UGC Minor Research Project

SUMMARY OF THE PROJECT

Title

“pH metric and thermodynamic studies of binary complexes of

Co(II),Rh(III),Pd(II),Pt(II),Ag(I),Zn(II),Cd(II)with Ibuprofen and

paracetamol ”

For the year- 2014-2016

File No:47-671/13(WRO) dated 20/05/2014

Name of Principal Investigator

Mr. G.D.Rawate

Assistant Professor

Department of Chemistry

Shri R. R. Lahoti Science College, Morshi

NAAC Reacreditated ‘B’ grade

Appendix C

Summary of Project

“pH metric and thermodynamic studies of binary complexes of

Co(II),Rh(III),Pd(II),Pt(II),Ag(I),Zn(II),Cd(II)with Ibuprofen and

paracetamol ”

CONTENTS

S.No. Title Page No.

1 INTRODUCTION 1-3

2 OBJECTIVE 4-14

3 METHODS OF DETERMINATION OF STABILITY CONSTANTS

15-25

4 EXPERIMENTAL AND DISCUSSION 26-32

5 CONCLUSION

32

6 REFERENCES

33-36

INTRODUCTION:-

The binary complex of Co(II), with medicinally important compound

paracetamol and ibuprofen have been described. The importance and physiological

characteristic of the pharmaceutical compounds use as ligands in the study of binary

complexes. The potentiometric study of binary complexes of most of these ligands with

chosen divalent transition metal ion have been reported for the first time in these project.

The metal ion chosen are the 3d series elements in their dipositive oxidation state. Fe(II)

was not chosen as it is more susceptible to the hydrolysis and readily get converted into

tri positive oxidation state. Metal ion used was Co(II). this element are the essential

elements in human physiology.

A brief survey of the study of binary complexes of these metal ions with

the selected pharmaceutical compounds is described below.

1

Interaction of pyridoxine with different metal ions has been studied using

polarographic method 1-4.Interaction of ehtanbutol with transition metal ions in solution

has been studied spectrophotometrically and potentiometrically by Bhattacharyya et al.5,6.

The formation constants of the complexes of Cu(II) 7and Ni(II) 8with

levodopa (LDP) have been investigated. The stability constants of heterobinuclear

complexes formed by LDP dopamine, histidine with Cu(II)-Ni(II), Cu(II)-Zn(II) and

penicillamine with Ni(II)-Zn(II) have been evaluated pH metrically by Nair et al9-11.

some of the chelates of penicillamine with different metal ion have been studied 12-14.

Nair and neelakantan 15have studied the complexes of Ni(II)with 6-aminopenicillanic

acid, ampicilin, L-cysteine and penicillamine.

Combined pH -metric and spetrophotometric study on the complex

formation of Cu(II) with ampicilin in aquous medium was reported16. Equilibrium study

of complexes of some metal ion with ampicillin has been reported17-21. Chakravarti et

al22. have studied the formation of biological chelates have divalent metal ion with

cephalosporins pH- metrically.

Bisht et al 23studied the interaction of Cu(II),Ni(II),Zn(II), and Cd(II) with

HQ anthranilic acid pH-metrically. The stability constants of some 3D divergent

transition metals with anthranilic and nicotinic acid have been repotorted pH metrically24.

Keemti Lal 25has studies the stability constants of Mn(II), Co(II), Ni(II),

Cu(II), and Zn(II) complexes of some sulpha drugs in 50% (v/v) ethanol- water medium

potentiometrically. Formation constant of Mn(II), Co(II), Fe(II) and Zn(II) with nalidixic

2

acid (an antibiotic) have been studied pH-metrically by Sharma et al26 compexes of

Mn(II), Fe(III), Co(II), Ni(II), Cu(II), Zn(II), Cd(II) and Hg(II) with

monophenylbutazone 27and oxyphenylbutazone28 have been studied . Sharma and

Joseph29 have investigated the interaction of Ca(II), Mg(II), Co(II), Fe(II), Ni(II) and

Cu(II) with tetracycline (broad spectrum antibiotic) pH-metrically.

In this project, the detailed investigation regarding the determination of

practical proton-ligand stability constants of paracetamol and ibuprofen reported. The

stability constants of binary complexes of paracetamol and ibuprofen with Co(II)

,Zn(II),Cd(II) were reported.

Generally complexes are designated as stable or unstable. The general meaning of

stability is supposed to be related with the concept, whether a particular complex can be

converted into other easily or not. As a matter of fact, this is kinetic aspect of stability;

which deals with the rate of the reaction and its mechanism. The other aspect of stability

is thermodynamic aspect. In which stability of a complex is related with the amount of

energy released during its formation or the amount of energy required to break it.

In this unit we describe complex forming equilibria in solution and the various

factors affecting it. We will also discuss the various factors affecting stability constants

for the formation of complexes in solution. In the end of the unit we shall describe the

method used for determining stability constants of the complexes formed in solution.

Which involves quantitative characterisation of the complex-forming reaction in solution.

You may recall what you have already studied about the basic concept of

chemical equilibria in solution.

3

OBJECTIVE

The main aim of this project is to study the complex formation equilibria in

solution. After going through this unit you should be able to:

describe stepwise and overall formation constants;

explain thermodynamic importance of stability constants;

discuss factors affecting stability of complexes; and

describe methods of determining stability constants for binary complexes in

solution.

STEP-WISE AND OVERALL FORMATION CONSTANTS

The term stability is a loose term, when the term stability is used without

qualification, it means that the complex exists and under suitable conditions, it may be

stored for a long time. The term can not be generalised for complexes. A complex may be

quite stable to one reagent and may decompose readily in presence of another reagent.

In studying the formation of complexes in solution, two types of stability of

complexes is found:

1. Thermodynamic Stability

This is a measure of the extent of which the complex will form or will be

transformed into another species under certain conditions, when the system has

reached in equilibrium. When we are concerned with this type of stability, we

deal with metal-ligand bond energies, stability constant etc.

2.Kinetic Stability

This refers to the speed with which transformation leading to the attainment of

equilibrium will occur. When we are interested in kinetic stability for complex

ions in solutions, we deal with rates and mechanism of chemical reactions.

These reactions may be substitution, isomerisation, recemisation and electron or

4

group transfer reactions. In the kinetic sense, it is more proper to call the complexes

inert or labile complex rather than stable or unstable complex. The complexes in

which the ligands are rapidly replaced by others are called labile, while those in

which substitution occurs slowly are called inert complexes. Stepwise and

Overall Formation Constants

According to J. Bjerrum (1941) the formation of a complex in solution proceeds

by the stepwise addition of the ligands to the metal ion. Thus the formation of the

complex MLn may be supposed to take place by the following n consecutive steps.

where M = central metal cation

L = monodentate ligand

n = maximum co-ordination number for the metal

ion M for the ligand

M + L ML K1 = ]][[)(LM

ML

ML ML2 K2 = ]][[

)( 2

LMLML

ML2 ML3 K3 = ]][[

)(

2

3

LMLML

Thus MLn-1 + L MLn Kn = ]][[

)(

1 LMLML

n

n

The equilibrium constants, K1, K2, K3, ..........Kn are called stepwise stability

constants.

5

The formation of the complex MLn may also be expressed by the following steps

and equilibrium constants.

M + L 1B ML, = ]][[

)(LM

ML

M +2L 2B ML2, 2 = 22

]][[)(

LMML

Thus M + nL nB MLn, n = nLMMLn

]][[)( ................(8.1)

The equilibrium constants, 1, 2, 3, .......... n are called overall formation or

overall stability constants. n is called as nth overall (or cumulative) formation

constant or overall stability constants.

The higher the value of stability constant for a complex ion, the greater will be its

stability. Alternatively 1/k values sometimes are called instability constant.

Stepwise and cumulative stability constants are also expressed as log10K1,

log10K2................log10Kn and log10n respectively.

Relationship or Interaction Between n and K1, K2, K3, ..........Kn

K's and 's are related to one another consider for example, the expression for 3

is:-

3 = 3]][[

)( 3

LMML

On multiplying both numerator and denominator by [ML] [ML2] and on

rearranging we get:

6

3 = ]][[]][[

]][[][

2

23

3

MLMLMLML

LMML

= ]][[

][]][[

][]][[

][

2

32

LMLML

LMLML

LMML

= K1 x K2 x K3

Thus n = ]][[

][]][[

][ 2

LMLML

LMML

.............]][[

][

1 LMLML

n

n

= K1 x K2..........Kn

or n =

nn

nnK

1

From above relation, it is clear that the overall stability constant n is equal to the

product of the successive (i.e. stepwise) stability constants, K1, K2, K3, ..........Kn. This in

other words means that the value of stability constants for a given complex is actually

made up of a number of stepwise stability constants.

Thermodynamic Importance of Stability Constants

In order to reach accurate conclusions regarding the nature of the forces acting

within complex species during their formation in solution, the energy changes

accompanying the reaction in question i.e. a complete thermodynamic characterisation of

the reactions is necessary at the very least, determination of enthalpy ( H ), entropy ( S

) and free energy ( G )changes accompanying complexation.

In the language of thermodynamics, the equilibrium constant of the reaction is a

measure of the change in free energy, heat content and entropy. A more useful manner of

stating equilibrium constant is in terms of the standard free energy change G , i.e. the

7

difference of free energy between the products and the reactants in a standard state,

which is related to equilibrium constants by the thermodynamic expression:

- RT log K = G = H - T S .....................................(8.2)

The reactions tends to go in the direction written, when G is negative.

Enthalpy change ( H ) gives the amount of heat either consumed or liberated per

mole of products and is related to the strength of the ligand to metal bonds, compared to

that of the metal to solvent bonds.

Entropy change ( S ) is related to the change in randomness (the disorder) of a

system. As is quite evident from the relation given above, complex formation is most

favoured by the negative enthalpy and positive entropy changes (either of the two or

both) as may be expressed by the equation:

log K = R

THS303.2

/ ......................................................(8.3)

In many reactions both the heat and entropy changes favour complex formation

but their relative importance changes markedly with minor variations from ML to M'L or

ML'.

FACTORS AFFECTING STABILITY

Factors related with Metal

The nature of the metal ions and the effect of the different physical properties of

the metal ions on the stability of the complex are:

1. Stability (or stability constant) increases with decreasing size of metal ion. K

generally varies are 1/r.

2. Stability constants for a complex increase with the charge of the central ion. The

8

K for the Fe(II) complexes will be less then the K for the corresponding Fe(III)

complexes.

3. The ions with high polarizability give complexes with higher stability constants.

Thus Cu(I) complexes have higher K values than the similar sized Na+

complexes, similarly of Ca2+ and Cd(II) or Al (III) and Ga(III) the former have

low K values for the complex formation.

4. Electronegativity increases the polarizing power and the ions with higher

electronegativity give stable complexes.

5. Ionization Energies: The electronegativity, covalent nature and ionic radii can

be related to the ionization energies of the atoms. It is found that the stability

constants for the metal complexes with a ligand increases with the ionization

energies of the metallic species.

Observations of Bjerrum Niecilson and others show that although most of the

metals of the periodic table form complexes, this tendency is the most with transition

metals. The reason being that the chelate effect is almost an entropy effect for the metal

ions of nontransitional group, while for the transitions metals it is partly an enthalpy

effect which increases the crystal field strength. The increase in crystal field strength

increases the points of attachment of the ligand to the metal ion imparting greater

chelating tendency to the latter (cf. CFSE). Fig 8.1

9

Fig. CFSE affecting stability of aquo-complexes

Chatt Ahrland classified the metals into a and b classes while a class metals form

stable complexes with ligands having the coordinating atoms, N, O, F (second period

elements), b class metals form stable complexes with ligands in which donor atom is P, S,

Cl (third or latter period elements).

The a class metals include H, alkali and alkaline earth metals; the elements from

Sc to Cr, Al to Cl, Zn to Br and lanthanides and actinides. While amongst b class Rh, Pd,

Ag, Ir, Pt, Au and Hg are included.

Elements from Mn to Cu, Tl to Po, Mo, Te, Ru, W, Re, Os, Cd are border line

metals.

It can be said with some approximation that increase in the ionic charge of the

metal ion and donor, will bring an increase in the chelating tendency while the increase in

ionic radius will decreases it. Thus small cation size, comparatively large ionic charge

and appropriate electronic arrangements are responsible for the maximum ability of

complex formation by transition elements.

10

Mellor and Maley have shown that the stabilities of the complexes of bivalent

metal ions follow the order: Pd > Cu > Ni > Pb > Co > Zn > Cd > Fe > Mn > Mg

irrespective of the nature of the ligand. Irving and Williams from the analysis of the data

on stability constants of transition metal ions, found that the order

Mn(II) < Fe(II) < Co(II) < Ni(II) < Cu(II) > Zn(II),

holds good. This order according to them follows logically from a consideration

of the reciprocal of ionic radius and second ionization potential of the metal, and is

known as 'Natural Order of Stability'.

Univalent ions have not been extensively studied but data on the complexes of the

univalent ions with dibenzol methanate ion shows the order of the stability as:

Ag > Tl > Li > K > Rb > Cs

For tetravalent metals much less information is available, the greater ease of

hydrolysis of these ions making potentiometric titrations more difficult. Irving and

Williams suggest from a considerable limited number of investigations that a rough order

of stabilities be:

Ti > Fe > Ga > In > Al > Cr > Sc

Factors Related With Ligands

The properties of the ligands which affect the stability of the metal complexes are

as under:

1. Basicity of the ligands: The greater is the Lewis base strength, higher is expected

to be the stability constant of the complex. Thus K values for the complexes are

11

expected to change in a manner similar to the changes in the proton association

constant (BH) for the ligands.

2. Dipole moment and polarizability of the ligands: Due to the greater

electrostatic interactions between the metal ion and the ligands, polarity and

ploarizability of ligand results in higher K for the complexes.

3. (ML) -bonding always increases the stability of the complex.

4. Steric factor: It play an important rule in determining the stability constants for

the complexes. Thus the 2 methyl derivative of 3 hydroxyquinoline gives much

less stable complexes then the parent compound because of the steric hindrance

caused by the methyl group adjacent to the site of co-ordination.

In complex formation hydrogen behaves just like a metal ion. Therefore, a ligand

with a larger affinity for proton will show the same behaviour towards the metal ions.

According to Riley any factor which can increases the localization of negative charge in

the co-ordinating ligands makes the electron more readily available and thus increasing

the co-ordinating ability of a base. The correlation between the basic strength of the

ligand and the stability constant of the complexes was pointed our first by Calvin and

Wilson.

Ring Formation and Size of the Ring

Ring complexes or chelates are very stable due to reduced strain. The number of

ring formed, the size of the rings and stabilizing or interfering resonance interactions are

determined by the structure of the chelating agent. The work of Ley on the chelates of

amino-acids showed that five and six membered rings are the most stable. Much evidence

has accumulated since then to prove that all chelates have either five or six membered

rings. Pfeiffer observed that in general the five membered rings is the more stable when

the ring is entirely saturated but when one or more double bonds are present, the six

membered rings is favoured. Schwarzenbach and Co-workers have observed that there is

12

a decrease in clate stability with the increase in ring size. The stability of a five

membered ring is not chiefly due to entropy but rather to the enthalpy of formation; the

example being 1, 2, 3 triamine- propane tetra chloroplatinum. Further the stability

increases with the increase in the number of rings in the molecule:

M(en) < M(trien) < M(EDTA).

(one ring) (two rings) (five rings)

Steric Effect:

Steric hindrance can influence stability in many ways, e.g.

(i) Metal-ligand bonds are weakened due to the presence of bulky group near the

coordinating site.

(ii) The substituting group prevents the ligand from assuming the planar configuration

and hence introduce strain in the metal-donor bond.

(iii)Steric hinderacne is also due to strained structure of the chelated ring, since it

breaks the usual linear configuration of the complexes.

From the study of the copper complexes of substituted malonic acids Riley

concluded that ethyl and propyl groups had a larger effect then methyl in reducing the

stability.

Resonance Effects

The stability of a chelated ring will depend on the possibilities of resonance in the

ring and on how these will fit in with resonance in the organic ligand itself. That

resonance may affect the formation of a chelate was first shown by Calvin and Wilson.

The double bond resonance has been attributed as a reason to be unusual stability of

histamine cobalt chelate.

13

Orbital hybridization

There are certain factors which serves to make a specific bonding arrangement

stable. As an example, the shape of , ', '' triaminotriethylamine is such that the

bonding atoms must be grouped tetrahedral round a metal atom. The ligand will therefore

tend to form a stable complex with a metal such a zinc, which favours sp3 hybridisation

in its 4-co-ordinate compounds, rather than with one such as copper which is limited to

dsp2 (planar) hybridisation. Similarly, triethylene tetra amine gives stable complex with

metal ions having dsp2 hybridisation, rather then sp3 hybridisation.

The chief factor responsible for the stability of the chelate ring is the entropy change

which can be viewed statistically or as probability factor. Considering the electronic

effect of the donor atom to be the same in the monodentate and the bidentatc ligands, it

can be seen that the dissociation of a monodentate from a complex will be higher than

that in the chelating bidentate. The dissociation of the M-L bond in monodentate will

release the ligand completely from the coordination sphere of the metal, so that it can be

easily swept off by the solvent. But the dissociation of one M-L bond for the bidenate

ligand does not release the ligand completely (for which simultaneous dissociation at

both ends is required). Hence the stability constant for metal chelate must be higher.

Consider the equilibrium reactions (Fig. 8.4):

[Co(NH3)6]3+ + 3en [Co(en)3]3+ + 6NH3 ...................(8.4)

Assuming that (i) Co-N bond strength in the two complexes is same (the f value

of ammonia and ethylendiamine are within 3%), and (ii) the entropy changes due to

structure making and structure breaking are negligible due to the similar size of the

complexes, it can be seen that the S o will increase for the reaction as the number of

moles of the products are more than those for the reactants. This will help the reaction to

go to the right.

14

METHODS OF DETERMINATION OF STABILITY CONSTANTS

There are many physical and chemical properties which may be used to detect the

formation of complex in solution and to measure the stability constants. The detection of

the complexes and the determination of the stability constants are very closely related.

Most of the methods used for the detection of complexes can also be used to determine

their stability constants.

The study of the complexes is supposed to be incomplete without finding the

stability or formation constants, because most of the properties and utility of the

complexes depend on it.

The value of stability constants may predict the conditions required for complete

formation of a given complex. This knowledge of the system is essential for correctly

interpreting its optical and kinetic properties of its partition equilibria and its biological

behaviour.

Further, it may also help in planning analytical and separation procedures. For

example in case where the species is highly coloured or can be precipitated from

solution, extracted into an organic solvent or absorbed on an ion exchange or

chromatographic column.

Stability constant is related with the thermodynamic parameters, as

-RT, Ink, = G = H - T S

Where, G , H and S are changes of free energy of enthalpy and of entropy

respectively.

The stepwise or overall stability constant, thermodynamic equilibrium constant

gives the value of free energy change, associated with the reaction. The corresponding

changes on entropy change of complex formation may be obtained by combining the

stability constants with the enthalpy change of complex formation, which is obtained by

15

determining the stability constant at a series of temperatures. The knowledge of entropy

is essential for the full understanding of many factors such as size, shape, electronic

structure of the central metal and the ligand, the temperature and the composition of the

solvent, which influence the stability of the complex.

Let us consider a reaction between a metal M and ligand L to form a complex

MmLn.

Mm + nL MmLn

K = nm LMLnMm

][][][

where 'K' is stability constant of the complex MmLn. The stability of the complex

is quantitatively expressed in terms of dissociation constant 1/k of the complex. The

latter is the tendency of the complex to split up into its components.

Some of the most important methods of determining the stability constants are

briefly described here.

1 pH - Metric Method

Bjerrum's Method

It is a potentiometric method for determining the stability constant for complex

formation. Although Bjerrum applied the method primarily to the binding of simple

molecules or negative ions to positive metal ions. It may be used with equal success with

chelating agents. The theoretical relationship outlined by Bjerrum are not restricted to

complex formation but may be applied to any equilibrium process regardless of the

nature of the interacting substances. Thus, it has been used with success on acid base,

and redox equilibria. Although the reactions to be considered involve ions that are more

16

or less completely hydrated, rather than the simple ions, but this fact does not affect the

validity of the conclusions, provided the activity of the water is maintained constant.

Formation or dissociation of a complex ion for molecule in the solution always

takes place in several steps, which can be easily determined by measuring pH in this

method.

Experimental Determination of Stability Constant by Bjerrum's Method

This is a potentiometric method. When the lignad is a weak base or acid,

competition between hydrogen ion and metal ions for ligand can be used to the

determination of the formation constant.

Let us consider the equilibrium in which an acid and metal ions are added to a

basic ligand in solution. Thus the following equation are obtained:

L + H+ Ka HL+, Ka = ]H][L[

]HL[

Basic Ligand Acid

L + M+ KF ML+, KF = ]M][L[

]ML[

Basic Ligand metal ion

Here Ka and KF are the acid association constant of the ligand and formation

constant respectively.

Now if CH, Cm and CL are the total amounts in moles/litre of acid (H+) , metal

(m+) and basic ligand (L), we have

CH = [H+] + [HL+]

17

CL = [L] + [ML+] + [HL+]

Cm = [M+] + [ML+]

Solving the last three equations given above and using the acid association

constant of the ligand, Ka. Then we get

[ML+] = CL - CH + [H+] - ]H[Ka]H[CH

[M+] = Cm - [ML+]

[L] = ]H[Ka]H[CH

Thus on putting the values of [ML+], [M+] and [L] from the above equation in

K1 = ]L][M[

]ML[

the value of K1 can be calculated. For the determination of [ML+], [M+] and [L],

the values of CH, CL, Cm, Ka and [H+], is generally determined potentiometrically using a

PH meter.

In order to get better results, the ligand must be a medium weak acid or base and

the formation constant, K1, should be within a factor of 105 of the value of the acid

association constant of the ligand, Ka.

Irving Rossotti Method

Calvin-Bjerrum pH titration technique as adopted by Irving & Rossotti is

18

generally used for determining the proton-ligand and metal-ligand formation constants.

The procedure consists of:

(A) Determination of the formation curve of the system. This is expressed as a plot

of n (formation function) against pL for metal ligand system and a plot of

n A

against pH for a proton-ligand system (Definitions of the terms n ,

n A and pL

are given below).

(B) The calculation of the values of formation constants by solution of the

formation function of the system or otherwise.

(C) The conversion or the stoichiometric constants into thermodynamic constants.

n term, was introduced by Bjerrum who called it the 'formation functions' or

'ligand number ' and is defined as the average number of ligand bound per metal atom or

ion present in whatever form.

n =

systeminpresentMofnumberTotalMmetaltoboundLligandofnumberTotal )()(

or n =

n

oi

n

oi

]MLi[.

]MLi[i. .....................................(8.5)

which can be written using equation (8.1) as,

19

n =

n

oi

i1

n

oi

i

[L]β.

β[L].i.[ = 1]................................(8.6)

A similar function for the proton-ligand sustems is n A, which defined as the

average number of protons bound per not complex bound ligand molecule, and can be

given by.

n A =

i

oi

iHi

i

oi

iHi

[L]β.

[H]β.i.[ H

o = 1]........................(8.7)

whereas, pL gives the free ligand exponent and may be defined as.

pL = log .][

1L

(A) Construction of the Formation Curves:

In Irving Rossotti method, this involves pH-titration of the following three sets of

mixtures (keeping total volume constant) against a carbonate free standard alkali:

(a) Mineral acid

(b) Mineral acid + Ligand solution

(c) Mineral acid + Ligand solution + Metal ion solution.

The ionic strength in each set is kept constant by adding appropriate quantities of

a neutral electrolyte solution. The temperature of the solution in each case is kept

constant. On plotting the observed pH against the volume of alkali, one obtains (a) and

20

acid titration curve, (b) a ligand titration curve and (c) a metal-complex titration curve,

corresponding to the above titrations. [Fig. (a)]

The calculation of n are made from the volume of alkali required to produce the

same pH value in the metal and ligand titrations. Similarly n A values are calculated from

the volume of alkali required to produce the same pH value in the ligand and mineral

acid titrations. According to Irving and Ressotti, n A and

n can be expressed as-

n A =

o

n

TLVV

ENVVTLoY)(

)()(1

1

....................(8.8)

n =

οTCM.n)V(V

TL)E(N)V(V

1ο

oοn1iii

....................(8.9)

Where Vo is the initial volume of the solution, Eo, TLo are the initial

concentrations of the mineral acid and the reagent respectively and V ', V'' and V''' are the

volume of alkali of a given normality, N, required during the acid, the ligand and the

metal titration respectively at a given pH (B). While the term Y gives the number of

titrable hydrogen ions arising from the chelating agent and TMo gives the initial

concentrations of the metal.

From the observed values of [L] for each n value, values of pL- are calculated

utilising the equation given by Irving and Rossotti:

pL- log10

0

0 .

)log1(

TCMnTCL

anti

jn

on

nHn

. o

iii

VVV 0

....................(8.10)

21

Values of proton-ligand formation constants, K H1 , K H

2 etc. obtained from the proton-

ligand formation curves plotted between values of n A and pH [Fig. 8.2(b)].

The pH value at n A = 0.5 gives the value of log K H

1 while the pH value at n A =

1.5 gives the value of K H2 and so on.

Similarly, the values of stepwise stability constants of metal-complexes are

obtained from the formation curve plotted between the values of n and pL- [Fig. 8.2(c)].

The value of formation constants are generally refined using least square method.

Fig. (a) pH - Titration Curves

(b) Proton-Ligand formation curve

22

(c) Metal-Ligand formation curve

In the present work the Calvin-Bjerrun pH-titration technique 30-31 as

adopted by Irving and Rassotti has been employed. the paracetamol (ligand) and complex

are soluble in water,while ibuprofen soluble in ethanol. The experiments were carried out

at the fixed ionic strength µ = 0.1M (KNO3) and temperature are298 k ,303 k and 308 k .

In order to investigate the equilibrium involving the formation of the complex, the

structure of proton ligand –equilibrium of various ligand employed is essential. The

proton –ligand stability constants of paracetamol have been calculated as ‘practical’

proton ligand stability constant following the method of Irving and Rossotti.

Paracetamol

Paracetamol, also known as acetaminophen or APAP, is a widely used over-the-

counter pain medication and antipyretic (fever reducer)Paracetamol is classified as a

23

mild analgesic. It is commonly used for the relief of headaches and other minor aches and

pains and is a major ingredient in numerous cold and flu remedies. In combination

with opioid analgesics, paracetamol can also be used in the management of more severe

pain such as post-surgical. Though paracetamol is used to treat inflammatory pain, it is

not generally classified as an NSAID because it exhibits only weak anti-inflammatory

activity.

While generally safe for use at recommended doses, even small overdoses can be fatal.

Compared to other over-the-counter pain relievers, paracetamol is significantly more

toxic in overdose but may be less toxic when used chronically at recommended dose

.Paracetamol is the active metabolite of phenacetin and acetanilide, both once popular as

analgesics and antipyretics in their own right. However, unlike phenacetin, acetanilide

and their combinations, paracetamol is not considered carcinogenic at therapeutic dose .

Ibuprofen

Ibuprofen, from isobutylphenylpropanoic acid, is a nonsteroidal anti-inflammatory

drug (NSAID) used for relieving pain, helping with fever and reducing inflammation.

About 60% of people improve with any given NSAID and it is recommended that if one

does not work that another should be tried Ibuprofen might be considered a weaker anti-

inflammatory than other NSAID.

Importance of stability constants

The stability constant of complexes has been found to be greater than zero, which is

24

perhaps one of the most convincing pieces of evidence for the existence of the complex

species MLn in solution. Moreover, if all the possible stability constants for a given

system have been Complex or coordination compound and any substance, which can

accept a pair of electron is called as Lewis acid whereas any substance donating a pair of

electrons is commonly called as Lewis base. When a ligand contains two or more donor

atoms close to each other, the metal complex formed is said to be a chelate and the

process is referred as chelating. The chelating ring may be ionic or covalent depending on

the nature of ligands. The history of complexes and the interpretation of complexes begin

with Alfred Werner. Coordinating agents are used in metal-ion sequestration or removal,

solvent extraction, dyeing, leather tanning, electroplating, catalysis, water softening, and

in other industrial processes. For example, vitamin B12 is a coordination compound of

cobalt, the hemoglobin of human blood is a coordination compound of iron, the

haemocyanin of invertebrate animal blood is a coordination compound of copper, and the

chlorophyll of green plants is a coordination determined, it is possible, in principle, to

calculate the equilibrium concentration or activity of each of the species present under a

known set of experimental conditions. Such exact knowledge of the composition of a

solution is essential for a correct interpretation of its optical and kinetic properties of

partition equilibria and its biological behavior.

25

EXPERIMENTAL AND DISCUSSION:-

a) Reagent:-

Water:-

Double distilled water was used which is free from carbon -dioxide

with pH of about 6.62

Oxalic acid :-

Dry Oxalic acid (AR) was used for the standardization of NaOH

Nitic acid:-

Nitric acid (AR) was used for the preparation of 0.01 N HNO3

and it was standardized by standard NaOH solution

Potassium Nitrate :-

1M potassium nitrate was prepared in distilled water .

Metal ion solutions:-

M(II) (0.01M) ionic solutions was prepared from metal nitrate of

AR grade and was standardized by standard methods.

Paracetamol :-

Ligand solutions of 0.05M was prepared in double distill water .

26

Ibuprofen:-

Ligand solutions of 0.05M was prepared in ethanol .

b) Instrument:-following equipments were used

Sr.

No.

Name of Equipment

1

2

3

4

5

6

7

Pranava Double distillation app with cut Off

DDU310350

BTI-digital water bath with constant temp

size-18*14*12

BTI digital magnetic stirrer with hot plate

capacity- 2 lit.

Nitrogen Cylinder 1.5 Cup cap

Nitrogen cylinder regulator single double

stage

Digital pH meter-Systronic model-361

Digital Electronic balance Contech CA-44

cap-40gm accuracy-0.1mg

27

Systronics microprocessor based instrument with accuracy of + 0.01 unit

with glass electrode and saturated calomel electrode was use for the

measurement. It was calibrated by buffer solution of pH 4.00, 7.00 and

9.20 at 270 C before proceeding the titration.

c) Calvin – Bjerrum Titration :-

The experimental procedure involved potentiometric titration of,

i) Free acid (0.01M) titration,

ii) Free acid (0.01 M) and ligand (0.05M) titration.

iii) Free acid (0.01M), ligand (0.05M) and metal ion (0.01M) against

std. NaOH solution.

The ionic strength of all solutions was maintain constant (0.1M) by

adding appropriate quantity of 1M KNO3solution.

The titration were carried out in a100 ml pyrex glass beaker kept in

a water bath maintain at constant temperature (25,30and 35 0C)

nitrogen gas was purged for chemically innert atmosphere. The

reading were recorded for each addition of 0.1ml. The graphs of

volume of alkali added against pH were plotted. The curve has

been designated as below:-

(i) Acid curve (A)

(ii) Ligand titration curve (A+L)

(iii) Metal – ligand titration curve (A+L+M)

28

Method for determination of stability constants:-

The dissociation constants of paracetamol was determined at 0.1 M ionic

strength pH – acid having only one dissociable H+ ion from –OH group and can therefore

represented

HL H+ + L-

The titration curve of the acid and the ligand deviates at about pH 3.0 and

then increase up to pH 12.0. The deviation between acid curve from ligand curve for the

systems showed the dissociation of H+ from OH groups of the ligands. (Table no 1 and 2)

Proton – ligand formation Number (nˉA):

Proton – ligand formation number (n A) were calculated by Irving and Rossotti

expression.

nA = γ – (E0 + N) (V2 + V1) / (V0 + V1) T0 L

Where, V0 = Initial volume of solution (50 ml)

N = Normality of sodium hydroxide

T0 L = Concentration of ligand in 50 ml solution

E0 = Initial concentration of free acid (HNO3)

γ = Number of dissociable proton from ligand

V1 and V2 – Volume of alkali consumed by acid and ligand on same pH

29

Metal- ligand formation number (n):-

The deviation of (A + L +M) curve from (A +L) started at about pH 3.5, It indicate the

commencement of complex formation .

Metal – Ligand formation number (n) was calculated by following expression.

n = (E0+ N) (V3 – V2) / (V0 + V2) (T0 M ) × nA

V0 = Initial volume of solution (50 ml)

N = Normality of sodium hydroxide

T0 M = Concentration of the metal ions

nA = Proton – ligand formation number

E0 = Initial concentration of free acid (HNO3)

Where, V2 and V3 – volume of NaOH consumed by ligand and metal ions

at same pH.

The values of n along with the difference between the volume of alkali

required for ligand and metal complex, titration curves (V3 – V2 )

30

Metal – Ligand Formation curves:Formation Curves were plotted between n and pH

.The metal-ligand stability constants were determination by half integral method.

The metal- ligand stability constants (log, k1 and log k2 values) are calculated from

formation curves. The values of n = 0.5 which corresponds to value of logk1 at n = 1.5

corresponds to log k2

Table. Stability constants and thermodynamic parameters of Co2+,Zn2+ and Cd2 +with paracetamol

System Temperature

(oC)

pKa logK -∆ H

(KJmol–1)

At 300C

-∆ G

(KJmol–1)

-∆ S

(KJmol–1 deg–1)

At 300C

Zn(II)Paracetamol 25 9.54 3.98

7.147

22.709 0.049

30 9.50

3.96

22.162

35 9.45

3.94

22.52

Co(II)Paracetamol 25 9.54

4.01

10.721 22.880 0.040

30 9.50

3.97

23.032

35 9.45

3.94

23.235

Cd(II)Paracetamol 25 9.54

4.429

17.86 25.271 0.025

30 9.50

4.428

25.689

35 9.45

4.427

26.107

31

Table: Stability constants and thermodynamic parameters of Co2+,Zn2+with Ibuprofen

System Temperature pKa logK -∆ H (KJmol–1) At 300C

-∆ G (KJmol–1)

-∆ S (KJmol–1 deg–1) At 300C

Co(II)Ibuprofen 25 5.7446 7.1956

3.573

41.057 0.12549

30 5.7385

7.17

41.597

35 5.7215

7.16

41.018

Zn(II)Ibuprofen 25 5.7446 7.2649

5.503 41.452 0.1194

30 5.7385

7.1849

41.683

35 5.7215

7.1695

42.280

CONCLUSION

The results obtained from the pH metric measurements, the values of pKa were found to decrease with increasing temperature. The values of the thermodynamic functions ΔG, ΔH and ΔS were calculated. The values of stability constants reveal that the stability constants decrease with increasing temperature, along with the pKa value.

ACKNOWLEDGEMENT:

Author are thankful WRO UGC for financial support to minor research project author are thankful to Principal Shri R.R.Lahoti science college Morshi for providing necessary facility.

32

REFERENCES

1) Khan, Firoj and Khan, Farid J Indian Chem, soc,74,171 (1997)

2) Jain, A.K. and kataria, H.C. Asian J. chem. 12(2), 574 (2000)

3) Sharma, G. and Chandel C.P.S. Orient J of chem. 17(2), 195 (2001)

4) Chaturvedi, D.N. and Gupta, C.M. Indian J chem. 1-12 (No5), 53(1974)

5) Bhattacharyya, R.G. and Paul, U.K. Indian J pharma. Sci, 43, 219 (1981)

6) Bhattacharyya R.G. and paul U.K. Indian J chem. 29A, 986 (1990).

And chatterjee, AB

7) Manjula, V and Bhattacharyya, P K J.chem. Soc, Dalton trans,

567(1989)

8) Manjula, V. and Bhattacharyya, P.K. Indian J. chem. 29 A, 400(1990).

9) Rajathiromoni, P, Arasu, P.T. Indian J. chem. 31A, 760 (1992)

and Nair, M.S.

10) Nair, M.S. Arasu, P.T. Fernando, T.L, Indian J. chem. 32A, 807(1993)

Pillai, M.S. and Chandran, M

33

11) Nair, M.S.,Arasu, P.T., Mansoor,S.S. Indian J chem., 34A, 365 (1995)

Shengbagavalli, P. and Neekantan,M.A.

12) Ritsma, J.H. and jellinek, F. Reel. Trav. Chim., 91, 923

(1972)

13) Lenz, G.R. and Martell, A.E. Biochemistry, 3. 745 (1984)

14) Roul,L.K. and Patnaik, R.K J. Indian chem. Soc.,

69,327(1992)

15) Nair, M.S. and Neelakantan, M.A. J. Indian chem., Soc. 77, 273

(2000)

16) Mukherjee, G.N. and Ghosh, T.K. Indian J. chem., 30A, 1033

(1991)

17) Mukherjee, G.N. and Ghosh, T.K. J. Inorg. Biochem., 57, 827

(1995)

18) Mukherjee, G.N. and Ghosh, T.K. J. Indian chem. Soc., 71, 169

(1994)

34

19) Mukherjee, G.N. and Ghosh, T.K. J. Indian chem. Soc. 74, 538

(1997)

20) Mukherjee, G.N. and Ghosh, T.K. J. Indian chem. Soc., 71, 249

(1994)

21) Mukherjee, G.N. and Ghosh, T.K. Proc. Indian Acad. Sci.,

108, 371 (1996)

22) Chakrawarti, M. and Maini, P. J. Indian Chem. Soc., 77,

217 (2000)

23) Bisht, J.P.S., Bisht, N.P.S. and Singh, Indian J. Chem. 28A, 812

S.P. (1984)

24) Srivastava, H.P. and Srivastava, R.K. J. Indian Chem. Soc., 72,

551 (1995)

25) Keemti, Lal. Indian J. Chem., 17A, 313

(1979)

26) Sharma, R.K., Joseph, S. and Kidwai, M. J. Indian Chem. Soc, 75,

93, (1998)

35

27) Suma, S., Sundarsanakumar, M.R., Nair, Indian J. Chem., 33A,

C.G.R. and Prabhakaran,p. 1107 (1994)

28) Suma, S. sundarkumar, M.R., Nair, C.G,R, Indian J. Chem. 33A,

and Prabhakaran. P. 775 (1994)

29) Sharma, R.K. and Sherly Joseph Indian J. Chem. 35, 639

(1996)

30) Bjerrum, J. Metal amine formation

in aqueous solution ,

Haase, Copenhagan (1941)

31) Calvin, M., and Wilson, K.W. J. Am. Chem, Soc, 67,

2003 (1945)

36

PROJECT COMPLETION REPORT

Generally complexes are designated as stable or unstable. The general meaning of

stability is supposed to be related with the concept, whether a particular complex can be

converted into other easily or not. As a matter of fact, this is kinetic aspect of stability;

which deals with the rate of the reaction and its mechanism. The other aspect of stability

is thermodynamic aspect. In which stability of a complex is related with the amount of

energy released during its formation or the amount of energy required to break it.

In this unit we describe complex forming equilibria in solution and the various

factors affecting it. We will also discuss the various factors affecting stability constants

for the formation of complexes in solution. In the end of the unit we shall describe the

method used for determining stability constants of the complexes formed in solution.

Which involves quantitative characterisation of the complex-forming reaction in solution.

OBJECTIVE

The main aim of this project is to study the complex formation equilibria in solution.

After going through this unit you should be able to:

describe stepwise and overall formation constants;

explain thermodynamic importance of stability constants;

discuss factors affecting stability of complexes; and

describe methods of determining stability constants for binary complexes in

solution.

EXPERIMENTAL AND DISCUSSION:-

d) Reagent:-

Water:-

Double distilled water was used which is free from carbon -dioxide

with pH of about 6.62

Oxalic acid :-

Dry Oxalic acid (AR) was used for the standardization of NaOH

Nitic acid:-

Nitric acid (AR) was used for the preparation of 0.01 N HNO3

and it was standardized by standard NaOH solution

Potassium Nitrate :-

1M potassium nitrate was prepared in distilled water .

Metal ion solutions:-

M(II) (0.01M) ionic solutions was prepared from metal nitrate of

AR grade and was standardized by standard methods.

Paracetamol :-

Ligand solutions of 0.05M was prepared in double distill water .

Ibuprofen:-

Ligand solutions of 0.05M was prepared in ethanol .

e) Instrument:-following equipments were used

Sr.

No.

Name of Equipment

1

2

3

4

5

6

7

Pranava Double distillation app with cut Off

DDU310350

BTI-digital water bath with constant temp

size-18*14*12

BTI digital magnetic stirrer with hot plate

capacity- 2 lit.

Nitrogen Cylinder 1.5 Cup cap

Nitrogen cylinder regulator single double

stage

Digital pH meter-Systronic model-361

Digital Electronic balance Contech CA-44

cap-40gm accuracy-0.1mg

Systronics microprocessor based instrument with accuracy of + 0.01 unit

with glass electrode and saturated calomel electrode was use for the

measurement. It was calibrated by buffer solution of pH 4.00, 7.00 and

9.20 at 270 C before proceeding the titration.

f) Calvin – Bjerrum Titration :-

The experimental procedure involved potentiometric titration of,

iv) Free acid (0.01M) titration,

v) Free acid (0.01 M) and ligand (0.05M) titration.

vi) Free acid (0.01M), ligand (0.05M) and metal ion (0.01M) against

std. NaOH solution.

The ionic strength of all solutions was maintain constant (0.1M) by

adding appropriate quantity of 1M KNO3solution.

The titration were carried out in a100 ml pyrex glass beaker kept in

a water bath maintain at constant temperature (25,30and 35 0C)

nitrogen gas was purged for chemically innert atmosphere. The

reading were recorded for each addition of 0.1ml. The graphs of

volume of alkali added against pH were plotted. The curve has

been designated as below:-

(iv) Acid curve (A)

(v) Ligand titration curve (A+L)

(vi) Metal – ligand titration curve (A+L+M)

Method for determination of stability constants:-

The dissociation constants of paracetamol was determined at 0.1 M ionic

strength pH – acid having only one dissociable H+ ion from –OH group and can therefore

represented

HL H+ + L-

The titration curve of the acid and the ligand deviates at about pH 3.0 and

then increase up to pH 12.0. The deviation between acid curve from ligand curve for the

systems showed the dissociation of H+ from OH groups of the ligands. (Table no 1 and 2)

Proton – ligand formation Number (nˉA):

Proton – ligand formation number (n A) were calculated by Irving and Rossotti

expression.

nA = γ – (E0 + N) (V2 + V1) / (V0 + V1) T0 L

Where, V0 = Initial volume of solution (50 ml)

N = Normality of sodium hydroxide

T0 L = Concentration of ligand in 50 ml solution

E0 = Initial concentration of free acid (HNO3)

γ = Number of dissociable proton from ligand

V1 and V2 – Volume of alkali consumed by acid and ligand on same pH

Metal- ligand formation number (n):-

The deviation of (A + L +M) curve from (A +L) started at about pH 3.5, It indicate the

commencement of complex formation .

Metal – Ligand formation number (n) was calculated by following expression.

n = (E0+ N) (V3 – V2) / (V0 + V2) (T0 M ) × nA

V0 = Initial volume of solution (50 ml)

N = Normality of sodium hydroxide

T0 M = Concentration of the metal ions

nA = Proton – ligand formation number

E0 = Initial concentration of free acid (HNO3)

Where, V2 and V3 – volume of NaOH consumed by ligand and metal ions

at same pH.

The values of n along with the difference between the volume of alkali

required for ligand and metal complex, titration curves (V3 – V2 )

Metal – Ligand Formation curves:

Formation Curves were plotted between n and pH .The metal-ligand

stability constants were determination by half integral method.

The metal- ligand stability constants (log, k1 and log k2 values) are calculated from

formation curves. The values of n = 0.5 which corresponds to value of logk1 at n = 1.5

corresponds to log k2

Table. Stability constants and thermodynamic parameters of Co2+,Zn2+ and Cd2 +with paracetamol

System Temperature

(oC)

pKa logK -∆ H

(KJmol–1)

At 300C

-∆ G

(KJmol–1)

-∆ S

(KJmol–1 deg–1)

At 300C

Zn(II)Paracetamol 25 9.54 3.98

7.147

22.709 0.049

30 9.50

3.96

22.162

35 9.45

3.94

22.52

Co(II)Paracetamol 25 9.54

4.01

10.721 22.880 0.040

30 9.50

3.97

23.032

35 9.45 3.94 23.235

Table: Stability constants and thermodynamic parameters of Co2+,Zn2+with Ibuprofen

System Temperature pKa logK -∆ H (KJmol–1) At 300C

-∆ G (KJmol–1)

-∆ S (KJmol–1 deg–1) At 300C

Co(II)Ibuprofen 25 5.7446 7.1956

3.573

41.057 0.12549

30 5.7385

7.17

41.597

35 5.7215

7.16

41.018

Zn(II)Ibuprofen 25 5.7446 7.2649

5.503 41.452 0.1194

30 5.7385

7.1849

41.683

35 5.7215

7.1695

42.280

CONCLUSION

The results obtained from the pH metric measurements, the values of pKa were found to decrease with increasing temperature. The values of the thermodynamic functions ΔG, ΔH and ΔS were calculated. The values of stability constants reveal that the stability constants decrease with increasing temperature, along with the pKa value.