SD0100031

Analytical Application of Aminohydroxamic Acids

By

atari \^Aibd (L*ud ^jradt L^l

B. Sc. Chemistry

A thesis submitted in fulfillment for the M. Sc. Degree in Chemistry

Department of Chemistry

Faculty of Science

University of Khartoum

November 2000

3 2 / 3 7 To

The soul of my mother

To my father,

My children,

My wife,

My brothers,

and sisters. DISCLAIMER

Portions of this document may be illegible in electronic image products. Images are produced from the best available original document Ac It n o w I e d g m e n ts /

I would like to express my deepest thanks and gratitude to my

supervisor Dr. Taj Lilsir Abass for his directions, invaluable advice,

encouragement and guidance throughout the course of this study.

Thanks are also extended to the slalTof chemistry department for

their permanent do-operation.

My thanks also to my friends for their assistance and support. Last my special (hanks to my IViend Abdelhafe'ez and Abd El Latif for typing

this thesis. Abstract

Anlbranilic !|ydroxamic ^acad was {prepared by coupling of methylanthranilate! (prepared by esterification of anthranilic acid with methyl alcohol using the Fisher—Speir method)' with freshly prepared hydroxylamine.

The ligand was characterized by the usual reaction of hydroxamic acid with acidic V( V) and Fe(III) solutions that gives' a blood-red colour in amyl alcohol and deep-violet colour in, aqueous solution, respectively.

The absorbance of Fe(III)-hydro'xamic acids complexes increases with increase of pH. In this study, the effect of pH on the absorbance of

Fe(lll)-anlhranilic hydroxamic acid was in accordance with this trend. The maximum absorbance was obtained at pH 5.0 at maximum wavelength of

482 HID.

For Cu(II)—anthranilic hydroxamic acid complex, the use of acidic or basic pi I lead to the precipitation of Cu(II)-ligand complex. But when using buffer pi I (acetic acid/sodium acetate)a clear green colour of Cu(II)-

Ligand complex was obtained. The maximum absorbance'was obtained at pi 1 6.0 at maximum wavelength of 390 nm.

V(V)-anmranilic hydroxamic acid complex was extracted in acidic medium in amyl alcohol at pH 2.0 because in aqueous solution V(V)- anthianilic hydroxamic acid complex has not clear colour. It was observed that the maximum extraction in acidic medium decreases sharply with the

Hi increasing' of pH yajue.. The maxi'mum; wavelength for maximum absorbance was recorded at 472 nm.

V(V) interfered with deterrnination of Fe(III) above concentration of

2 pprn^ whereas Cu(II) interfers slightly with'the determination of Fe(III) ions even at a high concentration of the Cu(II) ions.

i •

Both Cu(ll) and Ni(II) do ridt interfer with the determination of V(V) ions even at high concentrations, Fe(III) ion produced slight interference, while Mo(VI) ions have a pronounced interference.

Both. V(V) and Fe(III) iong ... interfered markedly with the deterniination of Cu(II) ions,and made impractical under these conditions.

However, the calibration curves for the three metal ions produced a practical linear dynamic range.

IV T^ble of Contents

The Contents Page Dedication I Acknowledgment II Abstract in English III Abstract in Arabic V The Contents VI List of tables IX List of figures X Chapter One Introduction i , 1 1.1 Complexation reaction I 1.1.1. Complexation ability of metals 2 1.1.2 Complexation ability of ltgand . 3_ 1.2. Interference , 4 1.3. Masking and demasking of chemical reaction 5 1.4 Solvent extraction 8 1.4.1. The Distribution ratio 9 1.4.2. The percent of extraction 10 1.4.3 The extraction process ; 11 1.4.4. The separation efficiency of metal chelates 12 1.5. Hydroxamic acids 13 1.5.1. Structure of hydroxamic acid 14 1.5.2. Nomenclature 18 1.5.3. Acidity of hydroxamic acids 19 1.5.4. Preparation of hydroxamic acids 20 1.5.5. Reactions of hydroxamic acids 22 1.5.6. Detection of hydroxamic acids 29 1.5.7. 1. R. of hydroxamic acids 29 1.5.8. Monohydroxamic acids as a ligand , 30 1.5.9. Aminohydroxamic acid as ligands • 32 1.5.9.1. Acidity of aminohydroxamic acid ' 33 1.5.9.2. Complex formation and analytical application of 34 aminohydroxamic acids .5.9.2.1. Copper complexes 35 .5.9.2.2. Nickel complexes ' ' .37 .5.9.2.3. Zinc(II) and cobalt(II) complexes • 39 .5.9.2.4. lron(lll) Complexes ' 39 1.5.9.3. The effect of side chain donors on the coordination 43 of a-amino hydroxamic acid .1.5.9.4. Derivatives of p-aminohydroxamic acid 46 1.5.9.5. Complexes of 8-aminohydroxamic acids 50

VI 1.3.10. Th6 Biological activities' of ftydroxamic acids 5,0 Chapter "two 2. Experimental ' 53 2.1. Instruments, 53 2.2. . Chemicals , 53 2.3. Synthesis of reagents. 53 213.1 Synthesis of anthranilic hydroxamic acid 53 2.3.1.1.. From the .coupling reaction of 54 N-hydroxyphthalimide anthranilic acid ester, and hydroxylamine 2.3.1.1.1. Preparation of N-hydroxyphthalimide • 54 2.3.1.1.2. Esterification of anthrahilic acid by N- . 54 hydroxyphthalimide \. 2.3.1.1.3. The coupling reaction between the N~ 55 hydroxyphthalimide and hydroxylamine 2.3.1.2.' Preparation of anthranilic hydroxamic acid from the 56 coupling reaction between methyl anthranilate and , hydroxylamine 2.3.1.2.1. Preparation of methyl anthranilate 56 2.3.1.2.1.1. Identification of methyl anthranilate . 57 2.3.1.2.1.1.1 Nitrogen contents (N%) .' 57 2.3.1.2.1.1.2 I.R spectra 58 2.3.1.2.2. The coupling reaction between methyl anthranilate 58 and the hydroxylamine 2.3.2. Characterization of the anthranilic hydroxamic acid 59 2.3.2.1. IR spectra ,.. 59 2.3.2.2. Nitrogen contents (N%) . , 60 2.4. Preparation of reagents : 60 2.4.1. Preparation of pH 1.0 to pH 10.0 buffer solutions 62 2,5. Analytical parameters for Fe(III),V(V) and Cu(II) 64 ions • 2.5.1. Determination of wavelength of the maximum 64 absorbance for 'Fe(III) 2.5.2. The effect of the ph on the absorbance using HC1, 64

HNO3 and H2SO4 acids , 2.5.3. Construction of calibration curve for Fe(III)- 65 anthranilic hydroxamic system ; 2.5.4. The effect of. foreign ions, on the absorbance of 68 Fe(lll) anthranilic hydroxamic acid 2.5.4.1. , The effect of V(V) on the absorbance or Fe(III)- 68 ligand 2.5.4.2. The effect of Cu(II) on the absorbance of Fe(III)- 69 anthranilic hydroxamic acid at pH5 and 482 nm.

VII 2.6. Analytical parameters for V(V) 69 2.6.1. Determination of V(V) wavelength of the maximum 69: absorbance 2.6.2. The effect of thei>pH on the absorbance ofVCV)- 70 Ligarid complex using HC1, HNO3 and H2SO4 2.6.3. Construptioh of j calibration curve for V(V)- 71 anthraniHc hydrox^mic acid cpmplex 2.6.4. The effect of foreign |8ns {Mo(VI), Fe(III) Cu(II), 74 Ni(II)} I on the •absorbance of V(V)-anthranilic hydroxamic acid system at pH 2 and 472 nm 2.6.4.1. Effect of Fe(III) on the absorbance of V (V)-ligand 74 •..: complex • ; • . , . ' ' 2.6.4.2. The effect k>f Mo(VI) on. the absorbance of V(V)- 75 ligand complex at pH2 and 472 nm 2.6.4.3. The effect of Ni(II) on the absorbance of V(V)- 75 ligand complex at pH2 and 472 nm 2.6.4.4. Theyeffect of Cu(II) on^the absorban.ce of V(V)- 76 ligand complex at?pfH2 and 472vnm • 2.7. Analytical parameters for Cu(II) ion 77 2.7.1. Determination of Cu(II) wavelength of the maximum 77 absorbance 2.7.2. -The effect of the*;:buffer: pH on the absorbance of 77t , ' Cu(II)-ligand complex : 2.8. Construction of calibration curve for Cu(II)- 78 anthranilic hydroxamic acid complex 2.9. The, effect of foreign ions Fe(III), V(V) on the '81 absorbances of Cu (II)-anthranilic hydroxamic acid at buffer pH6 and 390 nm 2.9.1. The effect of Fe(III) on the absorbance of Cu(II)- 81 ligand complex 2.9.2. The effect ofV(V) on the absorbance of Cu(II)- 82 ligand complex 1 Chapter three : , , T Discussion ' 83 References 87

VIII List of tables

Table 1 N%'o£methylanthranilate 58 Table 2 '.' :• The!I,R spectra for methyl anthranilate 58 Table 3 The IR spectra of anthranilic hydroxamic acid • 60 Table 4 N% of anthranilic hydroxamic acid 60 Table5 The effect of pH of an absorbance of Fe(III)-ligand with 65 different adijds at 482 nm Table 6 Calibration for Fe'(III) at 482 nm 66 Table7 The effect of V(V) on the absorbance of Fe(III) at pH5, 68 and482nm Table 8 The effect of Cu(II) on the absorbance of Fe(III)- 69 anthraru'lic hydroxamic acid at pH5 and 482 nm Table 9 ,• The effect of pH oh the absorbance of V(V)-ligand with 70 effect acids at 472 nm Table :10 Calibration for V(V) at 472 nm pH2 71 Table 11 The effect of Fe(III) on the absorbance of V(V) at pH2 74 and 472 nm Table !12 The effect of Mo(VI) on the absorbance of V(V)-ligand 75 Table 1 The effect of Ni(II) on the absorbance of V(V)- • 76 anthranilic hydroxamic acid at pH2 and 472 nm Tablet 4 The effect of Cu(II) on the absorbance of V(V)-ligand 77 complex at pH2 and 472 nm Tablel 5 The effect of pH on the absorbance of Cu(II)-ligand 78 complex Table 116 Calibration of Cu(Il)-Hgand complex at pH6 and 390 nm 79 Tablel 7 The effect of Fe(III) on the absorbance of Cu(II) 81 anthranilic hydroxamic acid at pH6 and 390 nm Table 18 The effect of V(V) on the absorbance of Cu-ligand 82 complex at pH6 and 390 nm

IX List of figures

Fig. 1 Proposed structures of various speeies in the Fe(III)-Glyha 42 Fig. 2 Calibration curve for Fe(III) 67 Fig. 3 a Caiibiration curve for V(V) 72 Fig. 3 b Calibration curve for V(V) 73 Fig. 4 Calibration curve for Cu(II) 80 CHAPTER ONE

1. Introduction

1.1 Compiexation reaction:

A complexation reaction involves the replacement of one or several of the solvent molecules by other groups'". The grbups bound to a central ion are called ligands. The complexation of metal ion in an aqueous solution is expressed by the following equation:

M(H2O)U+L> M(H2O)n.,L+H2O

The ligand L can be either a molecule or charged ion. Other L groups can successively replace;, the remaining aquo groups in the complex until the complex ML,, is formed' '. If a metal forms a complex ion ML,, the maximum number of bound ligands is the coordination number of the metal ion. It can may be noted, however, that a metal ion can have more than one characteristic coordination number depending upon the valence of the central atom and upon the nature of the coordinationvligand.

A ligand that is attached to the metal ion at only one point is called unidentate (one-toothed); , Water, ammonia and the halides are all unidentate .ligands1"'. Organic molecules or ions containing two or several bonding centres, i.e. donor groups, are able to replace two or more aquo or other unidentate groups from, a metal ion. Such ligands are termed bi, tri, quadri, or generally muitidentate. A muitidentate ligand may be bound to the metal by means of two or more (teeth) in such a way that a ring structure is foifmed. Such ring compounds are called jchelates and

-'• .it: . • multjtjentate reagent, chelating ^gents(2).,The chelates are characterized by .'t, their remarkably high stability as compared to the complex compounds formed by unidentate reagent(2\

According to Lewis, a complexation reaction is a kind-6f neutralization. The ligand is an electron pair donor and act as a base and the metal ion accepts the electron pair and acts as an acid.

1.1.1 Comple*ation ability of metals:

Valence, forces are electrostatic by nature, and it is obvious that the size and charges of the components are important factors. It is also obvious that if the ligand is a molecule, its dipole moment will be important among the factors determining its eomplexing power.

Actually, however, we must take into consideration that the electronic structure of an ion or molecule will be modified by other electric

•fields, it is therefore necessary to consider the deformabiHtyof the central ion and the polarizability of the ligand. The deformability of a metal ion increases with number of subshell electrons; a ligand is polarizable if it has a large electron shell and a low charge.

Taking transition metal ions with incomplete subshells as complexing metal ions and on basis of these factors- Hie charge, the size, and the ionization potential of. the ion, it is often possible to predict qualitatively the complexation ability of various metal ions. The most extensively investigated; is;:the series of divalent metals, Mn(II), Fe(II),

Co(II), Ni(II), CLI(II), and Zn(H).In this series, the ionic radius, decreases and the ionization potential increases up to copper. Accordingly, the stability increases progressively to reach a maximum with copper. Irving and Williams propounded this rule, and the series is' often named after them(2).

The complexation tendency of metals belonging to the Irving-

Williams series is particularly pronounced when nitrogen, carbon, or sulphur is the coordinating atom but less pronounced when oxygen is,the coordinatioi';i atom. Usually a low ionic charge and large ionic radius favoured covalent bonds. Therefore, low valence states are particularly stable.

1.1.2. Gomplexation ability of ligdnd:

According to the thermodynamic investigations, the stability of complex depends mainly on entropy factors if oxygen is the, donor atom and the bond are predominantly ionic, whereas enthalpy factors are decisive if nitrogen is the. donor atom and the bonds are more covalent in character. ; !

\ • '

The strongest complexing agents are those that are multidentate and form the particularly stable, five-membered rings. Effective multidentate chelating agents are molecules containing both oxygen and nitrogen as donor atoms. 1.2, Interference:

• 'i . •

The fact that two components of a reaction are able to combine not only in one but alsb in two or several ratios often reduces the accuracy of the analysis. Many reagents will form only one coloured complex, ML, with a mefaj ion if used in slight excess, but if large excess is used another complex e.g., MLT will be formed in considerable amount or even predominantly. If the constants are known, it is easy to estimate this interference and conditional Constants are appropriate for this purpose'21.

As a rule, no high degree of accuracy will be attained if spectrophotometric method js based on the measurement of the absorbance of a mixture containing several coloured compounds in equilibrium. The results are usually ,poorly reproducible and are affected by even small temperature variations which may shift the equilibria considerably, conditional constants calculated for spectrophotometric analysis are not always identical with the conditional constant valid for the complexometrj.c titrations12'.. The completeness of a complexometric tit-ration is favoured by the formation of any compound containing M and R in the right ratio (for instance acid and basic complexes). In spectrophotometric analysis, however, any compound containing M but with another composition than

MR,, has different light absorption and must thus be considered the product of an interfering side reaction. Only in exceptional cases may two different' compounds have the same particular wavelength e.g. at isobestic point(2). If the complexdmetric titration of a metal M with EDTA is based on the formation of the complex MY, iv^^^mplexes possibly formed required the introduction of an otMY(H) coefficient, which increases the conditional constant'2', On the other hand, if M is determined spectrophotometrically as

MY, the formation of MHY represents an interfering side reaction which can be taken into account by introducing a coefficient-ctM(HY) that decreases the conditional constants*'1.

1.3. Masking and demasking of chemical reaction:

The suppression of a reaction by addition of some complexing agent is usually expressed by stating that the metal ion is masked or sequestered.

Masking has been defined as a way of changing the usual course of a chemical reaction'1'. The fundamental equation for suppression of the formation of a compound ML by masking is:

KNn; = KM1 ,/aM (A), a L (B) where A and B represent masking agents.

Masking agents are extensively employed in complexometric titrations of solutions containing several metal ions. A masking reagent is one that lowers the concentration of a free metal ion or a free ligand to such a level that some or all of chemical reactions are prevented01. " :

In many instances only a weak masking effect is required, leaving the metal species free to react with more powerful complexing agents'"0. Varying the pH of a solution is the one of the most important'ways

* • ' r • ' ' of changing |he effectiveness with wljich'the m,etal ion is masked by a complexing species. A familiar example of pH dependent complexing ability is in the .use of hydrogen sulphide as a precipitant in the group separations of metals. Because hydrogen sulphide is a weak acid, proton compete strongly with metal ions for addition to the sulphide ion, S2", so that at a lower pH values the only metals that can be precipitated are those forming, very insoluble sulphides. With increasing pH of the solution proton competition becomes less important and metal ion forming less insoluble sulphides are also precipitated10.

Because most ligands are proton acceptors their masking ability varies with pH in a way that depends specifically on their pka values decreasing as , the solution, is ,made more acidic. Also in general, the stability constants of two different kinds of metal ions will not vary by a. constant amount from one .kind of complexing species to another. These: two factors have, important-consequences for masking and demasking because in most cases, use is made of these differences to form a complex between one kind of metal ion and a reagent (the principal reaction) whereas, because of their more favourable equilibria with the masking agent, the reactivity of the other metal ions is suppressed'" \

Masking procedures are usually of the following types(J): 1/ A masking agent' is added to a solution prior to determination so as to

form a stronger complex with interfering, ion, than with the desired species

which, instead is able to react normally. The amount of the metal ion

displaced in this way by the masking agent can then be determined from

the amount of free Ugand that is liberated.

2/ A masking agent changes the valence state of the interfering ion, for example, in acid solutions ks ascorbic acid reduces Fe(III) to Fe(II) and

Hg(II) is reduced to the metal.

3/ A masking agent precipitates the interfering ion in a form that can be left in the system; without causing trouble in further determination. Above pH 12 calcium can be titrated with EDTA in the presence of magnesium ion which is precipitated as the hydroxide.

4/ Advantage is taken of differences in' the rates of formation or dissociation of complexes by the ion to be masked and the ion to be determined. In the cold Fe(III) can be titrated with EDTA in the presence of Cr(III) because the latter reacts only very slowly(3).

The reverse procedure, demasking, occurs when a substance is added to counteract the effect of a masking agent which is already present in the solution; for example fluoride ion is a masking agent for Sn(IV) against precipitation as Sn(lV) sulphide, but the tin is demasked and can be precipitated in this way if boric acid is added to remove the fluoride ion by converting it to the very stable species BF4". Because masking is often achieved by pfJ adjustment or byadding excess of a suitable complexing agent touthe solution, the procedure may appear, to be deceptively simple.

The qualitative factors involved in the selection of masking agents can be readily understood from the properties of metal ions and complexing species. TriUs, in spectrophotorhetry there is additional requirement that the masked and the analytically important species must not absorb light in the same region of the spectrum, whereas, in gravimetric analysis it is necessary that ihe masked species is not coprecipitated(3).

1.4. Solvent extraction:

Solvent extraction is a method of separation based on the transfer of a solute from one solvent into another, essentially immiscible solvents' when the two solvents are brought into contact. The technique is extremely useful fbrvery rapidseparation of both organic and inorganic substances'4'

Extractions are based on the simple distribution law stating that "at a given temperature a solute becomes distributed between two,essentially immiscible solvents in such away that the ratio of the concentrations of the solute in the two phases will be constant". It is assumed that the solute is not charged and.has the same molecular structure in each phase'".

- _ S' (1) Kirk ; ~ s2

Where KD is the distribution coefficient, S, the concentration of the solute and subscripts represent solvent (j) which is organic solvent and solvent (2) which is an aqueous. Many substances are; partially idnized'in the aqueous solvent such as weak acids. This introduces a pH effect dn the extraction^. Consider for example, the extraction of benzoic acid from an aqueous solution. Benzoic acid (HBZ) is a weak acid in water with a particular ionization constant ka. The distribution coefficient KD is given by this equation

v _ (HBZ)L. p-, R { } °~ (HBZ)a :• ~

where (c) represent the ether solvent and (;1) representthe.aqueous.

However, part of the benzoic acid in the aqueous layer will exist as BZ" depending on the magnitude of Ka and on the pH of the aqueous layer<4).

1.4.1 The Distribution ratio:

It is more meaningful to describe a different term, the distribution ratio,- which is the ratio of the concentrations of all'the species of the solute in each phase. In the previous example it is given by:

A relationship between D and Kn can be derived:

The acidity constant for the ionization of the acid in the aqueous phase IS giv en by:

(H') (BZ-) Ka- a a . (HBZ)a + ^i{^Ka(HBZ)a/(H )a.:... (5)

n = KD (HBZ)a + (HBZ):1 + ka(HBZ)a/(H )a

1.4.2 The percent of extraction:

The distdbution ratio D is a constant independent on the volume ratio, however, the fraction of the solute extracted will depend on the volume ratio of the two solvents.

The fraction of.solute extracted is equal to the millimole of solute in the organic layer divided by the total numbers of millimoles,

Where Vo and Va are the volumes of the organic and aqueous

phases, respectively. The percent- extracted is related to the.

distribution ratio by

•rn/ - D+v;,/Vo

10 1.4L3 The extraction process:

The extraction proems ;can|b|,thojig^

steps each with equilibrium constant:

1/ The" chelating; agent HE distributed between the aqueous and the

organic phase

, (HR)0^=^ (HR)a

and SB •••<")

II The reagent in aqueous phase is ionized as follows:

HR^~—^ H+ + R"

and; , ' .

IU'1 ID1 . .{HRJ •

3/ the metal ion chelates with the reagent anion to form an uncharged

molecules

n N M - +.nR" v.: MR,,

and; •

4. The chelate distributed between the organic and aqueous phase

II and;

{MRp} —.....•..•...., MR and KDMR.I are the distribution coefficient of the reagent and the

chelate, respectiyely. Ka is the ionization constant of the reagent and Kf is

the formation constant of the chelate(4).

1.4.4. The separation efficiency of metal chelates:

The separation efficiency, of two metal ions at a given pH and

reagent concentration can be predicted from this equation,

L)2 tV|-(2) l^DMRh2 •• '

The . separation factoi; (3 is equal to the ratio of the distribution ratios

of the two metal chelates formed, with a given reagent. Since only K,-and

41 K|)MK,J will be function of the metal' .

The selectivity of an extraction can often be controlled by proper pH

adjustment. The separation efficiency depends on the relative formation

constants and on the relative solubilities of the chelate(4). One of the most

important application of solvent extraction is spectrophotometric

determination of metals in the visible region. Many organic reagents' form coloured chelates with metals but most of the chelates are insoluble in water. They are however soluble in organic solvents and can be easily, extracted.

Many complexes of metal ions with organic agents in aqueous solutions are coloured when extracted with organic solvent. The coloured

12 extract may be used directly for the determination of the concentration0 of

1 • • •

the metaj by spectrophqtometrie techniques. These techniques are"

particularly applicable with many chelate complexes(3).

1.5 Hydrqxamic: acids:

Hydroxamic acids may be regarded as derivatives of both

hydroxylamines and carboxylic acids)(6). The acyl portion of naturally

pccurring derivatives is usually simple and is often acetyl or originates

biogenetically from acetate*

Hydroxamic acid group containing compounds are intimately associated with iron-transport phenomena, the selectivity of this mechanism is critical since numerous metal ions which may not be essential or which may have a toxic effect on the organism are present in the environment'''0.!

Hydroxamic acids are also known as coristituents, of growth factors, food additives, antibiotic antagonists, tumor inhibitors, antifungal agents, and cell divisions factors, and several of them have been used as drugs'71.

Hydroxamic acids are also potent and specific inhibitors of urease activity, themolycin, elastase and aminopeptidase. These enzymes, are metalloproteinase and the mechanism of inhibition appears to involve chelation of metals, at their active sites161.

Hydroxamic acids have also received considerable attention as reagents in analytical chemistry for gravimetric analysis and for solvent^

13 extraction and spectropiotometric determination of metalsw. The reagents

•• '!\ >! . -,:i;- •&,-. '•'• .. . • ' v/ ' ' ..•• • .•.•.'•• •. • • • • (•' are also/useful jn. the aha^sis of trace metals by flow injection analysis'91 and high performance liquid chromatography(l0).

. Hydroxamicacid resins have attracted attention as analytical reagent, owing to their good complexing behavior with a broad range of metal

Discovery of oscillation phenomena in the fluorescence intensity of some aromatic hydroxamic acids suggests that they can undergo photochemical reaction(l2).

1.5.1 Structure of hydroxamic acid:

The term hydroxamic acid is given to the monoacyl derivatives of hydroxylamine.and indicates most readily a compound with hydroxylamide structure. Even though, the term hydroxamic acid is used commonly to denote all the monoacyl derivatives of the hydroxylamine there is tautomeric hydroximic acid with hydroxylimide structure .

14 • 9 ?H R-CNHOH R-C=NOH

Hydroxamic' ac;id Hydroximic acid

9 ?H Hydroxamic acids have the general formula R-C - N - R' where R and R' could be hydrogen, alkyl or aryl groups. In solution hydroxamic- acid exist in the two tautomeric forms shown in scheme (1). If there is restricted rotation about the C-N bond, then the both Z and E forms or isomersi of the ketoforms in scheme 2 exist'61. R OH R \ / . \ C-N •// ••- \ • /. c=n\ • O H HO OH

(1) (2)

: Scheme (1)

/0H

O OH O H

(Z) (E)

Scheme (2) •

NMR spectra confirm that both Z and E forms exist in solution, as do the enol forms (the Z - E isomers ratio is solvent dependant)(6).

15 The stabilities of the different forms of isolated and hydrated hydrbxamic acids ; were examined using abinitio molecular orbitafl calculations. The calculation show that the isolated E-keto isomers of formohydroxamic acid (Fha) and acetohydroxamic acid (Aha) have the lowest energy, however, oh hiydration the Z-keto isomers become the more stable forms due to hydrogen bonding'14'. The X-ray crystal structures of

(Aha) hemihydrate: reveal tha,t, its crystals',. the Z-keto form is present similar to the prdpionohydroxamic acid (Pha) molecule.

Structure (1), scheme (1), contains one easily replaceable proton, monobasic acid, while structure (2) may dissociate two protons thus behaving as dibasic acid. This keto-enol tautomerism provides a number of sites which are available for metal ion coordination16'.

The keto-form predominates in acidic solutions while the enol form is the dominant form in alkaline medium* 3>. There are several possibilities for the mode of dissociation, the mono anion of these forms give rise to quite equilibria and proton dissociation may follow the paths shown in scheme 3.

•16 p /0H R-C : v V R-C 2 (1) NHOH ( ) NOH

p :>p- •. XM /0H R-C :,..-•, : R-C R-C

\ •:•••••"• \ ; :; ^ NHO" N'-OH N-O" (3) (4) ^ (5)

O

../••• R-C V N-OH (6)

Scheme (3) ;

The possibility for the existence of the several different monoionic

forms .depend on the ligand concerned, it was suggested for example that

for benzohydroxamic acid (Bha), structures (3) and (4) occur in essentially

equal concentrations( r>" '.

Exner and; Simon concluded from IR and UV spectra that

hydroxamic acids with common substituents form N-acids practically

exclusively"6'. The existence of such forms in a complex may depend on

the metal ions"71.

Unless the structure is hindered, most hydroxamic acids will be

hydrogen bonded and exist in the keto form, bound to a transition metal via the oxygen atom"8"19'. '

17 R-N-OH R-N-O. r n I + M I R-C=O \ R-C = O

Thiohydrbxamic acid is obtained by replacing the oxygen atom with sulphur atom.

H-N-OH Tautomeric N-OH

R-C=S R-C-SH

1.5.2 Nomenclature:

The hydroxamic or carbohydroxamic acid nomenclature (IUPAC O II

Rule C-451.3) may be used for this structure R-C -NHOH for example

CH3CONHOH is named acetbhydroxamic acid and cyclo-C6H||CONHOH cyclohexane carbohydroxamic acid, and 1,4—d o (CO-NHOH)2,

1,4-naphthalene dicarbohydroxamic acid. Alternatively, compound of this type R-CO-NHOH, may be named as N-hydroxycarboxamides according to Rule C841.3. Monoalkyl derivatives R-CO-NHOR and.R-CO-NR'OH are, esters and N-substituted hydroxamic acids respectively forexample

CH^CO-NHOR can be designated as alkyl acetohydroxamate (or N-alkoxy acetamide) and CH;,-CON(R)OH as N-alkyl acetohydroxamic acid or

N-alkyl-N- hydroxyacetamide. Compouhds'of this tyfce R-C(OR')=N: OH and R-C(OR')=NOH

;: :(E) •;•;, r , (Z) •>;; ./;•• are best named as derivatives of hydroxamic acids, of this .type.

i '! j :,.-. - j • . •' .

R-C(OH)=NOH for example, (Z) C6H5-C(OCH2C6H5)=N-OH would be benzyl (Z) benzohyroximate.

The acyl derivatives' R-CO-NHOCOR' and R-CO-NCOR'OH can be . (i) (ii) named as hydroxamic acid derivatives e.g, (i) and (ii) as acetohydroxamic aceticanhydride and N-aeetylacetohydroxamic acid respectively.

CH3-CO-NHOCOCH3 and CH3-CO-NCO-CH3OH.

Dialkyl derivatives ,of this type R-CO-NR'OR" are referred to as hydroxamates (i.e. as esters) therefore this type CH3-CO-NCH3OCH3 can

2:) 26 named;methyl N- methyl acetohydroxamate( ' ).

1.5.3 Acidity of hydroxamic acids:

The unexpected relatively high acidity of hydroxamic acids of this form R-CONHOH is one of their most striking properties. A number of pKa values for RCONHOH are reported in the literature and these are of the order of 9, pKa units i.e. approximately 6 units more acidic than amides

R-CO-NH,(:il2|i). - • .' O 11

The structure ot the anion of R-CNHOH is still the subject of some controversy. Apriori, three possibilities, (3, 4,5) are considered in scheme

(3) and the ultimate aim is to determine the dissociation constant that is solvent dependant. Some. evidence has1, accumulated that structure 4 represents the

structure of the anion of (\). In any one series the acidity of either NH or

OH protons are( affected to some degree by the nature of the substituents

("" ). An IR study of a salt of deutrated benzohydroxamic acid in the solid

state and in dioxarie solution' detected 0-D vibration frequency which

could point to the presence of structure (4). An intramolecular hydrogen

bond may also contribute to the stability of structure (4) as in structure (7).

O"'

• •• / . \" ' R—C H \ /

' N

' (7)

An intermolecuiatiy hydrogen-bonded species (6) could explain the

stability of (4) and the fact that nucjeophilic reactions seem to take place prefrentially on the NOH oxygen atom. .

1.5.4 Preparation of hydroxamic acids:

Hydroxamic acids, aliphatic or aromatic, are in genefal prepared by the reaction of an activated acyl or aryl group with hydroxylamine in the presence of an alkali as a catalyst. '"•.'.

The methods, which are commonly utilized, for the preparation of the hydroxamic acids are:

20 (1) Blatt Methdd:

In this method alkyl or aryl ester react with hydroxylamine in the

.presence of alkali* %e acid was obtained by the addition of mineral

acid in appropriate quantity, to the solution'281. :

RCO2Et+NH2OH+KOH ,—-> RCONHOK+EtOH+H2O

RCONllpK+MX- > R- qONHOH + KX

(2)The reaction between carboxylie acid carbodiimide and hydroxylamine'" ':

RCOO11 -i R'N- ONiUNIkOil -> ; RCONHOI 1+R'NH-CO-NHR

(3)The reaction of acid anhydride and hydroxylamine:

v. The, iionhal product is Q, \\, diacyl hydroxyjamine :RCO;NHOCOR . which is easily hydrolysed by alkali to the monohydroxamic'241.

(RCO)2 () + 11-N11O11—-> RCO-NMOM+RCOO11

(4) Acid chloride procedure:

In this process, the hydroxylamine is acylated with an acid chloride to produce'hydroxamic acid derivative1""'.

RNHOI-l'-r-R'COCl -> RCONROH + HC1.

(5)The reaction between carboxylic acid and hydroxylamine in the presence of Ni as a catahst'"'":

RCOOH r H:NOI 1 -iii->' RCONI-IOH + H2O.

(6) Reaction hehveen isocyanate and hydroxylamine'271: i

R-N=C=() + Nil2Ol 1 ->R-NHC(=I!1OH)OH f==^ RNHCONHOH

•> i 1.5.5. "Reactions olf hydroxamic acids:

'mm ' 3t

In. many aspects the reaction of hydroxamic acids resemble those of amides. The presence of the second oxygen atom in O=C-NO.chain of hydroxamic acid, however, alters both their nucleophilic reactivity and the pattern of substitution compared to amide. Thus hydroxamic acids, are stronger hi&leophiles than am|de with (N)-0 atom being most reacti've'"Sl.

The acidity of hydroxamic acids may be attributed essentially to the

-OH group anil the basic characters of the- nitrogen atom is suppressed as it is in amide1""1. I lydroxamic acids are'very week, though several time stronger than phenol, the suppression of acidic character may, be attributed, to intramolecular hydrogen bonding as shown by infrared studies12'"1'.

(1) Hydrol>:sis:

Acid .or base caial_\sed hydrolysis of hydroxamic acids and their derivatives to carbox\lic acid and hydroxylamine derivatives proceed readily and il would appear logical to compare these hydrolysis with corresponding reactions of amides'711. ::

i ' . ' ; ••'',,•

Kinetic studies do indeed suggest that the mechanism of acid and base-catalysed hydrolyses of b'enzohydroxamic acid resemble those of

amides.; • . ;. • • • . "\

Attack b> water leads to a tetrahedral intermediate with a kind usually associated with nucleophilic acyl-substitution ' reaction, this intermediate yields the Una I products. (2) Alkylafibn 6f hydroxaink acids:

T • "••' •"• '•s ?'

Alkyl hydroxamates are/the major products, resulting from the action of an alkylating agent on ;the hydroximate ionl2(J|.

O O / M'RX" / C,,M5 - C -T-^ C6H5 - C + CfiH5 - C - OR . ; : • -MX • •••••• • ' N~OR : • : - N-^OR NOR

R1

TJiis; reaction represent alkyiatitin of the potassium and silver salts of the alkyl esters C(1HsCO-NHCR afforded a mixture of products from which the esters and the two isomers, were isolated(29). The ratios of the products are largely dependent: upon the solvent, • the nature of the accompanying cation, and the electrophilieity of the carbenium ion centre in the alkylating reagent'" '.

(3) Acylation:

The reaction of hydroxamie acids with acid halides or anhydrides produces ntixed: anhydrides (O-acslhydr'oxamic acid). More reactive acid halides such a^ sulfon\ I or phosphoryi halide induce an almost spontaneous l.ossen, •rearrangemeiu1'"". The mixed anhydrides are acids that form stable sails, which rearrange on heatintz. .0

R-C 0-COR \: / ,' Nil (0-Acyl hydroxamic acid)

(4) The reaction with Ire(111) chloride:

Hydroxamie acids containing CONHOH group react in weakly acidic solutions with l:e(lll) uivirm red or bluish red inner- complex salts;

•' ' '•''.' '•.'• . ... > " .; "• • ;< •

This reaction' provides a method for detecting the ester group-alter conversion to hydroxamic acid group1 lh.

N -- O: Fed II // R -C Fe

v. .0

'This, reaction is a characteristic of enol form,: R-C(0H)=N0H. of hydroxamic acid1'"'. The most probable structure of the l'e( III )-hydro\amate complex is the metal ion chelate shown in the following formula proposed'by Werner1'1".

2) R NH

O

Fe O

HN Ol C-R \ C I I R . Ii

, In aqujeoUs [solutions a maximum colour yield is obtained at 8 C but lor best reproducible result room temperature is used. In non-aqueouijp solvents the reaction is slow and higher temperature must be used for a short; period', since prolonged healing causes, decomposition of the hydroxamate. The two most important factors for the formation and stability of the complex are pH and the excess of Fe(III) ion used. After the' formation of the liydroxaniate anion the solution is acidified and Fe(Ill) ions are added in excess to produce the coloured complex.

3RCOMHOH+ •r'3 1-1,0: Fe(RCONHO).^ + 3H,0' Red-violet A considerable excess of Fe(lll) ion concentration is required t^i;. atiaiiiinen.t of maxmiuhi colour .yield, but the blank, values also increase, limiting the accurac} of determination of small quantities. The pH of final solution has a profound effect '' '. It has been shown that there are three possible coloured fCOiTiplexes depending on the pHk>f the..;.final solution as

shown in equation below.

(i) \ih), (in) (iv)

As shown by the above equation, a very high acidity will inhibit the

formation of the, complex ;.and favours the existence of free hydroxamic

acid, (IV). Structure (I) is lfeddish-brown and is formed at weakly alkaline

solution. Structure' (11) is chery-red and is formed at weakly acidic

solutions, while structure (HI) is red - violet with an absorption maximum

at 520-540 nm, is formed at relatively strongly acidic solution (pH 1-3) has

the greatest stability of the three, and is the preferred complex. This

indicates that metal cation complexes or chelates are produced in steps and

that the formation, composition, and stabilities vary with the pH of the

reaction mixture. \

>••••'. f- •

(5) Nucleophilic reactions:

These are complicated by three possible sites of substitutions, N, N-%

andC=O.

Treatment of a series of hydroxamic acids with mesylchloride

(MesCl) in'the absence of 2-equivalent of triethylamine at 0C gives

2-chloroamide'in good yield1 :>). O . R,

Et3N . .: N MSCl I : OH . OMS ; ci H

The treatment of hydrbxamic acids with trietliylarmne and methene sulphonyi chloride at 0°C in dichloromethane gives 2-chloroamides.

O 1 EtiN Ri MeSCI I

Ql- '• i OMeS Gl

Also the hydroxamate anion is particularly effective nucleophile for attack on the phosphorous atom on phosphoric antfphosphonic anhydrides and halides. In the reaction between hydroxamate unions and phosphohofluoridates was suggested that the very high reaction rates might be due to participation of hydrogen bonding in the transition state.

R—(T-H3H O R—C—O" - F I NO' . NOH O

(5)With hypobromo.us acid:

Monohydroxamic acid gives carboxylic acid.

2 RCONHOH+2HOBr 2RCOOH+2HBr+NO-»+H-,O (6)ThreTLossein rearrangement:

In 11! 872fjjbossen discovered the- rearrangement bearing his name,

when h^'Qh'^lveclUhatipyrolysis of the mixed anhydrides affords phenyl

isocyanate(";^f^Byl derivatives only undergo pyrolytic rearrangement under fairly str|rigtr)^t)nditions<30).

The"usuaj-fTidHe jis to convert this compound

O // R-C ' \ • N - OCOR1

by mean of; strpng base at tow temperature into the salt, preferably in a medium from which the salt precipitates. Rearrangement of the anion to the isocyanate involves the departure of the carboxylale ion with concomitant migration of the group R from carbon to nitrogen

M

> • > R-N = C = CH RCO2"M

N-O-COR1

i I he rate of this base-catal)zecl-l ossenrearrangement depends upon the electronic nature of the group R and R'and the strength of the acid

K'l 0,11'r> 1.5.6. Detection of hydroxafhic acids:

Hydrpxamic acids can'.be identified by their characteristic reaction with Fe(III) ion which yield an intense dark violet complex and they give green blue insoluble Cu(Il) salt when reacted with copper(II) acetate, also hydroxamic' acids, react with vanadium(V) and give a violet complex when extracted into chloroform12"1.

1.5.7. I.R. of hydroxamic acids:

The JR. spectra of hydroxamic acids and their complexes are generally very complex, though some characteristic bands for different l.igand structures were suggested1'8'. For example,the band around 3200cm"1 was assigned as the NH valence frequency, while those observed in the

3080-3060 cm"1 region were attributed to the NH deformation and the CO valence vibrations1:^. The broad band around ;l 610-1585 cm"1 that was observed in,.metal free ligand is assigned to the ketonic carbonyl vibration.

Its broadening originates front intramolecular hydrogen bonding an^l it undergoes an energy shift of 40 - 60 cm"' when ketonic oxygen coordinate to the metal ion,'1*1. The band containing the NH planar deformation 60% and CN valence 40% frequencies is centered at 1 570 cm"1, there is also a band around 1400-1440 cm'1 that is also assigned to the deformation of the

NH moiety1'?1. ' :

When hydroxamic acids are dissolved in an inert solvent the band at

2770 cm"1 disappeared and a new one is observed at 3420. cm"1, also the band at 3280 jcm"-I is shifted to 32^0 cm" . These vibrations, suggest that

there are considerable changes in the hydrogen bondirfg system during

dissociation process'61.

These IR results are supported in several cases by NMR spectra. The

NMR spectra of acetohydroxamic acid exhibit the signals of both NM and

Oil protons while in the case of sodium acetohydroxajnate only the NH

proton is observed in the spectrum'61. This result supports the existence of the structure^)'in contrast the NMR spectra of the Sn(IIl) acetohydroxamic acid complex show the presence of onl\ the OH proton'6'. Thus the NMR spectra results seam to exclude structuie (3) in the complex system while the |R measurements suggest the formation of structure (6) in the Sn (III) acetohydroxamic acid species'6*.

O" ] / R-C X N - OH

(6) •'

i .5.8. MonohycJrovattiic acids as a ligand:

The single h\dio\amic acid group behaves as a typical bidentate donor towards various metal ions'" Mo'nohydroxamic acids form octahedral complexes with a number of

metal ions, coordinating via two oxygen atoms of the deprotonated

hydroxamic ,acid group. This has been proved in the X-ray studies of the

trisbenzohydroxamic acid iron(III) dihydrates(j9) and trisbenzohydroxamic

acid chromium(III). Brown also, studied solid state complexes of Fe(III),

Co(U), Cu(II) and'Ni(II) with monohydroxamic acids'7'. Infrared spectra

showed shifts of about 40-60 cm'1 in the broad bands at 1619-1583 cm"1

when compared with the free-ligand. This strongly suggests the

complexation of the ketonic oxygen atom. Bands in the 1445 cm"1 region

can be assigned qualitatively to the, N-C stretching vibration with

contributions from C-0 and C-R modes those at about 1300 cm"1 to C-R

stretches and that about 1100 cm"1 to a practically pure N-O stretching

mode"11 The general pattern of the infrared spectra support normal coordination via the ketonic oxygen atom and the oxygen atom of the deprotonated NHO group'h).

The s'olid state magnetic moment and the electronic spectra provide

further support for an octahedral structure in the case of Fe(III), Co(II) and

Ni(ll) complexes and close to tetragonal geometry for the Cu(ll) complexes'6'. The' hjghbr Stabilities of the propionohydroxamic acid in comparison with those obtained for the acetohydroxamic acid presumably reflect the greater inductive effect of the ethyl group compared with methyl group1381.

In solution, equilibrium studies of monohydroxamic acid, with

Ni(II), Cu(II), Zn(II) jand Fe(III) have shown the existence of different species. Except for Cu(II) the acetohydroxamic acid, propionohydroxmic acid (Pha) and benzohydroxamic acid (Bha) ligand, were assumed to form octahedral complexes as found in the solid state.

According to the Irving-Wiliam series Ni(II) complexes should have higher stability constant than those for Zn(II) complexes'61.

1.5.9.Aminohyclroxamic acid as ligands:

Amino hydroxamic acids are sub class of hydroxamic acid driving from amino acids'of general formula' '.

NH, : / R - Cl 1 \ \ O=C-NHOH In the a-amino h\dro\amic acid, the amino group NHi and the ' V hydroxamic acid group - C-NHOH are in an a-position with respect to

l(l1 each other R-CH-(NH2)CONHOH .

These compounds possess the abilits to form two types offive- membered chelat'e either via -their nitrogen atoms or through the hydroxamilc oxygens and consequently they, are good ligands for various

metal ions(6\

1.5.9.1.Acidity of aminohydroxamic acid:

Hydroxamic acid derivatives of aminoacids have even more complicate abid base properties than the monohydroxamic acid. The most simple aminohydroxamic acids (alky! aminohydroxamic acids) can liberate two protons, one from the protonated amino group and one from the hydroxamic. acid group161, some contradictions,.concerning the acid- base chemistry of aminohydroxamic acid were published where, some have concluded that the NH/ is more acidic than the hydroxamic function'61.

While others have taken the opposite view"'". Very recent papers suggest however, that the two protons of aminohydroxamic acid ligand be liberated in overlapping process. This means the pKa values determined by the pM- metric titrations can not be unambiguous! v assigned to either of the proton dissociating group. Instead, only the dissociation microconstants are characteristic for the real acidity of the individual ligand function'61. The dissociation microconstants were determined for L u-alanine hydroxamic acid (L u-Alaha) and [3-alanine h\droxamic acid ((3-Alaha)(42) by pH-

C|lNN'1R spectrophoto'metric titrations.. According to the results obtained, the NH? group of the u-derivatives is some what more acidic than the

CONHOH function while the acidit\ sequence is: opposite for the p-deiivat||fej;,.(p|iese results are in gopd agreement with the results obtained

<6) by Karji^fS'|g^'j,Hi91 a^ek from UV absorption curves .

;Re|ejillyV;|HlNMR measurements published for sarcosine hydroxamic acid (Sarcha) indicated that the CONHOH group is more acidic than the amirio group. The differences in the results published for (a-Alha) and

(Sa'rcha) 'can be explained by electron donating effect of the methyl group(6).

X-ray resujts indicate that the above situation is not the same in the solid state.' Namely the structure of glycinehydroxamic acid (Glyha) reveal that the molecule crystallized from neutral solutions contain a deprotonated hydroxamic nitrogen while its amino group and hydroxyl oxygen remain protonated.

1.5,9.2Complesx formation ami analytical application of aminohydroxamic acids:

Most papers dealing with complexes of aminohydroxamic acid, related to u-derivathes"1'1^ and specially to derivatives of simple amino acid, such as (Glyha)'K1"' and (u-Alaha)141'4'-"1*1. Iron (III) was involved in most imestigaiions because of the well known biological importance of

l 1 • i in i i i ll > -I ' it>> r^ 1 ,-p /n, (42-44. 16 the l-e(lll)-h\droxamate complexes . Complexes of Cu(II) S| <:' Ni(II)'r' u<" Ml Coill)1*' " '"• r- •'" and Zn( I l)(5a55) ifave also being in\ L'^

u Very few papers contained results on Mn(II) and Cd(II)-Glyha and

« • '- a • • ' a-Alaha complexes of these metal ions were found to be rather week and precipitation usually occurred at pH'values around 8. Glyha and a-amino hydroxamic acids which contained alkyl side chain, (a-Alaha), norvaiinehydroxamic acid (n-Valha), norleucinehydroxamic acid

(n-l.euha), (Sarha)) form complexes with identical stoichiometry and also the same stabilities"". •

The lower stability of (Sarha) complexes may be due to the steric hindrance caused ,by the

1.5.9.2.1. Copper complexes:

The first reliable result of (Glyha) complexes of Cu(II)' were obtained by Paniago and Carvalho who reported the formation of four different complexes including dimeric species whose existence was supported by ESR spectra1'"1,

Potentiometric data were used to evaluate the stability constants for the four 'proposed species. The authors suggest that in the dinuclear ' complex (Cibl-2H.|), the Oil group acts as the bridging group between two metal ions. l.eporati, in his later work, suggested an even more complicated

1 i 1 "* i equilibrium containing a tetrameric complex (Cu.(L.O , however, in most papers the authors have returned to the suggestion of a hydroxo-bridged

binuclear spe6ies(6). Although the fact that tlie dimeric species exist at low

• pH(4 5—5.5) is strong argument against a hydroxo-bridged species'6'. A

similar bonding mode was suggested earlier by•Karlieck and Polasek1''1 for

the Cu2L: species formed in the Cu(II)-Glyha system for which

spectrophotomeric measurements were used to study the solutions with

copper(II), These reports indicate that the complexation of copper (II) to

a-aminohydroxamic.acids start via (N,N) as well as the (0,0) donor set"".

This mixed, bonding mode changes at pi 1 (4.5-5.0) and above p'H 6 the

1 1 >l) only species is (CuL2) complex with 4N coordination * ' . This latter

coordination "was supported, by X-ray studies in the solid state1'11061 and in

solution by visible absorption (A. max 535nm) and ESR spectra (g = 2.212,

A- l95G)(5a54>. •

Increasing the pi I .up to 9-10 in these systems causes a new

consumption of base, this process can lead either to the formation of

a mixed species or to the deprotonation of the bound hydroxamic acid

group NOH in the (Oil.2) species. According to the spectrophotomeric

studies this proton release dose not change the 4N-coordination mode

although the d-d transition shift from 535 to 500 nm. The result suggests

thai the stoichiomeiiA of (CULTH-I) species formed be most possibly as

siiven in'the scheme (4). CH3—HQ—£=0 ;j U N. N—-O- "i \ 7 . Cu H 7 \ '• •' H2N N — C '

CH3—HC — C = 0

Scheme (4)

This coordination mode is specially favoured when square planar

complexes are formed(6).

1.5.9.2.2.Nickel complexes:

\ Brown and Roche have made a comparison between the complexes

formed in the Ni(II)-alkylhydroxamic acid • systems with those created in

the Ni(II)-aminohydroxamic acid system, both in the solid state and

solution'4'1". Aha, Pheha, Glylia, and seriiiehydroxamic acid (Serha) were

all involved'6'. Obvious differences were found in the solid state,

e!g. in colours and in solubility properties. For instance the

bis-(aminohydroxamato)-nickel(II) complexes were red or orange in colour

and diamagnetic planar species, while the bis-(alkylhydroxamato)

nickel(II)-dihydrates were green octahedral. The complexes formed in the

former system were insoluble in most solvents while those of the latter system were freely soluble in polar solvent"11.

37- The infrared spectra df bjbth bis-(arninohydroxamato)-nicke|(II)i

T * i •••'•?' complexes in the 3.000-3500cm" pregion were -,\£rv.' 'simifar to those of corresponding amino acid complexes'"'.Three bands assigned to NH stretching mode are observed to shift on deuteration. These results suggest the coordination of the amino group to the nickel ion(:o). In contrast to the bis-(alkylhydroxamato)-nickel(II) dihydrate the carbonyl frequency in bis-

(aminohydroxamato)-niekei(Il) remains unchanged at 1604cm"1 when compared with free ligand i.e. the ketonic oxygen is not involved in coordination to the nickel atom X-ray crystallographic studies of the bis-(jlyha-niekel(II) complexes and bis-(Alaha) nickel(ll) complexes showed that' these complexes are square-planar with trans-geometry and coordination via both the amino and hydroxamic nitrogens1'1"1.

In solution depending •/•on.pi i different species are formed'" o()1. The spectrophotometric results show a square-planar environment to be present in all complexes except the Nil., species and the complexes

4N coordination. The energies of the d-d transition are characteristics- for: square-planar donor set of four nitrogens, the charge transfer pattern resemble ver\ closeK as reported for Cu(II)-Alaha but the.respective C.T . transitions are shifted about 50 nm to shorter wavelength'61. This shift of energ) \alue is normall> found in Ni(ll) complexes when compared to the concsponding Cu(ll) species. At a high pH9 a new base consumption process similar to that for Cu(II) containing system begins. Deprotonation of (NiL2j complex lead to the formation of ifie (NiLoH^) species. X-rays supported these studies of the (NiL2) complex.

1.5.9.2.3. Zinc(II) and cobait(II) complexes

Both metals ions begin complex formation with a-aminohydroxamic acids at pH5.5, Precipitation however, occurs at pH8 even at a metal- ligand ratio 1:5 (47o°. The species most generally found are the ML and

ML2 complexes. The species formed just before precipitation (MLH i) and

(MLiH-i) were suggested to be mixed.complexes'6'.

In a new study proton NMR spectra were used to establish the binding mode in zinc(Il) aminohydroxamic acid complexes, it was shown that the zinc(Il) ion initiate its coordination via the two oxygens of the hydroxamic group. An increase in pi 1 leads to deprotonation of an amino group and the formation of binuclear (ZiiiLO species together with (ZnL)

1 and (Znl.2) complexes. A mixed type of 'coordination via both nitrogens and oxygens was found in the (ZmL;,) and (ZnL:)1'1'-

1,5.9.2.4: Iron(IIf) Complexes:

The pH-metric and spectrophotometric results agree that there is a measurable complex formation between iron(III) and a-aminoh\droxamie - acids even .below pH2INIX>l. Very surprising bonding modes were presented for the iron(III)-Glyha and iron(IIl)-methion;inehydroxamic acid complex15"1, instead of bonding by two hydroxamate oxygens the involvement of the amino nitrogen and the OH oxyggtvof the bydroxanpate groups coordinate^ to the iron(III) lias been* proposedl rThei.proposal was-baseu oh the -formation,of an unprotQnateel species (ML) -in the acidic region. However, it seems more likely that the iron(III) as a? typical hard acid would favour an oxygen donor rather than the very.basic! nitrogen. On this basis, it is clear that (ML) is in fact a {M (HL)(OH)) complex i.e. a mixed hydroxo species with the amino group in .-itjsi .protonated form. This . assumption is supported by the, spectrophotometric result1M ".

Species distribution is shown as a function of pH and stability constants of the various species for the system Fe(III) Glyha'^'1. The following species were delected for l;e(IIl)-gl>c\ Iglycinehydroxamic acid

(56) system M2L, ML2H2, ML3H2,' ML, ML,, ML2(OH)2 , where

M represents metal ion and L to the ligand. In the acidic region a protonated complex ML2H2 of the complex M2L are formed, which

1 predominates and no iVH. was detected'" "'. The protonated ML2H2 species which is represent at pH5.7 is precurso of 1:2 complex and the protonated

M1.-,H2 complex which is the .major complex species in equilibrium at close to neutral pH(5f1).

In the alkaline region the major peaks belong to the ML.^ species, above pH9 there is evidence for the formation of the hydroxo complex.

Visible spectra of the le( IIl)-Glyha was recorded in the pH range 2-8, the system shows absorption bands in the 380-430 nm region. The colour of

40 the solution changesTrcfyi-wine-red at pH2-4~6 to o&nge-rerd at pH 5.2 and

• - *

<:>6) remains this colour to pH9 .. (

The magnetic moments of solutions of Fe(III)-Glyha are given as a function of pH for various ligand metal ratio"'". The species distribution and the stability constants are given for this system Fe(III)-Glyha system. in the normal aqueous titration range, Glyha can liberate two protons one from the protonated amino-groupjMH}" and the other from the OH group of the hydroxamic group" M.

The presence of the u-amino group in the aminohydroxamic acid increases the acid character of the Oil group in comparison with that of acelohvdroxamie.acid.

41 Maximum

Species Structure ,r proton Liberation H

O — N II. / \ Fe C =0 \ / NIL — CI I.

II .

N—0 N

••([., Fe-. C = 0 4 \ CU> N1F Nil,—CH,

— 0 N \ c==o

Nilr—CFF

Fig.l Proposed structures of \arious species in the Fe(M)-Glyha.

The distribution shows various species iVIill^ ML, MHL2 and occurring in the pi I region 2.0-6.8. For the species ML3 at pH8 approximate!) ^o proions ' Mo! o\~ Fe( 111) are liberated and at this pH the

N11O11 hulrowl has lost 90% of ii is proton, above .pH9 two hydroxyl species were detected

•1: The potentiometric, spectrophotometric and magnetig&studies show that Glyha coordinates to F.e(III) via the p.-7-amino nitrogeitand the hydroxyl oxygen of the NHOH group'10'.

1.5.9.3.The effect of side chain donors on the coordination of a-amino hydroxamic acid:

'The results obtained for Serha and threoninehydroxamic acid (Thrha) both containing an alcoholic OH group tyrosine hydroxamic acid (Tyrha) and 3,4-dihydroxyphenylalaninehydroxamic acid (Dopaha) (with one and two phenolic 011 group respectively). Methioninehydroxamic acid (Metha)

(with thioeiher sulphur) and histidine hydroxamic acid (Hisha).(cqntaining an imidazole nitrogen atoms) have shown that there is no involvement of the side chain donor in Serha, Thrha, Tyrha and Metha. These ligands coordinate in a manner similar to the simple u-aminohydroxamic acid'61.

Dopaha contain two phenolic Oil groups in the ortho positions which lead to -the possibility of the formation .of catechol-like (O, O) chelaies in addition to the normal ami no - hydroxamate - like chelates.

Results published for the nickel(II) and cobait(H)-Dopaha systems indicate thai phenolic groups do not affect the coordination equilibria throughout whole pH regions studied1'".

The electron paramagnetic resonance EPR measurements performed on the copper(If)- Dopaha complex indicated the formation of4N species ut pi 16.9I"V'. A now species however, with less than four coordination was

43 formed' above pH9.5. Comparison of these result^ with those obtained for

•.. •••••% the Copper(II)-(Dopa) .system'?-7?, suggest the formation of.& catecholate- type (O, O) chelate. The comparison also shows that in the case of copper

(II)-Dopa the catecholate- type of the chelation begins to form at pH 5-6, while in the copper(II)-Dopaha system this happens only above pH9.5p7).

This difference indicates that copper(II) ions form' much more stable species when bound io ami.nohydroxamate moities than in the parent ami no acid donors1"1.

One of the most effective sides chains for binding of metal ions is certainh the imidazoie moiety of histiciine residue141'49'. Studies on the coordination modes of histiciine hsdroxamic acid (Hisha) with Cu(II) ions suggest the Jbrmation of dimeric species (CibLiH-l.), (CibLiM) and

(Cu:l.:).

The remarkable fact that coordination of Hisha began at a lower pH compared .with previously mentioned.aminohydroxamic acid ligand suggest different coordination pattern i.e. also involving the imidazoie nitrogen owing to its favourable pka value'. ''.

The complexes formed below pH 9 in the nickel(II)-Hisha system are octahedral or pseudooctahedral. The first species (MHL) start form at pH3 this is if chelate complex witn (Nimi,NH;i) coordination mode t and octahedral geometry this coordination mode was also proposed tor (Nil.^llj) complex which is the second species assigned in the Ni(il)-Hisha<431. Deprotonatibn of. fhis complex lead • to the foKijnatioaof,

•:'-.••' >' '••','•.•'. : • :••.-•'• •"*';*"1

(NiL2H) and (NiL2) species which: are also octahedral but-with different coordination modes. The ligand molecule becomes tridentate with

(NH2, Ni,,,,N ) 5 binding set. Again the octahedral symmetry of the (NiLi)- complex is favoured by the presence of the competitive side-chain donor.

The required stability being provided by involvement of the imidazole ' nitrogen"". ;; ; • • .

pH-nietric and proton NMR measurement performed on zinc—Misha system have shown that zinc ion coordination begins in a similar manner to the simple u-derivatives( "^ '. i.e. via two hydroxamic oxygen and with the formation of (Zni.H) and (Zn(LI-I)i) species'61. Above pH6, one of the

•bound ligand "molecules become completely deprotonated to form (ZnLiH).

'The proton NMR indicates the presence of the other binding mode, which exhibit its own spectrum that . is present until precipitation occurs around pi I1)"". The pinion chemical shift strongly suggests the involvement of the aminu nitrogen in the binding of the metal ion. The change in the chemical shift of the imidazole protons indicates the additional involvement of the imidazole nitrogen. Strung support for the tridentate coordination of the ligand with the imokemeni of both aminoand imidazole nitrogens, plus one ul the Indroxamic gty-up donors, is obtained from consideration of the rotamer population of free and bound ligands|6). The strong stabilization of the rotational isomer, in which three donors (Njm)Namjno and the hydroxamic group) are in the gauche position with respect to each other may.be used as'

• •*,••• • • '- •

. • •"" , > . , " . : ' • • strong evidence for the tridentate coordination of His residue toa Zn(II)

(hi ion . 1.5.9.4.Derivatives of f^-aminohydroxamic acid:

When the NHi group and the hydroxamic acid group are in the

(3—position with respect, to each other, then the amino and hydroxamic nitrogens may form a six-membered chelatej-ing'61. Its formation is usually less favourable than that of the iive-membered ring possible in the case of a-derivatives1'". Thus the five membered c.helate rings created by two hydroxamic oxygens may be more competitive 10 (N, N) coordination in- complexes with •[3-analogues">l. This fact can be clearly seen in the

Cu(H)-(i -Akiha system in the1 scheme-.(5).

•if) Hp

I Cu N " o \

i .- 7 • \ c - O CM \ y N \ x '• y A Clio M \.' Cu

\ K O O - Cu c I I ' \ I •N 0 Cu n c -•

1-lK Cl H. Above pH5, the solution becomes grefen in colour, tin addition iQ the V: t '' ? ' " • ' ' ' ''•„'. « > . ' •' •.?'.•'." nitrogen atoms the oxygen atoms of the hydroxamafe moiety must;a/s?)'be

involved in metal ion coordination in the copper (II)-|3-Alaha system'61.

Above pH'4.5 the c\-6 bond at 616 nm increases in intensity until pH 9

without any change in energy. This indicates that the major species in the

pH range 4.5-9. remains the same. In this pH range, the major complexes

do not show any HPR spectrum at liquid nitrogen temperature or above. Thus, the coupling of the copper(ll) centres through NO group is very strong and it is now evident that the earlier suggestion for its involvement

in dimer formation in the Cu(li) a aminohydroxamic acid system was correct. It is ob\ ions from potentiometric results that only one complex species exists at a measurable concentration in the pHrange 4.5—9 whifch is

pentanuelear species (CIK-L,1 l.() and it is confirmed by an X-ray structure

determination'01.'

The coordination power of ligands which form complexes with

sequence of alternating ii\ e- and six-membered rings is greater than that of

ligands- forming complexes with all five-membered and six-membered

chelate rings'6'.

1 ha iinusuall} stable pentanuclear species exist in solution with a ver> high stability constant and contains 12-five and six-membered rings of different conformation"11.

48 'A similar situation to; that discussed'also exists in the aspartic acid ••'•... i ''. ' ' '" ' .••.',. .* derivatives7 having a ,. hydroxamic group placed, at the (3-cafboxylate

'(aspartic-p-hydro/amic acid (ASP-j3-ha) copper(II). Coordination starts at ground pH3. 7 'ie BSR species clearly indicated the formation h jpf an oligomeric structure with strong metal-metal antiferromagenetic

!coupling°'S). Abo- -t pH6 the only oligonuclear complex present in solution

is (CU.1L4 H.V^'sp-r-ies, the stability of this complexes is very high. '

In addition to (N, N) and (O, O) coordination, glycine type chelates

can also formed 1 the stable oligomeric structure161. Some very interesting

results were obta. '-ed with nic|el (II)—(3—Alaha system. Complex formation

begins at pH5 • ith the formation of an octahedral species. However,

'visible spectra slv A the geometry to be planar above pH7 when the .species

1 '(NiL>H) (Nil.-.) -.."'d (NiL2M-i) are formed"' . The most probable bonding mode in the I.: ."-^i" complexes is 4N coordination. Steric factors may additionally lave.-1' the planar .geometry with the two six-membered (N, N) chelate rings"".

•.' The Asp-j:- -ha ligand can also form glycme type chelate in addition to those created \> -h (N,N) and (0,0) donor. The visible absorption spectra indicated that tru- Ni(II)-Asp-(3-ha complexes have octahedral geometry knd the stabilit> -onstants of these complexes strongly suggest tridentate coordination thr -''Ugh the carboxylate oxygen, amino nitrogen and t hydroxamie nitrogen of the Asp-|3-ha(6)

-49 Ad interesting bonding mode was proposodJ'for the complexes • •..•. . » .. if- " »• 7- • <& formed in iron(JII)Msp+-J3-hak- solution*3**. ,The higher stability of the** complexes. and the considerable resistance' to the hydrolytic process suggested the presence of (O, O, O) type coordination via hydroxamate and carboxylate oxyge'hsp81. ..

1.5.9.5.Complexes of 5-aminohydroxamic acids: '

Glutamic-6-hydroxamic acid (Glu-8-ha) form the same type of coordination^ thai found for.' (Asp—(3—ha). The resulting (N, N) chelate, however, is a seven membered ring. As the seven membered chelate is much less stable than'the six membered one, the Glu-8-ha essentially has • two separate metal binding sites, i.e\ the glycine—like and the hydroxamate- like despite this separation (O, O, O) coordination with the hard iron(III) was found to be likel\. Nickel(II) and cobalt(Il) ions on the other hand were found lo prefer a gl\eine type of coordination, which is stabilized by the week coordination of the hydroxamate nitrogen"'1.

The stabilities of the Zinc(lf) complexes are higher than those proposed. This ma\ be explained b\ assuming that both (glycine-type) and

(hydroxamale- type) chelate complexes are formed in Zn(II)-Glu-8-ha

. I'M system .

1.5.10.The Biological activities.of'hydroxHmic acids:

The isolation of several naturally occurring and the synthesis of a number of medicalh acti\e hydroxylamine derivatives have stimulated 1 50 .recent progress-in hy'droxasnic acid chemistry.;I$otable;among these are the

• •:• >•••••• ' • v$r • antibiotic-cycloserine,the * antitumor antibiotic • . hadacidinef and heteroaromatic antibiotic aspergillic acid(IA). A series of.o-, m-; and p-alkoxybenzohydroxainic acid was found to be highly effective against pathogenic fungi1'16' while salicohydroxamic acids and their derivatives are effective antibacterial ,and antifungal agents. J3-alkylaminopropiono hydroxamic acid show hypotensitve properties and a number of hydroxamic acids possess hypocholetermic activity, p-butoxyphenyl Aha

(Bufcxamac) is in actual use as an anti inllammatory agent in humans(>fl).

A series of letraphihalohydroxamic and other dicarbohydroxamie acids have been investigated as potential antim.alaria'1'"'61. In quantitative structure, activity relationship study Lising the Marisch approach, a series of aliphatic and m-and;pfiubstituted benzohydroxamic acid were investigated for their relative power to inhibit urease activity. Among the alky I hydroxamic acid, maximum activity was observed with heptanohydroxamic acid, which the authors attributed;to slereospecific (hydrophobic bonding). They concluded that electronic effects do not play a significant role in this activity.

However, the steric effect of a bulky substituents becomes very clear in a series of aliphatic hydroxamie acids in which a phenyl group is moved along the fatty acid chain, a remarkable decrease of inhibitory power is )bser\ed as the phenyl group, approaches the hydroxaniic acid group (36) .

51 .1

Hydroxamic acids ifire generally potent and specific 'thhibitqrs 6f

urease activity and have become important as therapeuticals in the

• • •• [• '' ••>'•' -•'*'•• •.

treatment of hepatic coma iriiaddition to the trace elements that are required

for may forms of life and the essential elements'4'1.

Although the mechanism of the clinical1 action of the metal ligand

complex is not completely understood!, there is a reason to suspect that the

interaction of these ligands with metal ions is significant'4'1. Hydroxamic acids number among a few compounds effective as nucleophillic reaclivators of sarine- in activated chymotrypsin or aeetylcholine estei'ase '' N ' Experimental and Results

2. Experimental:

2.1. Instruments:

1 - Melting point apparatus.

2- Perkin Elmer UVVVIS spectrophotometer model 55OS.

• 3- Perkin Elmer IK spectrophotometer model 1330.

4- Jenwary, pi I meter, model 3030.

5- Rotavapor R 1 10.

2.2 Chemicals:

All reagents were Analytical Grade unless otherwise stated. The water used was double dislilled deionized.

1 lydroxylamine h\ clrocliloride, sodium metal, absolute methanol, absolute ethanol, anthranilic acid, diethyl ether, sulphuric acid 98%, hydrochloric acid 36%. nitric, acid 70%, sodium carbonate, sodium sulphate, potassium, carbonate, triethylamine, hydroxylamine sulphate, phthalici anhydride, glacial acetic acid, N,N-dimethylformainide (DMF), dicyclohexylcarbodiimide (DCC). sodium hydroxide, triethylamine (TEA), were purchased from BDII.

2.3 Synthesis of reagents:

2.3.1 Synthesis of anthnj&iilie hydroxamic acid:

The anthranilic hydroxamic acid was prepared by two methods: 2.3.1,l.Froni the^ coupling reaction of N-hydroxyjphthalimide t ; i • \ • • •

anthranilic acid ester, and hydrox^lafnine:

2.3.1.1.1 Preparation of N-hydroxyphthalimide:

To a vigorously stirred solution of hydroxylamine, prepared by adding a solution of 4.8 g (0.065 mol) of sodium hydroxide in 15 cm3 of distilled water to a solution of 9.9g (0.065 mol) of hydroxylamine sulphate in 15cm1 of distilled water, was added 14.3^ (0.065 mol) of finely powdered phlhalic anhydride.; As soon as crystals began to appear, an' additional 35cm' of distilled water was added and the mixture was allowed to stand over' night at 4°C. The solid was filtered, washed with ice water containing a little acetic acid, and air dried.

The experimental weight was I2.7g (80%). The product was reci\stallizeci from acetic acid to give needles colourless crystals, m. p. 235

(lit. 235- 237CC)

2.3.1.1.2. Ksterifiea'tion of anthranilic acid by N—hydroxyphthalimide:

10 a solution of ().c) g (0.05 mol) of anthranilic acid and 9.9 g

(0.0b mol) of N--h\drox\phthalimide in 150cm1 dimethylformamide

(DMlj, was added lOg of dicyclohexylcarbodiimide (DCC). After 4 hours< of Mirring. the dic\clohe\\ lurea (DCU) was removed by filteration. After additional 4 hours, more (DCU) was precipitated and this was also removed. The reaction mixture was concentratedon a rotary evaporator and

• 5-1 the N-hydroxyphthalijTJrderijister wa's precipitated by the addit-ian of ether and alcohol16"": - -

C NOH N—O—G

The product was pale yellow crystals from melting point of 180 C.

2.3.1.1.3.The coupling reaction between the iN-hydroxyphthalimide and hyilroxylumine:

To a solution of 4.2 g (0.06mol) of hydroxylamine hydrochloride in

30cm1 DMF, was added 6.07g (0.06 mo I) of triethylamine (TEA) and the

•mixture was stirred for 15 min and then cooled. The TEA hydrochloride was removed by filteration. and the filtrate was added to a solution of 5.5g of-N-hydroxyphihalimide ester in 120cm' DMF. After stirring at room

. I • .. . temperature for 4 hours, the solution was concentrated on the rotary evaporator. On addition.of diethyl ether, a pasty brown solid was produced,

It gave the general test of hydroxamie acid with Fe(III) and. V(V) solutions'" '. 0/

+ H.NOH H 0r\ I +

The purification of the product, had proven to be difficult, so another method of preparation was tried.

2.3.1.2.Preparation of anthranilic hydroxamic acid from the coupling reaction between methyl anthranilate and liydroxylamine:

2.3.1.2.1. Preparation of methyl anthranilate:

7.0g of ant h rani lie acid were dissolved in 50.mol'of methyl alcohol, which has previously been dried over dry pottasium carbonate. Dry hydrogen chloride gas was passed in Until the solution, which became hot, was saturated. Then the solution was retluxed for one hour. When the solution was. cooled, to . room . temperature methyl anthranilate hydrochloride crystallized. The mixture was diluted with about 200 cm"'of water and made alkaline with sodium carbonate. The oily ester was shaken out with ether ami the ethereal layer Was washed, first with 5% sodium carbonate solution and finally with water. The ethereal layer was dried over t sodium sulphate ami evaporated to a small bulk. Then the ester was distilled under reduced pressure. It formed a crystalline mass on cooling.

The boiling point of the ester was 25(1 C (lit. 256 C)"1"'. XItT 2—OH ; C—OCH3 V NH, X- HCi NH2

+ CH:, —OH Dry T Q | + H2O

2.3.1.2.1.1. Identification of methyl anthranilate:

2.3.1.2.1.1.1. Nitrogen contents (N%):

• The nitrogen content of. methyl antlt^anilate was determined using

Kjeldhal method, 0.401 Og! of methyl anthranilate was placed in a Kjeldhal

flask and digested with 10 cm'1 of concentrated sulphuric acid. One gram of sodium sulphate and 0.2g .of manganese dioxide was added. The mixture was diluted with distilled water and transferred quantitatively into a distillation apparatus, few antipumping granules were added, the excess acid was neutralized with 4M NaOH using litmus paper. The mixture was distilled and all ammonia gas evolved was received in 50 cmJ of 0.1 M HCI in the receiver flask, which was adjusted so that the tip of the condenser just dipped in the acid solution. The distillation was continued till all the ammonia gas absorbed by HCI. Few drops of screened methyl orange were added to the receiver and the excess acid was titrated with standard 0.1 M

NaOH"111. Results are shown in Table 1.

Calculation:

N:l()lli 1-4 Weight of sample / r i [(50X0.1)- (23..8X0.1)] 1.4 % ot tth e estet r = -U / 0 401& —~ = ! • ) Table 1: N% of methyl arithranilate:

Nitrogen content Found Calculated

N% 9..15 9.27

2.3.1.2.1.1.2. .I.R spectra:

The I.R1 spectra of methyl • anthranilate were recorded in IR spectrophotometer. The spectra were measured at room temperature and are shown in Table 2.

Table 2: The I.R spectra for methyl anthranilate:

Croup Vibrations in cm" (Found) N -l-l. 3440 O-H 3000* ~T675~ 300 c6 240 ••This shift is due to the hydrogen bonding.

2.3.1.2.2 The coupling reaction between methyl anthranilate and'the hydroxylamine:

The procedure followed was that of Pope" 1. To a cooled solution of

24.()g of sodium hydroxide (0.6 mol) in 150cm' of distilled water, was added slo\\l\ with stirring 20.8g (0.3 mol) of hydroxylamine

5.S hydrochloride. To this, solution was added 22.%g (0.15 mol)?* of methyl >t ' it'

anthrani late, and -enough j-nethyFaJcohol lo bring it., into, solution.'The

•••••'. -,";-••• -' %' ;.•...•*, ••;

solution was allowed to stand for three days at room temperature, then

distilled under reduced pressure to remove the excess solvent, leaving

about 50 cm' of the mother liquior in the flask where sodium salt of the

hydroxamic acid was precipitated. The salt was filtered by suction and washed with'ether, the fi Iterate was made acidic with hydrochloric acid where the free hydroxamic acid was precipitated. The crude product was recrvstallized from ether to give light brown crystals, yield 14.0g. (61.5%) m.p. 148C (lit. 140 (')'''. It compound gave a good test for hydroxamic acid.s with l-'e (HI) chloride. It is iairk stable up tol40°C.

MI. .rorih + CH3— Oil

2.3.2 Characterization of the anthranilic hydroxamic acid:

2.3.2.1 IR spectra: 5

The IR spectra of the compound were recorded in IR specirophoiometer. The spectra were measured at room temperature in KBr disk. I'he results are ^hown in Table 3. Table 3: The IR spectra of anthranilic hydroxamic acid: I , • • V 0

-1 Group •I • Vibration in cm (Found) N- i-1 3150 0- H 3040* Q = 0 1640 Q _ N 1320 l N- o 900 -This shift is due .to.the hvdroaen bondine.

2.3.2.2. Nitrogen contents (N%):

It was done as in 2.3.1.2.1.1 using 0.4120g anthranilic hydroxamic

acid. Results are shown in Table 4.

Calculation:

.VXM)iKi-(VXM) ,ii| X 1.4 N%=. Nat • • Weight of sample

1 00X 0.1 )-(50X0. l)jx 1.4 N% of the ester = - -- = 17.5 0.4120

Table 4; N% of anthranilic hydroxamic acid:

Nitrogen content Found Calculated

N°/n 17.5 8.4

2.4. Piej)aiati()i] of reagents:

1X10 "M of anihranilieih\droxamic acid \v|s prepared by dissolving

0.38Og of the acid in 250cm1 uilumetric ilask, using water.

60 3 1X10 M of the anthranilicacid was prepared by dissolving:0.0760g of the acid in 500 cm3 volumetric flask using arriyl alcohol.

100 ppm stock solution of each of the Fe(IlI), V(V), Cu(II), Ni(II), and Mo(VI) ions were prepared from the corresponding salt by dissolving the following weights in 100 cm3 volumetric flasks; 0.0290g of anhydrous ferric chloride, 0>0229g of ammonium metavanadate, 0.0392g of Cu(II) sulphate, 0.0495g of nickel nitrate hexahydrate and 0.0187g of ammonium molybdate, respectively.

The working solutions1 w.ere prepared by serial dilutions of the stock solutions with water.

6M, 10M stock solutions of hydrochloric acid and nitric acid were prepared by taking 50.13 Cm'of concentrated HC1 and 37.5 cm1 of.HNO.y, respectively, in 1.00 cm' volumetric flask and completed with distilled water:.:

1M, 2M. 3M • HC1. and HNO^ were prepared by dilution of appropriate volume of the 6M HC1 stock solution and 10M HNO3 stock solution. pH'•• 1, 2, 3. 4. 5.. 6 were prepared by diluting appropriate volumes oflM 11C1. UNO; orll^SO,.

6 1 t

2.4.1.Preparation of pH 1.0 to pH 10.0 buffer solutions:

Buffer pHl.O:

This was prepared by mixing 250 cm'of 0.2N KC1 with 485 cm3 g of

0.2 N.HCI in 1000 cm"1 volumetric flask and diluted to the mark with

distilled water.

Buffer pH2.0:

This was prepared by mixing 322.5 cm'of 0.2N HC1 then diluted to

mark with distilled water.

Buffer pH3.0:

Was prepared b\ adding 500cm' of 0.IM potassium biphthalate to

203.2 cm"1 of 0.1N HCTin ldm'. volumetric tlask and diluted to mark with distilled.

Buffer pH4.():

Was prepared as the same as pll 3.0 except the addition of 147cm1 of

0.IN KC1 instead of 203.20 cnr1 and diluted- to. the mark with.distilled water.

Buffer pH5.0:

Was prepared by mixing 226 cm' of 0.1 M sodium hydroxide with

:> 500cm' of 0.1 M potassium hydrogen phthalate (C8H5O4K) in 1000cm volumetric flask, then completed to the mark.with distilled water.

62 Buffer pH6.0:

This was prepared by adding'76.4cmJ of O.OSlVf'sodium acetate to

36cm1 of 0.05M acetic acid in dm3 volumetric flask, and completed to the

mark with distilled water.

Buffer pi 17.0:

It was prepared by adding 291cmJ of 0.1 M NaOH to 500 cm3 of

potassium hydrogen phthalaie in 1000cm' volumetric flask, and diluted

with distilled water to the mark.

Buffer pHS.0:

461.0 em; of 0.1M NaOll added to 500 cm' of 0.1 M potassium

h\d,rngen phthalate in, I dm' volumetric llask and completed to the mark

\\ iih distilled water. ' •

Buffer pH9.l):

. 2UiS cm" of O.IM NaOll was added to 500 cm'of 1:1 mixture of

0.1M KCI and 0.1 M boric acid (ihBOj in a dm1 volumetric llask, and com|~)leied to the mark with distilled water.

Buffer pH 10.0:

It was prepared b\ abiding 500 cm~ of 0.1M boric acid in 0.1M KCl to 4.W cm' ol' O.IM \a( )| 1 in 1000 \olumetric llask and completed to the. mark \\ ith distilled w alei 2.5. Analytical parameters for Fe(T,JI), V(V) and Ctt(II) ions:

2.5.1.Deternfination ofc wavelength of the njaxjnium ?4bsarbancer for '•'•', • .-T-: '••••'' -' -• $ .%:A*A •••• . •

Fe(III):

The spectra of Fe(IIl)-anthranilic hydroxamic acid complex using

1X1O"2M of the ligand and 100 ppm of the metal ion solution were measured. The maximum absorbance was observed at 482 nm.

2.5.2.The effect of the pH on the absorbance using HCI, HNO3 and,

H2SO4 acids:

To study;the effect of the.pH on the absofbance on the wavelength of maximum absorbance. three sets of mixtures were prepared by keeping the concentration of the ligand and the metal ions constant at 1x10"" M of the ligand and 15ppm of th.e Fe(III). ion, respectively, and taking variable volumes of the acids to gi\e variable pH values. The results are shown in

Table 5.

64 TableSi1,, The .effect of pH of an absorbahce of Fe(III)-ligand \yith different acids at 482 nm: .

1 Absorbances PH HNO3 HC1 H2SO4 1 0.185 0.090 0.080

-> • 0.224 0.160 0.140 3 0.251 0.201 , , 0.198 4 0.286 0.245 . 0.242 ,5 0.3 17' 0.302 0.294

'6 ' 0.288 • 0.265 0.260 7 0.165 0.143 0.142 1M 0.127 0.120 0.117 2M: 0.086' • 0.060- 0.005 3M 0.002 ' 0.002 0.002 J

2.5.3. Construction of calibration curve for Fe(III)-anthranilic hydroxaniic system:

The calibration curve for Fe(lll) was constructed by placing 0.50,

1.0. 1.50, 2.00. 2.50.3.00 and 3.75cm" of the 100 ppm Fe(III).stock solution in 25 cm' volumetric tlasks. 2.50 cm' of I X10"' M nitric acid were added, followed b\ 2.5.cm' of anthranilic hydroxamic acid in each flask tb give

0.20. 0.50. O.SD. 2.4. o. S. 10. 12. and 15 ppm of Fe(III), respectively, when completed to mark with water. The absorbances were recorded at 482 nm.

The re^uli^ are shown in'-the fable 6; and Fiu 1. Table 6: Calibration fofr Fe(III) at 482 nm:

Concentration (ppm) AbsorbancC* 15.00 0.288 12.00 0.224 10.00 0.196 8.00 0.148

6.00 • 0.124 4.00 0.076 2.00 ' 0.037 0.80 0.016 0.50, 0.015. 0.20 0.014

0(1 0, 0.30- 0.25 - Ig 0.20- 1,0.15 -\ < 0A0- ' 0.05 '- 0.00 0 4 6 8 10 12 14 16 Concentration (ppm

Fig 2. Calibration curve.for Fe(III)

67 2.5.4. The effect of foreign ions on the absorbance of F£(III) anthraniiic hydroxamic acid:

2.5.4.1 The effect of V( V) on the absorbance or Fe(III)-ligand:

To study the eiTeet of foreign metal ions on the absorbance of

Fe(lll) 0, 0.5, 1, 1.5. 2. 2.5 3.0 and 3.75 cm3 of 100 ppm of V(V) solutions were placed in 25, em"1• volumetric flask., 3.0 cm' of 100 ppm of Fe(III) solution were added to each flask, followed by 2.5 cm3 ofIX10"4 M nitric and 3.00 cm? of anihranilic hydroxamic acid. The volumes were completed to the mark; with distilled water. The absorbances were:recorded at 482 nm,

The results are shown in Table 7.

Tabie7: live effect of Y(V) on the absorbance of Fe(NI) at pH5, and

482 n m:'

i Concentration of V(V) j Concentration of Fe(II) Absorbance

(ppm) (ppm) ojxf 2 ppm 0.250 "0.235"

4.00 0.205 0.1O4

0.182 0.00 7 0.178" 2.00 0.162 5.00 1T(J6(F 2.5.4.2. The. effect,of Cu(II) on the absorbance of Fe(III)-anthranilic

hydroxamic acid at pH5 and 482 nm.

The same procedure as in 2.5.4.1 was used except V(V) solution was replaced by Cu(H) solution. The results are shown in Table 8.

Table 8: The effect of Cu(II) on the absorbance of Fe(III)-anthranilic hydroxamic acid at pH5 and 482 nm:

Concentration of Cu(II) Concentration of Fe(III) Absorbance (ppm) (ppm) . . o.oo 12.00 . 0.256 - 2.00 0.230 4.00 0.220 6.00 0.210 8.00 "* 0.209 10.00 0.205 12.00 0.204 15.00 0.220

2.6. Analytical parameters for V(V):

2.6.1. Determination of V(V) wavelength of the maximum absorbance:

To The* specrra of V(V)-anthranilic hydroxamic acid using lxlO"2M of the ligand disseised in amyl alcohol were measured in the visible absorption region. The maximum absorbance was observed at 472 nm. •

6') 2.6.2.The effect of the pH on the absorbanc°eof V(V)-Ligandcomplex

• t ; • . • • using HCl, HNO3 and H2SO4:

Three sets of mixtures were prepared by keeping concentration of the ligand and the metal ions constants at lxlO"3 M for the ligand and 15ppm for the V(V) ion, and take variable concentrations of HCl, HNO3, H2SOj acids to give different pH values. The mixtures were transferred to 100cmJ separator}' funnels. Shaken and allowed to stand for 5 min. The organic layers were separated. The absorbances were recorded at 472 nm. Results are shown in Table 9.

Table 9. The effect of pH on the absorbance of V(V)-ligand with effect acids at 472 nm:- Absorbances pi I UNOj MCI H2SO4 1 . 0.432 . 0.524 0:420 -> 0..X16 0.893 0.805 0.624 0.707. 0.612 4- 0.491 0.520; 0.408 5 0.307 0.341 0.295

6 0.200 : 0.204 0.200 1M 0.120 0.165 0:. 110 2M (V.OSi) 0.107 0.060

70 2.6.3.Construction or calibration curve for V(V)-anthranilic hydroxamic acid complex:

The calibration curve was constructed by placing 0.50, K00, 1.50,

2.00, 2.50cm1 of lOOppm ofV(V) in 25 cm volumetric (Tasks, followed by

2.50 cm3 of 0. 1IV1 11C1 and then cpmpleted to the mark with distilled water.

The solutions were transferred quantitatively to 100 cm' separately funnel and to each funnel, 5cm' of Ixl0"M anthranilic hydroxamic acid soluiion in amyl alcohol were added. The mixtures were shaken gently for two min then allow to separate. The organic layers were separated and the al>M)rbances were recorded at 472 nm. The results are shown in Table 10 and figs. 2a and 2b.

Table II); Calibration for \'(V) at 472 nm pH2

Concentration of V(V) (ppm) 1 Absorbaiice 10.00 0.953 r 8.00 0.770 6.00 0.581 : • -Loo 0.400 | 2.00 0.385 0.80 0.183' • 0.W) 0.137 0.40 0.082 ' 1 0.20 0.040

U.(JJ>{) 0.016 Bkink 0.010 VII. Metabolism of Ammpnia 239

B. Overall stoichlometry of urea cycle Protein Aspartate + NH3 + CO2 + 3 ATP

Urea + fumarate + 2ADP + AMP + 2 P, +PP,+ 3 H2O Four high-energy phosphates are consumed in the synthesis of Amino acids

each molecule of urea: two ATP are needed to restore two ADP to (i-Ketoglutarale -J Aininolmns/graso two ATP, plus two to restore AMP to ATP. Therefore, the synthesis of urea is irreversible with a large, negative AG. One nitrogen of •&~~~~~~^~ ot-Ketoacids the urea molecule is supplied by free NH3, and the other nitrogen Glutamate by aspartate. Glutamate is the immediate precursor of both NAO'^ Oxafoacetate ammonia (through oxidative deamination by glutamate dehydroge- Glulamate \ Aspartate nase) and aspartale nitrogen (through transamlnation of oxaloac- dohydiogonaso 1 \ aminotransloraso elate by aspartate aminotrans(erase). In effect; both nitrogen o-Keloglutarate u-KetorjUilarale atoms of urea arise from glutamale, which in turn gathers nitrogen " + NADH + H1 / from other amino acids (Figure 21.13).

C. Regulation of the urea cycle N-Acetylglulamale is an essential activator for carbamoyl phos- phate synthetase I, the rate-limiting step In the urea cycle (see Figure 21.1.1). N-Acelylgiulamate is synthesized from acetyl CoA Carbamoyl cilrulline and glutamate (Figure 21.14).: Th.eintrahepat.ic concentration of phosphate this compound increases after ingestion of a protein-rich meal, leading to an increased rale of urea synthesis. Omilhine UREA Arginlnosucdnate ' CYCLE A

. . . Fumarate \/ll. METABOLISM OF AMMONIA : , Argmine

Although ammonia is involved in the formation of urea in the liver, the evel of ammonia in the blood must be kept low because even slightly Figure 21.13 alevaled concentrations (hyperammonemia) are toxic to the certtral Flow of nitrogen from amino acids to lervous system. There must, therefore, be a metabolic mechanism by urea. Amino groups for urea synthesis vhich nitrogen is moved from peripheral tissues to the liver for ulti- are collected In the form of ammonia nale disposal as urea, at the same time, maintaining'low levels of cir- and aspartate. :ulaling ammonia. i. Sources of ammonia Ammonia is produced from tITe metabolism of a variety of com- pounds. Amino acids are quantitatively the most important source of ammonia, because most western diets are high in protein and provide excess aniino acids, which are deaminated to produce Glutamate ammonia. Acetate Acetyl CoA 1. From amino acids: Many tissues, but particularly the Ijver, form ammonia from amino acids by the aminotransferase and\ gluta- Hydrolase male dehydrogenase reactions previously described (see pp. 234-235). N-Acetylglutamate 2. From glutamlne: The kidneys form ammonia from glutamine by the action of renal glutaminase (Figure 21.15). Most ol this Figure 21.14 ammonia is excreted into the urine as NH.<\ which is an impor- Formation and degradation of N- tant mechanism for maintaining the body's acid-base balance. acetylglutamate, an allosteric Ammonia is also obtained from the hydrolysis of glutamine by iactivator of carbamoyl phosphate synthetase I. 20. Cholesterol and Steroid Metabolism

#0.7 Which ofiao|jfi^fpl(bwing statements about familial 20.10 A patient has a genetic defect resulting in a defi ||jp|ef)i3,is INCORRECT?, . ' ciency of lipoprotein lipase. After eating a meal con r { taming a large amount

...... , .v.... f. ( ilia cjefect is a lack of binding of Correct answer - A, Chylomlcrons are produced "j, ;jll'pj.:i9;cel|,§iiiface receptors due to a lack of LDL by the,Intestinal mucosal cells from dietary llpld. \^ apolipoprotelrt B-100. They corttaln primarily trlacylglycerol, which Is E', flqmozygotes for .this disease are able lo metabo- normally degraded by lipoprotein lipase. If this •; Itee chyiornfqrpns normally, • i enzyme is deficient, chylomicrons would accumu- late In the plasma after a llpld-rlch meal.

20.11 Which one of the following steroids is synthosi7f>tl 1 from cholesterol WITHOUT boing hydroxylnlod by I /' u-hydroxylase? . A Corticosteiono 0.8 Which one of the following changes would you expect B. Cortisol 1 ,|n a patient with decreased activity of lipoprotein ! C. Testosterone lipase? \ D Estradiol A. Elevation of plasma chylomicrons on|y. E Androsleneujone B. Elevation of both plasma chylomicrons arid very low density lipoproteins. ( Correct answer = AT C. Eievation of plasma HDL only. D. Elevation of plasma LDL only. 20 12, Which one of the following steroids is synthesized E. Elevation of both plasma HDL and LDL. ifrom cholesterol WITHOUT being hydioxylnlnd by 21-hydroxylase? A Corticoslerone B Cortisol C Aldoslerone _ D 11-deoxycorticosterone •E, 17 u-hydroxyprogoslerone

f Correct answer« E, 1 9 The lipoprotein particles that have the highest percent- age concentration of cholesterol are: A Chylomicrons. B VLDLs. C LDLs. D. HDLs E. Serum albumin-associated lipid.

[, Correct answer = C. \)^{ ',:J:J,\B1^1

SITS H i Mm pms in ml mm s i i 11 mtff' 20. Cholesterol and Steroid Metabolism

#0.7 Which b'nep wing-statements about familial 20.10 A patient has a genetic defect resulting in a deli hypercholesjl PfrjNCQRRECT? ciency of lipoprotein lipase After eating a meal con t a- * i . i _ i A. This iiraf Isprder's characterized by elevaf taining a large amount of fat, one would expect to tions xHftb^jfjiMfa1 arid especially IDL eholes- see a plasma elevation of , A Chylomicrons B. We .disease is^rnanifested clinically by xanthomas B. VLDLs and accelerated atherosclerosis. C. LDLs C. Hete'rozygotesjor this disease may have plasma D HDLs. qjjpjestepl lev^fs that are Approximately tvyice nor- E. Serum albumin-associated lipid •''jqaU''-"'!"*?••".',•,• • -y, i;i D. The primary metabolic defect Is a lack of binding of Correct answer « A. Chylomicrons are produced LDL to cell surface receptors due to a lack of LDL by the,Intestinal mucosal cells from dietary llpfd. apolipoprotein B^IOO. They doritalh primarily trlacylglycerol, which is E. Homozy^otes for this disease are able to metabo- normally degraded by lipoprotein llpase. If this lize chylomicrons normally. enzyme Is deficient, chylomlcrons would accumu- late In the plasma after a llpld-rlch meal.

20.11 Which one of the following steroids is synthesized Irom Cholesterol Wl 11IOUT bainrj hyrlioxylnlnd by I i (x-hydroxylase? A Coiticosterono 20.8 Which one of the lollowing changes would you expect B Cortisol in a patient with decreased activity of lipoprotein C. Testosterone lipase? D ,Estradiol A. Elevation of plasma chylomicrons only. ' E Androslenedione B. Elevation of both plasma chylomicrons and very low density lipoproteins. Correct answer = A. C. Elevation of plasma HDL only. D. Elevation of plasma LDL only. 20 12 Which one of the following steroids is synthesized E. Elevation of both plasma HDL and LDL. ,from cholesterol WITHOUT being liydioxylatod by 21-hydroxylase? A Corticosterone B Cortisol C Aldostorono D 11-deoxycorticosterone E 17 (K-hydroxyprogesterono

( Correct answer» E. 20.9 The lipoprotem particles that have the highest percent- age concentration of cholesterol are: A Chylomicrons. B VLDLs. C. LDLs. D HDLs E. Serum albumin-associated lipid.

:.HI< , •'.• -" •; '••• ).''\ '" -

t 1.00 -i

0.80 -

o c 0.60 - a o en X) 0.40 -

0.20 -

0.00 - 0.0 2.0 4.0 6.0 8.0 10.0 Concentration (ppm)

Fig 3a. Calibration curve for V(V) 0.0 0.5 1.0 1.5 2.0 Concentration (pp

Fig 3b. Calibration curve for V(V)

73 2.6.4.The effect of foreign ions (Mo(VI), Fe(III) Cu(II), Ni(H)} on the absorbance of V(V)-anthranilic hydroxamic acid system at pH 2 and 472 nni: 2.6.4.1 Effect of Fe(III) on the absorbance of V (V)-ligand complex: To study the effect of Fe(III) ion on the absorbance of V(V) 0, 0.50,

1.00, 1.50, 2.00, 2.50, 3.00, and 3.75cm3 of 100 ppm of Fe(III) were placed

in 25cm1 volumetric flasks, 3.0 cm of lOOppm V(V) solution were added

-i

followed by addition of 2.5cm' of 0.1M HC1 then the flasks were

completed to the mark with distilled water. The solutions were transferred

quantitatively to lOOcnv' separatory funnels and to each funnel 5cm of

anthranilic hydroxaniic acid solution in amyl alcohol were added. The

mixtures were shook gently for two minutes. Then the organic layers were

separated and the absorbances were recorded at 472 nm. The results are

shown in Table 1 1.

Table! I; the effect of Fe(IIl) on the absorbance of V(V) at pH2 and 472 nm: Concentration of Fe(III) (ppm) Concentration ofV(V)(ppm) Absorbance 0.00 12 ppm 0.702

2.00 Vi. 0.650 -

4.00 Ib 0.583

6.00 k* 0.572

8.00 H. 0.570

10.00 U 0.564 12.00 0.560 15.00 0.555

74 2.6.4.2.The effect of Mo(VI) on the absorbance of V(V)-ligand complex at pH2 and 472 nm: The same procedure as in 2.6.4.1 was used except Fe(III) solution

was replaced by Mo(VI) solution. The results are shown in Table 12.

Table 12: The effect of Mo(VI) on the absorbance of V(V)-ligand:

Concentration of Mo(VI) Concentration of V(V) Absorbance

(ppm) (ppm)

0.00 4ppm 0.470

2.00 0.365

4.00 0.282

6.00 vv 0.238

8.00 vi 0.214

10.00 ct 0.180

12.00 vt 0.172

15.00 0.149

2.6.4.3.The effect of Ni(II) on the absorbance of V(V)-ligand complex at pH2 and 472 nm: The same procedure as in 2.6.4.1 was used except Fe(III) solution was replaced by Ni(II) solution. The results are shown in Table 13. TableO; the effect of Ni(II) on the absorbance of V(V)—anthranilic

hydroxamic acid at pH2 and 472 nm:

Concentration ofNi(II) Concentration of V(V) Absorbance (ppm) (ppm) 0.00 6ppm 0.601 2.00 a. 0.600 4.00 0.598 6.00 0.600 8.00 0.599 10.00 0.596 12.00 vt 0.598 15.00 0.597

2.6.4.4.The effect of Cu(II) on the absorbance of V(V)-Iigand complex

at pH2 and 472 nm:

The same procedure as in 2.6.4.1 was used except Fe(III) solution was replaced by Cu(II) solution the results are shown in Table 14.

76 Table 14: the effect of Cu(II) on the absorbance of V(V)-ligand

complex at pH2 and 472 nm:

Concentration of Cu(II) Concentration of V(V) Absorbance ppm (ppm) 0.00 10 ppm 0.865 2.00 it 0.873 4.00 0.874 6.00 0.860 8.00 0.859 10.00 0.886 12.00 0.896 15.00 0.830

2.7 Analytical parameters for Cu(ll) ion:

2.7.1.Determination of Cu(II) wavelength of the maximum absorbance:

The spectra of Cu(II)-anthranilic acid complex using lxl0""M of the

ligand and lOOppm of the metal ion solution were measured in the visible absorption region the maximum absorbance was observed at 390 nm.

2.7.2.The effect of the buffer pH on the absorbance of Cu(II)-Iigand complex:

In series of eight 25cm' volumetric flasks 3.0cm"'1 of 100 ppm of

Cu(II) solution were added followed by addition of 7.0cm' of buffer solutions pH2-pH9 and then 10cm'1 of lxl0"~M of anthranilic hydroxamic were added and completed to the mark with distilled water. The mixtures

77 were shaken and the absorbances were recorded at 390 nm. Also the effect

of acidic and basic pH as in 2.5.2 were studied for Cu(II) ligand but a

precipitation was produced in all cases. The results are shown in Table 15.

Table 15: The effect of pH on the absorbance of Cu(II)-ligand complex: Absorbances

pH Buffer HNO3 H2SO4 HCI NH4OH

i 0.085 ppt ppt ppt ppt

3 0.100 it u tC

4 0.120 it CC

5 0.172 4. V

6 0.108

7 0.1 70 U Vi.

8 0.160 CC

9 0.143

Where ppt refers to precipitate

2.8.Construction of calibration curve for Cu(II)—anthranilic

hydroxamic acid complex: *

It was constructed by placing 0.5, 1.00, 2.00, 2.50, 3.00, and 3.75 cm' of lOOppm Cu(II) solutions in 25cm1 volumetric flasks, followed by addition of 5cm1 of buffer solution of pH6 then 5cm1 of 1X1O""M anthranilic hydroxamic acid were added. The flasks were completed to the

78 mark with distilled water and the mixtures were shaken. The absorbances were recorded at 390 nm. The results are shown in Table 16.

Table 16: Calibration of Cu(II)-ligand complex at pH6 and 390 nm:

Concentration of Cu(II) (ppm) Absorbance

15.00 0.186

12.00 0.134

10.00 0.120

8.00 0.088

6.00 0.068

4.00 0.044

2.00 0.010

blank 0.01 4 6 8 10 12 14 16 Concentration (]

Fig 4. Calibration curve for Cu(II) 2.9.The effect of foreign ions Fe(III), V(V) on the absorbances of

Cu (II)- anthranilic hydroxamic acid at buffer pH6 and 390 nm:

2.9.1.The effect of Fe(HI) on the absorbance of Cu(II)-ligand complex.

To study the effect of Fe(III),on the absorbances of Cu(II) ion 0,

0.50, 1.00, 1.50, 2.00, 2.50, 3.00, and 3.75cm' of IOOppm of Fe(III)

-i solution in eight 25cm' volumetric flasks followed by addition of 100 ppm of Cu((I) solution and 5cmJ of buffer solution pH6 were added and 3.0 of lxl0"2M anthranilic hydroxamic acid were added. The flasks were completed to the mark with distilled water and the mixtures were shaken.

The absorbances were recorded at 390 nm, the results are shown in

Table I 7.

Tablel7:The effect of Fe(III) on the absorbance of Cu(II) anthranilic hydroxamic acid at pH6 and 390 nm: Concentration of Fe(III) Concentration of Cu(II) Absorbance (ppm) (ppm) 0.00 12 ppm 0.132 2.00 t.i 0.207 4.00 >,t 0.276

6.00 v* 0.349

8.00 *v 0.436 10.00 0.501 12.00 0.604 15.00 0.641

81 2.9.2 The effect of V(V) on the absorbance of Cu(II)-Iigand complex:-

The same procedure as in 2.8.1 was used except the Fe(III) solution was replaced by V(V) solution the results are shown in Table 18

Tablel8: The effect of V(V) on the absorbance of Cu-ligand complex at pH6 and 390 nm:

Concentration of V(V) Concentration of Cu(II) Absorbance (ppm) (ppm) 0.00 12 ppm 0.132 2.00 u 0.134 4.00 cc 0.187 6.00 u 0.251

8.00 k t> 0.475 10.00 U 0.544 12.00 bt 0.674 15.00 0.807

82 CHAPTER THREE

3. Discussion

Anthranilic hydroxamic acid was prepared by coupling of

methylanthranilate (prepared by esterification of anthranilic acid with

methyl alcohol using the Fisher-Speir method) with freshly prepared

hydroxylamine.

The Fisher-Speir method, in which dry hydrogen chloride was passed in a mixture of anthranilic acid and methyl alcohol till the mixture was saturated, is more favourable than the ordinary method of esterification in which concentrated sulphuric acid was used as a dehydrating reagent.

With anthranilic acid, which is an amino acid, the ami no group would undergo salt formation, and thus the hydrogen sulphate of anthranilic acid methyl ester would result H:SO4.NH2C6H4COOCH3, which would be very difficult to isolate from the excess of sulphuric acid. Also, considerable sulphonation of the benzene ring may result. These difficulties were overcome by the Fisher-Speir method. The function of hydrogen chloride is primarily catalytic, since Fisher found that 5% of hydrogen chloride in the reaction mixture would cause efficient esterification.

The advantage of this method is that the hydrogen chloride does not usually add onto unsaturated grouping. It does not affect the aromatic groups, and the excess hydrogen chloride of the ester can be readily removed by direct evaporation. A further advantage is that the

corresponding ether is not formed as by-product.

The ligand was characterized by the usual reaction of hydroxamic

acid with acidic V(V) and Fe(III) solutions that gave a blood-red colour in

amyl alcohol and deep—violet colour in aqueous solution, respectively.

The absorbance of Fe(III)-hydroxamic acids complexes increases with increase of pH. In this study, the effect of pH on the absorbance of

Fe(III)-anthranilic hydroxamic acid was in accordance with trend. The maximum absorbance was obtained at pH 5.0 at maximum wavelength of

482 inn, These results resembles those obtained by Brown and Glennon(|y>.

Above pH 7.0, both A.m:iN, and £max remain constant, consistent with the formation of the MHi species which predominates in this pH range, while at pH > 9.0 there is a marked decrease in absorption'l9).

For Cu(II)-anthranilic hydroxamic acid complex, the use of acidic or basic pH lead to the precipitation of Cu(II)-ligand complex (Table 15). But when using buffer pH (acetic acid/sodium acetate) a clear green colour of

Cu(II)-Ligand- complex was obtained. The absorbance increases subsequently with the increase of pH value and the maximum absorbance was obtained at pH 6.0 at maximum wavelength of 390 nm. These results are comparable with the work of Kurzak and Kurzak(Dj). According to Narian and Shulka, the position of the absorption

maximum for Cu(II) complexes in water solution depends on the number of

Cu-N bondsD>). The maximum at 620-630 nm is typical for Cu(II)

coordinated by two nitrogen atoms of the amino or amido groups of the

cc-amino acid or peptides, respectively. Since the dinuclear complex

r [Cu2H"'L2] exhibit A.m;lx at 654 nm, it is possible to conclude that one

copper cation is coordinated by two nitrogen atoms of ligand.

V(V)-anthranilic hydroxamic acid complex was extracted in acidic

medium in amyl alcohol at pH 2.0 because in aqueous solution V(V)-

anthranilic hydroxamic acid complex has not clear colour. It was observed

that the maximum extraction in acidic medium decreases sharply with the

increasing of pH value. The maximum wavelength for maximum

absorbance was recorded at 472 nm. This is comparable with that of

Priyadarshini and Tandon"'01.

All coloured products have broad absorption bands, it appears that

only two coloured complexes of V(V) with N—benzoyl—N—phenyl

hydroxamic acid have formed violet colour in strongly acids solution. A

methyst coloured systems are seemingly produce by the interaction of

violet and mahogany red complexes.

As shown in Table 7, V(V) interfered with determination of Fe(III) above concentration of 2 ppm. This may be attributed to the competition of both cation towards the ligand, whereas Cu(Il) interfers slightly with the determination of Fe(III) ions even at a high concentration of the Cu(II) ions, this may be due to the that Cu(II) ions need buffer pH to form its complexes and not the acidic pH (Table 8).

Both Cu(II) and Ni(II) do not interfer with the determination of V(V) ions even at high concentrations, Fe(III) ion produced slight interference, while Mo(VI) ions have a pronounced interference.

Both V(V) and Fe(III) ions interfered markedly with the determination of Cu(II) ions, and made impractical under these conditions.

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8')