Metal Chloride

5. 0.2 M NaOH

Strength of stock solution of perchloric acid was determined

by titrating against standard NaOH solution. Calvin-Bjerrum

titration technique, as modified by Irving and Rossotti was

followed [30-32]. The following solutions were titrated separately

against standard NaOH (0.2 M). 1 Molar NaCl solution was added

to maintain the constant ionic strength (|i = 0.1). The total volume

12

in each case was made up to 50 mL by adding appropriate volume

of CO2 free double distilled water. The metal contents were

estimated by standard methods using EDTA titration [33].

Solution A: 5 mL, 0.1 N HCIO4 + 4.5 mL, 1 M NaCl + 40.5 water

Solution B: 5 mL, 0.1 N HCIO4 + 4.5 mL, 1 M NaCl + 20 mL, 0.005

M Ligand + 20.5 water

Solution C: 5 mL, 0.1 N HCIO4 + 4.5 mL, 1 M NaCl + 20 mL, 0.005

M Ligand + 5 mL, 0.005 M MCI2 + 15.5 water.

The titration curve of solution A, B and C with alkali is called

acid titration, ligand titration and metal-ligand titration curves,

respectively. A plot between volume of alkali added and the

corresponding pH values was made for each set. The complex formation

results in the lowering of buffer region due to proton displacement.

Calculation of n (Proton Ligand Formation Number):

The acid, ligand and metal-ligand titration curves were obtained by

plotting pH values of solutions A, B and C against volume of NaOH

added, respectively. The ligand curve deviates from pure acid curve at

pH-2.5, which suggests that deprotonation starts from this pH. The value

of proton ligand formation number has been evaluated by the pointwise

calculation method using the following equation:

13

^, ( V " - V ) ( N ° + E ° )

( V ° + V ' ) T °

where Y is the average number of dissociable protons present in the

ligand. V and V" are the volume of NaOH consumed during acid

and ligand titration respectively. V° is the total volume of the

mixture. N° is the concentration of NaOH. E° is the concentration

of acid. T°L is the initial concentration of Ligand.

The proton ligand formation curves were obtained by plotting

proton ligand formation number (HA) against pH. The approximate

value of the proton ligand stability constants, LogK"and LogK" were

evaluated by pointwise calculation using the following equations:

LogK," = p H + L o g , „ - % - (n, <1)

LogK," = pH+ Log,„ ^ ^ (1< n, < 2)

Calculation of n (Metal ligand formation number):

Metal ligand formation number is the average number of

ligands attached per metal ion and was evaluated by using

following equation:

_ ^ ( V " - V " ) ( N ° + E < ' )

( V + V j i l . T "

where T°M is the initial concentration of metal ion. V" and V" are

14

the volumes of NaOH consumed during ligand and complex

titration respectively. The values of n were determined at various

pH by pointwise calculation method.

Calculation of pL (Free ligand exponent):

The free ligand exponent was evaluated at various n values

using the following equation:

pL = Log 10 i = 0

>"

anti log pH^ rv°+v"

Where |3"n is the overall proton ligand stability constant. The

pL values were also calculated at different pH by pointwise

calculation method.

Calculation of Log K (Stability Constant);

The formation curves of the metal ligand complexes were

obtained by plotting ii versus pL and stability constant can be

directly estimated from them by half-integral method [34,35]. The

stability constant was also evaluated by using following equations:

n LogK, = pL + Log— -1-n

LogKj =pL + Log—— 2 - n

15

RESULTS AND DISCUSSION

A. pH metric titration of Sulfamethazine,

In the case of sulfamethazine, proton ligand formation

number (w A) shows that the molecule contains only one replaceable

hydrogen atom.

H2 \ /

S O 2 — N H — ^

Sulfamethazine

The proton ligand formation number ( « A ) was obtained by

potentiometric titration curves of acid and ligand at different pH

(Fig. 1-5) by using pointwise-calculation method. Since « A is less

than one (Table-1) there is only one replaceable hydrogen atom and

therefore only LogK" is obtained.

16

12

10

« 6

-Connplex Titration witliMn(ll)

I I I I 1 I I I I I I I I I I I I — I — \ — I — I — I — I — \ — \ — I — I — I — I — I — I — I — r

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

Volume of NaOH (ml)

Fjg.1 Potentiometric Titration Curves of Sulfamethazine with IVIn(ll)

17

12

10

0)

« 6 X

a

- Ligand Titration

-)<—Complex Titration withFe(ll)

T I I I t I 1 1 ; I I I 1 I I I I I 1 1 1 1 1 — I 1 — I — 1 \ 1 1 1 — \ r

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

Volume of NaOH (ml)

Fig.2 Potentiometric Titration Curves of Sulfamethazine with Fe(ll)

18

12

10

I I I I — I — I — I — I — I \ I I I — I — I — I — I — I — I — I — I — 1 — I — t — I — I — I — I — I — i — I — I — I — I

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

V o l u m e of NaOH (ml)

Fig.3 Potent iometr ic Titration Curves of Sui famethazlne

w i t h C o ( l l )

19

12

10

I I \ t I ] 1 I I i I 1 I I I I I I I 1 1 1 — i — I — \ — I — I — 1 — 1 — ] — 1 — \ — 1 — I

0 O.S 1 1.S 2 2.S 3 3.5 4 4.5 5 S.S 6 6.5 7 7.5 8

Volume o f N a O H (ml)

Fig.4 Potentiometrjc Titration Curves of Sulfamethazine

with Ni(il)

20

12 -,

10

-[—I—I—I—I—I—r T ! I i I — I — I — I — \ — r 1 I I I I I I

0 0,5 I 1.5 2 2.5 3 3,5 4 4.5 5 5.5 6 6.5 7 7.5

Volume o f N a O H j m l )

Fig.5 PotentJometric Titration Curves of Suifamethazine

wi thCu( l l )

21

Table-1

Proton Ligand Formation Number («A)Of Sulfamethazine

(SMZ) at different pH.

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

n A(sulfamethazine)

0.94

0.80

0.72

0.66

0.58

0.56

0.54

0.52

0.48

0.42

0.38

0.32

0.29

0.21

The protonation constants (Table 2) of sulfamethazine LogK," and

LogK," were directly obtained by plotting n^ versus pH value. The pH at

fi, =1.5and iT =0.5 refers to LogK," and LogK * respectively. This is

known as Bjerrum^s half- integral method.

22

Table-2

Protonation constants of sulfamethazine

Half-integral method

Pointwise calculation method

LogK"

7.00

7.02

LogPi:

7.00

7.02

Proton ligand formation curve (Fig, 6) were obtained by plotting

n^ versus pH of the system. The protonation constant can be directly

obtained by using Bjerrum's half-integral method [38, 39].

1.0-,

0.9

I 0.8-E

0.7-

0.6-

•O 0.5-1 c 3> = 0.4 H c o 1 0.3-Q.

0.2.

- | ' 1 ' 1 9 10 11

Fig.6 Proton ligand fonnation curve of Sulfomethazine at u = 0.1 M

23

The metal ligand formation number (n) was calculated by

potentiometric titration curves of ligand and complex from the n ^

value and protonation constant (^^8*^2) by pointwise calculation

method at different pH values.

Since metal ligand formation number (ii) is less than one,it is

presumed that the complex is formed in 1:1 (M:L) ratio, except in

the case of Co(II) and Ni(II) complexes where sulfamethazine to

metal ratio is 1:2 [40]. The complex forming tendency increases

on moving from manganese through copper which is indicated by

metal-ligand formation number (n).

The free ligand concentration (pL) of sulfamethazine was

obtained by the pointwise-calculation method at different pH

values with the help of ii value and, protonation constant (LogK^^

The values of proton ligand formation number (n) and free ligand

concentration (pL) are given intable-3.

24

Table-3 Metal Ligand Formation Number and Free Ligand-

Concentration (pL) of Sulfamethazine.

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

Mn

n

0.16

0.18

0.20

0.26

0.30

0.37

0.42

0.49

0.52

0.58

0.62

0.64

0.66

0.69

pL

8.68

8.19

7.72

7.21

6.25

6.06

5.19

4.32

3.81

3.50

3.10

2.88

2.62

2.40

Fe

n

0.17

0.19

0.21

0.30

0.33

0.37

0.44

0.5

0.55

0.6

0.64

0.67

0.70

0.72

pL

8.81

8.35

7.86

7.39

6.88

6.35

5.60

5.12

4.60

4.21

3.80

3.36

2.80

2.43

Co

n

0.18

0.22

0.29

0.35

0.46

0.61

0.71

0.89

1.05

1.32

1.40

1.56

1.70

1.88

pL

8.84

8.42

7.90

7.50

6.92

6.45

5.60

5.13

4.56

4.20

3.82

3.41

2.90

2.56

Ni

n

0.21

0.30

0.34

0.41

0.56

0.78

0.88

1.02

1.24

1.47

1.61

1.72

1.83

1.95

pL

8.92

8.51

7.95

7.56

6.96

6.48

5.66

5.19

4.39

3.98

3.53

3.05

2.70

2.31

Cu

n

0.39

0.46

0.52

0.59

0.61

0.66

0.70

0.75

0.79

0.82

0.85

0.89

0.92

0.96

pL

9.12

8.61

8.11

7.68

7.19

6.71

6.30

5.71

5.28

4.75

4.29

3.76

3.25

2.71

Since metal ligand formation number (n) increases with

increasing pH, it suggests that complex formation is favourable in

basic medium.

The metal-ligand formation curves (Fig.7) were obtained by

plotting n versus pL values [41,42]. The stability constants of the

sulfamethazine-complexes can also be directly obtained by half-

25

integral method (Table 4) in which n=0.5 for Log Ki and n=1.5 for

Log Ki

2 . 0 -

1 .8 -

1 .6 -

I J I 10-

i 0.8

0 . 6 -

0.4 -

0 . 2 -

0.0

• . ^ >

-•-Mn(l l ) • Fe(ll) A Co(ll) • Ni(ll) • Cu(ll)

—I 10 2 3 4 5 6 7 8 9

Free ligand concentration (pL)

Fig.7 Formation Curves of Sulfamethazine witfi different metal ions at ^ = 0.1 M

Table 4

Stability constant of sulfamethazine with divalent metal ions

Metal ions

Mn

Fe

Co

Ni

Cu

Stability constant (Log K)

Pointwise calculation method

4.01

4.07

4.14

4.20

4.28

Half integral Method

4.90

5.38

5.80

7.30

7.52

26

The stability constants (Log Ki and Log K2) were also

calaulated by pointwise calculation method (Table 5) at different pH

values. The stability constant decreases gradually with increasing pH.

Table 5

Stability constants of sulfamethazine (Log K) by pointwise

calculation

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10

Mn

6.61

6.12

5.76

5.40

4.96

4.58

4.12

3.80

3.39

3.02

2.68

2.31

1.92

1.61

Fe

6.72

6.25

5.84

5.46

5.01

4.64

4.20

3.85

3.45

3.10

2.71

2.36

1.94

1.65

Co

6.81

6.30

5.90

5.51

5.05

4.69

4.36

3.92

3.52

3.15

2.78

2.40

1.97

1.68

Ni

6.92

6.41

5.95

5.62

5.12

4.73

4.41

3.96

3.62

3.18

2.84

2.42

2.02

1.75

Cu

6.98

6.51

6.04

5.68

5.18

4.82

4.48

4.04

3.76

3.25

2.91

2.52

2.06

1.79

27

The stability constant increases gradually from Mn through

Copper. It follows the Irving -Wil l iam's order [43].

Mn < Fe < Co < Ni < Cu

A plot of LogK of the sulfamethazine-complexes vs e /r of

metal ions (Table 6) showed a straight line (Fig.8) indicating that the

metal-ligand(M-L) interaction is predominantly electrostatic in nature

and results in the formation of ionic complexes although the

possibility of some covalent interaction cannot be excluded [44,45].

Table 6

Log K of sulfamethazine and e /r value of divalent metal ions

Mn"^

Fe^^

Co^^

Ni ^

Cu^^

LogK

4.01

4.07

4.14

4.20

4.28

e ' / r

4.12

4.35

4.52

4.82

4.60

4.3

4.25

4.2-j

^ 4.15

o 4.1

4.05-

4

3.95

4 4.1 4.2 4.3 4.4 4.5 4.6 4.7 4.8 4.9

e'/r

Fig.8 A plot of e /r verus Log K of Sulfamethazine

28

B. Conductometric Titration

In the conductivity titrations, 20ml of sulfamethazine(lxlO'^M)

was titrated against different metal chloride (IxlO^^M) solutions in

aqueous medium. It was found that a 2:1 (M.L) complexes were formed

with Co(Il) and Ni(Il) although 1:1(M:L) complexes are obtained with

other metal ions. Precipitation of the complexes does not occur under the

experimental conditions at room temperature.

14

13

12

? i i

5 10 c 3 « 9 •o § 8 O

7

6

5 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70

Volume of Metal ions (ml)

Fig. 9 ConduGtometriG Titration Curves of Sulfamethazine with different metal ions

The thermodynamic parameters (Table 7) were also calculated

from the stability constants by using the following formulas [46].

AG=-2.303RT Log K

AG=AH-TAS

29

With increasing stability constant of complexes, free energy

(AG) decreases gradually from manganese to copper. It suggests

that the reactions are spontaneous i.e. irreversible. As free energy is

inversely proportional to the temperature, the stability constant

decreases with increasing temperature of the system.

Enthalpy changes (AH) accompanying the complexation of metal

ions in solution is characteristic property of reaction. The enthalpy of

reaction is negative and decreases from Mn to copper with increasing

stability constant (Log K).It suggests that the reactions are exothermic.

There will be an increase in the number of particles in the system

leading to an mcrease in the entropy (AS) of the system. Since entropy is

a measure of the extent of disordemess of the system, more positive value

of entropy was observed in all the systems studied.

Table-7

Thermodynamic parameters for sulfamethazine

Metal Mn Fe Co Ni Cu

Log K 4.01 4.07

4.14 4.20 4.28

-AG (KCal) 5.47 5.55

5.65 5.73 5.84

-AH (Cal 18.35 ^ 18.62

18.94 19.22 19.59

AS (Cal/degree) 18.28 18.56

18.88

19.15 19.51

30

A. pH metric titration of Pyridoxal 5-phosphate

The "A value for pyridoxal 5-phosphate is more than one

which suggests that there are two replaceable hydrogen atoms in

the molecule.

CHO OH, 1 Oil

\ / ^

CH3 ^ cf ^OH

(Pyridoxal 5-phosphate)

The pyridoxal 5-phosphate gives a pale yellow-colour in aqueous

solution. Since precipitation occurs in basic medium the experiment was

performed below that pH. It was noted that the addition of alkali was

accompanied by a gradual change in the colour of pyridoxal 5-phosphate

from yellow to deep yellow with increasing pH. The colour change

(Table 8) observed in presence of metal ions are almost similar to those

observed in their absence,[47,48].

Table-8 Colour change in PLP-metal system

pH

3.0-5.85

5.85-10.0

10.0 above

Colour

Pale yellow- deep

Reddish

Precipitation occurs

31

The proton ligand formation number (« A) was calculated by

pointwise calculation method at different pH values using the acid

and ligand titration curves (Fig. 10-14). In this experiment, the w ^

value is greater than one (Table 9) and therefore, protonation

constants, LogK" and LogK" were obtained by pointwise

calculation method.

Table-9 Proton Ligand Formation Number of Pyridoxal 5-phosphate (PL?)

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

« A Pyridoxal 5-phosphate (PLP)

1.71

1.58

1.43

1.36

1.26

1.15

1.06

0.92

0.88

0.80

0.71

0.58

0.49

0.41

32

12

10

??^»4^

.6 1 1,S 2 2.S 3 3,6 4 4.S i (.6 G

V o l u n i e o f N a O H ( m l )

(.6 7 7.5

Fig.10 Potent iometr ic Titration Curves of Pyr idoxalS-phosphate with Mn(l l )

33

12

10

- ^ f - C o m plex Titralion with Fel

0.6 1 1.< 2 2.6 3 3.6 4 4.6 6 6.6

V o l u m e of NaOH |m I)

6.6 7 7.6

F i 9 . l l Potent iometr ic Titration C u r v e s of PyridoxalS>pl\ospl \ate with Fe(ll)

34

12

10

-T \ n

- C o m p lex T i t ra t ion wi th Co ( l l )

T I I I [ r~n I I 1 r i 1—~i 1—i 1 r—~r

0 O.S 1 1.5 2 2.6 3 3.S 4 4.5 5 5.S .6 7 7.6

V o l u m e Of N a O H (m I j

Fig.12 Potent iometr ic Titrat ion C u r v e s of Pyr idoxai S-pl iosphate

with Co(l i )

35

12

10

"1 ] r — I [- -1 1 n "1 1 1 r-

0 0 .5 1 1.S 2 2.S 3 3.S 4 4 . 5 5 5 .5

Volume of NaOH (ml)

.5 7 7.5

Fig.13 Po ten tic me trie T i t ra t ion C u r v e s of Pyr idoxa l 5 - p h o s p h a t e

with Ni(l l)

36

12

10

— I r—\ r"~i r—1 i \ i — T I I I I I 1 1 1 1 1 1 1 1 1 1 1 \ ] 1 \ 1 —

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

Volum« of NaOH (ml)

F ig .14 P o t e n t i o m e t r i c T i t r a t i o n C u r v e s of P y r i d o x a l 5 - p h o s p h a t e

with C u d ! )

37

The protonation constants (Table 10) of pyridoxal 5-phosphate,

LogK" and LogK" were directly obtained by plotting n^ versus pH value

(Fig. 15) by noting the pH at which n^ = 1.5 and n^ = 0.5 respectively.

Table 10

Protonation constants of Pyridoxal 5-phosphate.

Half integral method

Point wise calculation method

LogKf

4.52

4.68

LogK;"

9.35

8.96

Logpr

13.87

13.64

1.8

1.6

I.., c .2 12 •s

, .

0.8

I 0.6-

0.4-

- 1 1 1 1 1 1 1 1 1 . 1 1 1 1 , 1 1

3 4 5 6 7 8 9 10 11

pH

Fig. 15 Proton ligand formation curve of Pyridoxal 5-phosphate

Since metal ligand formation number (n) increases (Table 11)

with increasing pH, the complex formation is favoured in basic

medium. The metal-ligand formation number indicates that there is

38

a weak interaction of divalent metal ions of first transition series

with pyridoxal 5-phosphate.

Table-11

Metal Ligand Formation Number and Free Ligand-

Concentration (pL) of Pyridoxal 5-phospliate

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

Mn

n

0.12

0.14

0.22

0.27

0.31

0.37

0.41

0.46

0.49

0.52

0.55

0.58

0.61

0.63

pL

12.44

11.92

11.45

10.96

10.48

9.96

9.51

8.97

8.50

8.03

7.59

7.10

6.61

6.18

Fe

n

0.16

0.18

0.23

0.29

0.33

0.40

0.43

0.48

0.50

0.53

0.57

0.60

0.63

0.66

pL

12.45

11.95

11.48

10.84

10.46

9.98

9.50

8.96

8.50

7.98

7.48

7.01

6.58

6.10

Co

n

0.17

0.19

0.26

0.34

0.36

0.40

0.44

0.51

0.59

0.65

0.71

0.73

0.77

0.80

pL

12.50

12.06

11.62

11.15

10.61

10.20

Q.68

9.21

8.62

8.24

7.77

7.26

6.80

6.38

]

n

0.20

0.24

0.28

0.36

0.41

0.43

0.46

0.53

0.60

0.68

0.72

0.75

0.79

0.83

Ni

pL

12.60

12.10

11.75

11.20

10.63

10.26

9.72

9.30

8.71

8.30

7.80

7.31

6.84

6.41

Cu

n

0.28

0.32

0.37

0.41

0.44

0.49

0.52

0.59

0.62

0.70

0.73

0.77

0.80

0.86

pL

12.65

12.15

11.78

11.29

10.69

10.28

9.75

9.32

8.76

8.36

7.88

7.35

6.85

6.50

The metal-ligand formation curves (Fig. 16) were obtained by plotting n

versus pL values. The stability constants of the PLP-complexes can also be

39

directly obtained from half- integral method (Table 12) in which n=0.5 for

Log Ki and n=1.5 for Log Ki.

0.9 •

0.8-

S 0.7 •

E n 0.6-c p

I 0.5.

E 0.4. D)

I 0.3 •

0.2-

0.1

—•-- • -- A -

—T-•

-Mn(ll) Fe{ll)

-Co(ll) -Ni(ll)

Cu(ll)

—r-10

—r 11 12 13

Free ligand concentration (pL)

Fig. 16 Metal ligand formation curves of PLP at ^ = 0.1 M

Table 12

Stability constant of Pyridoxal 5-phosphate with divalent metal ions

Metal ions

Mn

Fe

Co

Ni

Cu

Stability cons

Pointwise calculation method

7.91

8.03

8.15

8.45

8.58

tant (Log K)

Half integral Method

8.35

8.50

9.30

9.42

9.86

40

The stability constants (Log K ) of pyridoxal 5-phosphate were

calaulated by the pointwise calculation method (Table 13) at

different pH. It was observed that the stability constant decreases

gradually with increasing pH of the system.

Table-13 Stability constants of pyridoxal 5-phosphate (Log K) by

Pointwise Calculation Method

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

Mn

10.79

10.18

9.87

9.38

8.92

8.47

8.21

7.56

7.28

6.64

6.41

5.96

5.67

5.44

Fe

10.81

10.21

9.88

9.40

8.98

8.52

8.26

7.60

7.36

6.78

6.60

^ 6.12

5.76

5.50

Co

10.83

10.22

9.90

9.50

9.08

8.60

8.55

8.06

7.79

7.25

6.67

6.20

6.16

5.77

Ni

10.86

10.24

9.94

9.68

9.15

8.91

8.46

8.17

7.85

7.42

6.83

6.50

6.22

5.87

Cu

10.99

10.44

10.06

9.78

9.26

9.05

8.76

8.36

8.00

7.53

6.96

6.56

6.30

5.98

The stability constant increases gradually from Mn through

Cu according to the Irving -William^ s order .

Mn < Fe <Co < Ni < Cu

41

A plot of logK of the PLP-complexes vs e^/r of different

metal ions (Table 14) shows a straight line (Fig. 17) indicating that

the metal-ligand (M:L)interaction is predominantly electrostatic in

nature and results in the formation of ionic complexes although the

possibility of some covalent interaction cannot be excluded [49].

Table 14

Log K of Pyridoxal 5-phosphate and e /r value of divalent metal ions

Mn"'

Fe ^

Co ^

Ni"^

Cn''

LogK

7.91

8.03

8.15

8.45

8.58

eVr

4.12

4.35

4.52

4.82

4.60

8.6 n

8 .4 -i^

— 1

8 .

78-1 . U

i

-

1

\ 42

• ^ ^^—•

^

^ 4

1 - 1 1 • - — 1

4.4 46 48 I e^/r

Fig.lTAplotofe^/rvslogKofPLP

42

The plot suggests interaction of electrostatic type between

the metal ion and the pyridoxal 5-phosphate.A 1:1 (M:L) complex

is formed between PLP and transition metal ions.

It has been shown by Viswanathan and coworkers from NMR

spectral data that Mn(II) and Co(II) complexes are simultaneously

bound to the phosphate group and aldehyde-oxygen in the PLP-

complex system and stability constants of complexes indicate a

definite interaction with ring[50,51].

C H 3 ^ r ^

where M = Mn (II), Co (II)

B. Conductivity Titration of Pyridoxal 5-phosphate:

In the conductivity measurements, 10ml of PLP(lxlO~'*M) was

titrated (Fig, 18) against different metal chloride (lxlO'"*M) solutions

in aqueous medium. It was found that a 1:1(M:L) complexes were

formed with all the metal ions without any precipitation under the

experimental conditions at room temperature.

43

55-,

50 -

•^ 4 5 -

a> o c (0

ts = 40 c o O

35-

30

- • -Cu( l l ) - • - Co(ll) - A - Nb((()

— I — 10

—T—

20 —r-30

I 40

Volume of metal ions (ml)

Fig. 18 Conductometiic Titration curves of PLP-metai system

As free energy (Table 15) decreases gradually from manganese

to copper with increasing the stability constant,it suggests that the

reactions are spontaneous.

There is a decrease in the enthalpy of reaction from Mn to

copper suggesting that reactions are exothermic. As entropy is the

measure of randomness and it increases gradually from Mn to

copper suggesting that complex formation tendency also increases

from Mn to copper.

44

TabIe-15

Thermodynamic Parameters of Pyridoxal 5-Phosphate.

Mn

Fe

Co

Ni

Cu

LogK

7.91

8.04

8.15

8.26

8.39

-AG

(KCal)

10.78

10.88

11.11

11.26

11.44

-AH

(Cal)

36.20

36.52

37.30

37.80

38.39

AS (Cal/degree/mole)

36.08

36.14

37.17

37.67

38.26

45

A. pH metric titration of DL-Cystine

The w A value for DL-cystine is greater than one (Table 16) which

suggests that there are two replaceable hydrogen atoms in the molecule.

COOJ ' \

COOH /

CH-CH2-S-S-CH2-CH Mi{ NH2

(Cystine)

In acidic medium, the amino acids exist as HsN' RCOOH.The pH of

initial solutions of amino acids is higher than the free acid because

TabIe-16 Proton Ligand Formation Number of DL-Cystine

pH

3.5

4.0

4.5

5.0

5.5

#.0

• ^ 1 - 1 ^ ^ . 5

/ \ ^ . 0

WA(Cystine)

1.89

1.72

1.61

1.57

1.45

1.28

1.13

0.95

0.88

0.71

0.58

0.42

0.36

^ 0.29

A. pH metric titration of DL-Cystine

The « A value for DL-cystine is greater than one (Table 16) which

suggests that there are two replaceable hydrogen atoms in the molecule.

COO^ ^ /

,CH-CH2-S-S-CH2-CI^

COOH

NH-/ NHo

(Cystine)

In acidic medium, the amino acids exist as HsN^RCOOH.The pH of

initial solutions of amino acids is higher than the free acid because

TabIe-16 Proton Ligand Formation Number of DL-Cystine

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

/7A(Cystine)

1.89

1.72

1.61

1.57

1.45

1.28

1.13

0.95

0.88

0.71

0.58

0.42

0.36

0.29

46

of the protonation of the basic -COO group and the dipolar ion has

been converted to the acid form, HsN^RCOOH.

The colourless complexes of DL-cystine formed in solution in

acidic medium, and readily gives precipitate in alkaline medium.Due to

very low concentration of metal ions used, the formation of polynuclear

species was considered to be unimportant[52,53].

The proton ligand formation number ( « A ) was calculated by

pointwise calculation method at different pH values using the acid

and ligand titration curves (Fig. 19-23).

The protonation constants (Table 17) of DL-cystine, LogKf

and LogK" were directly obtained from a plot of n^ versus pH (Fig. 24)

the value corresponds to n^ = 1.5 and n^ = 0.5 respectively.

TabIe-17

Protonation constants of DL-Cystine

Half integral method

Pointwise calculation method

LogKj^

4.80

4.86

LogK^

8.78

8.24

LogPi:

13.58

13.10

47

12

10

8

o 3 7i > 6 z a.

- Complex Trtration w jth lUh(n)

~T 1 ^ I I I I i I I I I I i I 1 1 1 1 1 1 1 1 i 1 \ 1 1 1 1 1 1 1

0 0.6 1 1.6 2 2.6 3 3.6 4 4.6 6 6.6 6 6.6 7 7.6 8

Volume of NaOH (ml)

Fig.19 Potentiometric Titration Curves of DL-Cystine with Mn(ll)

48

12 n

10

»=#=•=»

I.S 1.25 2 2.75 3.5 4.25 5 5.75

V o l u m e o f N a O H (ml)

6.5 7 ,25

F i g . 2 0 P o t e n i i o m e t r i c T i t r a t i o n C u r v e s o f D L - C y s t i n e

w i t h F e ( l i )

49

12

10

3

— [ — ( — ~ [ — 1 1 1 1 ( — I 1 1 [ — I 1 — I — I — I 1 — I 1 1 ; — \ — T — ; — ; i — ; — i ; i — \ — i

0 0.6 1 1.S 2 2.6 3 3 .6 4 4.6 5 5.6 6 6,6 7 7.6 8

V o l u m e of NaOH (ml)

Fig .21 P o t e n t j o m e t r i c T lira tie n C u r v e s of D L - C y s t i n e with C o ( l l )

50

12

10

~\ 1 1 i 1 r- n r '1 r~-i r~T~i—i—n *n—1—I—r-

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

Volume of NaOH (ml)

Fig.22 Potentlometrjc Titration Curves of DL-Cystine

with Nj/ll)

51

12

10

"T""' J'" "I r-^r—r

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8

Volumt of NiOH (mil

Fig.23 P o t e n t i o m e t r i c T i t ra t ion C u r v e s of DL-Cyst ine

with Cu( l l )

52

E 3 C C o

•S3

n

• D c (0

c S 2

2.0-1

1.8

1.6

1.4-

1.2-

1.0-

0.8-

0.6-

0.4-

0.2-

- i — I — r -»—r 6 7 8

pH value 10 11

Fig. 24 Proton Ligand Formation Curve of OL-Cystine

The metal-ligand formation number (Table 18) indicates that

there is a weak interaction of divalent metal ions of first transition

series with DL-cystine. No precipitate was formed during the

titration of metals with DL-cystine upto pH 8.0.

The metal ligand formation number (« ) increases with

increasing pH which suggests that complex formation is favoured in

basic medium. The free ligand concentration (pL) decreases with

the increasing pH.

Table-18

Metal Ligand Formation Number and Free Ligand

Concentration of DL-Cystine.

PH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

Mn

n

0.09

0.1

0.12

0.16

0.2

0.28

0.33

0.39

0.42

0.44

0.49

0.52

0.55

0.58

pL

12.88

12.4

11.9

11.39

10.9

10.4

9.91

9.38

8.95

8.43

7.96

7.44

6.92

6.51

Fe

n

0.062

0.1

0.13

0.21

0.32

0.35

0.4

0.43

0.47

0.49

0.51

0.53

0.58

0.61

pL

12.9

12.4

11.92

11.44

10.93

10.45

9.96

9.45

8.96

8.46

7.98

7.51

6.98

6.55

Co

n

0.1

0.12

0.15

0.23

0.33

0.37

0.41

0.44

0.48

0.52

0.55

0.58

0.61

0.64

pL

12.93

12.4

11.95

11.43

10.96

10.45

9.99

9.5

9.03

8.58

8.1

7.52

7.18

6.57

Ni

n

0.15

0.16

0.19

0.28

0.39

0.41

0.46

0.5

0.53

0.58

0.61

0.64

0.69

0.72

pL

12.96

12.5

11.96

11.46

11.02

10.51

10.02

9.42

9.02

8.46

8.06

7.5

7.12

6.6

Cu

n

0.21

0.25

0.29

0.32

0.41

0.45

0.49

0.52

0.55

0.6

0.65

0.69

0.72

0.78

pL

12.98

12.52

12.08

11.55

11.05

10.55

10.1

9.58

9.11

8.61

8.26

7.78

7.41

6.92

From a plot (Fig. 25) of « and pL, the formation curves of

metal-complexes were obtained from which the stability constants

(Table 19) can be obtained by half-integral method.

54

0.8 •

0.7 •

5 0.6 E 3 c 0.5 g TO

0.4-

I 0.3

0) 0.2

0.1

0.0.

'''^ l^n^ytx^'^

X —•-- • -- A -

— • -

- Mn(ll) - Fe(ll) -Co(ll) - Ni(ll) -Cu(ll)

- I — ' — I — ^ 11 12 6 7 8 9 10 11 12 13

Free ligand concentration (pL)

Fig. 25 Metal ligand formation curves of DL-cystine

Table 19

Stability constant of DL-Cystine with divalent metal ions

Metal ions

Mn

Fe

Co

Ni

Cu

Stability constant (Log K)

Pointwise calculation method

8.75

8.88

9.24

9.34

9.43

Half integral Method

7.80

8.22

8.82

9.53

9.95

55

As the pH value increases, the stability constant decreases

(Table 20) gradually. The stability constants were also calculated

by pointwise calculation method at different pH.

Table-20

Stability Constants of DL-Cystine (Log K) by Pointwise

Calculation Method

pH

3.5

4.0

4.5

5.0

5.5

6.0

6.5

7.0

7.5

8.0

8.5

9.0

9.5

10.0

Mn

11.65

11.17

10.79

10.27

9.90

9.40

8.95

8.41

8.06

7.71

7.24

6.71

6.19

6.06

Fe

11.76

11.36

10.88

10.45

9.96

9.56

8.98

8.58

8.12

7.84

7.44

6.84

6.32

6.16

Co

11.96

11.53

11.14

10.71

10.33

9.82

3.49

9.00

8.60

8.13

7.73

7.32

7.06

6.44

Ni

12.03

11.65

11.26

10.84

10.47

10.02

9.60

9.18

8.76

8.30

7.80

7.36

7.08

6.48

Cu

12.26

12.50

11.62

11.11

10.62

10.14

9.65

9.25

8.84

8.35

7.92

7.38

7.20

6.56

The stability constant increases gradually from Mn through

Copper which follows Irving -William's order:-

Mn < Fe<Co < Ni < Cu

56

A plot of logK of the Cystine-complexes vs e /r of metal ions

(Table 21) shows a straight line (Fig.26) indicating that the metal-

ligand interaction is predominantly electrostatic in nature and

results in the formation of ionic complexes although the possibility

of some covalent interaction cannot be excluded.

Table 21 Log K of DL-cystine and eVr value of divalent metal ions

Mn"2

Fe^^

Co^^

Ni"'

Cu^

LogK

8.78

8.95

9.25

9.36

9.47

eVr

4.12

4.35

4.52

4.82

4.60

9.6 9.5 9.4 9.3

? 9.1 9

8.9 8.8 8.7 — , J —

4.1 4.2 4.3 4.4 4.5 4.6 4.7 4.8 4.9 e»/r

Fig. 26 A plot of e /r vs log K of DL-cystine

The free energy (AG) decreases gradually from manganese to

copper with increasing stability constant of complexes. It suggests

57

that the reactions are spontaneous i.e. irreversible. As free energy is

inversely proportional to the temperature, the stability constant

decreases with increasing temperature of the system.

The enthalpy of reaction is negative and decrease from Mn to

copper with increasing stability constant (Log K).It suggests that

the reactions are exothermic. Enthalpy changes (AH) accompanying

the complexation of metal ions in solution is characteristic property

of reaction

There will be an increase in the entropy (AS) of the system

during experiment. Since entropy is a measure of the extent of

disordemess of the system, more positive value of entropy was

observed in all the system studied.

Table-22

Thermodynamic parameters for DL-cystine

Metal

Mn

Fe

Co

Ni

Cu

LogK

8.75

8.88

9.24

9.34

9.43

-AG

(KCal)

11.93

12.24

12.60

12.71

12.86

-AH

(Cal)

40.04

40.45

42.28

42.65

43.15

AS

(Cal/degree/mole)

39.91

40.26

42.14

42.50

43.00

58

Stability of sulfamethazine, pyriodoxal 5-phosphate and

DL-cystine upon complextion with divalent metal ions decreases with

an increase in pH value.

The values of Gibbs free energy change is negative which

indicates that the reaction is spontaneous.The enthalpy change is

also negative which suggests that the reactions are exothermic[56-

58]. As entropy is positive so it also favours the complexation

reactions.

59

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