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2

EXPERIMENT At

2.0 CHEMICALS

In the present study, the following chemicals were used, whose details are given

below as follows :

2.01 Acids

The acids used in the present investigation were hydrochloric, sulphuric,

perchloric acid and deuterium chloride.

[A] Hydrochloric acid (HCI) : (Qualigens ExcelaR)

Sp. gr. 1.18 gm ml-1, 11.6 mol dm-';

[B] Sulphuric acid (H,SO 4) : (Qualigens ExcelaR)

Sp. gr. 1.83 gm ml-1, 18.0 mol dm_,;

[C] Perchloric acid (HCI04) : (Qualigens ExcelaR)

Sp. gr. 1.70 gm ml-1, 10.0 mol dm-';

[OJ Deuterium Chloride (OCI) : Procured from Bhabha Atomic Research Centre

(BARC), Mumbai (India). Isotopic purity > 95%

DCI was diluted with 0,0 to produce the solutions of required molarity.

All the acids were standardised with sodium hydroxide, which in turn was

standardised with potassium hydrogen phthalate (BDH, AnalaR)

2.02 Bases

[A] Sodium Hydroxide (NaOH) : (Qualigens ExcelaR) mol. wt. 40

(37)

Ex onmontal

[B] Potassium Hydroxide (KOH) : (Oualigens ExcelaR) mol. wt. 56.11

[C] Lith1um Hydroxide (LiOH H,O) : (S1gma) F.W. 41.96

[D) Sodium deuteroxide (NaOD) : Procured from BARC (Isotopic purity > 95%)

[E) Ammonia solution : (s.d. fine AR) sp. gr. 0.91, mol. wt. 17.03

Sodium deuteroxide was diluted with D,O to produce the solutions of required

molarity.

2.03 Solvents

[A] Water : Double distilled water was used for preparation of solutions for kinetic

studies. It was prepared by redistilling deionised water in distillation glass

assembly (BHANU) over alkaline potassium permanganate. Triple distilled water

was used for the preparation of metal complexes of hydroxamic acids and in the

study of acid-base equilibria.

[B] Deuterium Oxide (020) : Procured from BARC Mumbai (India)

Isotopic purity > 99.85%

[C] Other organic solvents : Other organic solvents used for the kinetic studies were

: 1 ,4-Dioxane (Oualigens ExcelaR), DMSO (Merck GR), DMF (BDH AR), Acetone

(Qualigens ExcelaR), Acetonitrile (Merck), Tetrahydrofuran (Merck), Methanol

(Merck HPLC). Ethanol (Qualigens ExcelaR) 2-propanol (Qualigens ExcelaR).

The above solvents were used as obtained without further purification.

2.04 Solutions

[A] Buffer solutions : pH meter was calibrated by using saturated solution of

potassium hydrogen tartarate (pH = 3.56 at 25°C), 0.05M potassium hydrogen

phthalate (pH = 4.02 at 25°C) and 0.01 M Borax (pH = 9.18 at 25°C).

(38)

Ex omnontal

[B] Ferric Chloride Solutions :

(1) Acidic :The ferric chloride solution used in the colorimetnc procedure for acidic

hydrolysis was prepared by dissolution of 44 gms of anhydrous salt (mol wt. 162.21,

Oualigenes) m 1 litre of distilled water containing 10 mi. of concentrated hydrochloric

acid to prevent ferric hydroxide formation. Solution so prepared was filtered through

ordinary filter paper.

(ii) Alkaline : Ferric chloride anhydrous (mol. wt. 162.21 Qualigenes), 44g

dissolved in 1 litre distilled water containing 100 ml concentrated hydrochloric acid, was

used for developing purple complex with hydroxamic acids for their analysis in kinetic

runs in case of alkaline hydrolysis.

2.05 Salts

Some anhydrous salts such as lithium chloride (Sigma), sodium chloride (Merck),

Potassium chloride (Qualigenes AR), Sodium bromide (Merck), Sodium perchlorate

monohydrate (Merck), Potassium sulphate (Merck) and Potassium bisulphate (Merck)

were used in the· measurement of salt effects. Each of these salts were dehydrated

before use.

2.06 Esters

Following esters were used in the preparation of dihydroxamic acids.

[A] Diethyloxalate : (Aldrich)

F.W. 146.14, b.p. 185°C, d 1.076

[B] Diethylmalonate : (Aldrich)

F.W. 160.17, b.p. 199°C, d 1.055

[C] Diethylsuccinate : (Aldrich}

F.W. 174.20, b.p. 217.7"C, d 1.047

(39)

2.07 Acid Chloride

[A] Hydroxylamme hydrochloride : (Aidnch A.C.S. reagent)

F.W. 69.49, m.p. 159°C

2.1 APPARATUS

Following apparatus were used in the experimental investigation.

2.11 Thermostat

Ex onmontal

Acidic hydrolysis as well as alkaline hydrolysis reactions were carried out in

thermostats at 35-85°C with automatic temperature control ±0.1 oc at required

temperature by a precision thermometer (0-50°C and 0-1 oooc, least-count 0.1 oc). The

temperature setting remained undisturbed throughout the course of investigation. For

low temperature CRYOSTAT (Remi CB-700) was used.

2.12 Mechanical Stirrer

(Remi motors 220 I 230 V, 50 - AC), magnetic stirrer (Remi equipments, 2 MLH)

and vacuum pump (Wiswo, RPM 5400) used for synthesis purpose. Melting point

apparatus (Tempo S.No. A 233) was used to determine melting points.

2.13 Spectrophotometer

[A] Kinetic measurements were made by the use of spectrophotometric method

using a Systronics UV-VIS 108 spectrophotometer.

[B] Unicam UV-2 300 spectrophotometer was used for UV spectra measurement to

study the Acid-Base equilibria (protonation and deprotonation behaviour) and for

metal complexation.

(40)

Ex onmental

2.14 pH-Meter :

Digital pH-meter (Systronics model 335) was used for pH measurements, pK,

determination and to observe the effect of pH on metal complexation.

2.15 Computer :

Least squares analysis were carried out on a Wipro 386 Acer Entra Pentium

Computer under MS Dos. The GW Basic language was used for preparation of

program.

2.16 Glasswares :

[A] Reaction vessels :Tubes with (B-14, B-19 and B-24) stoppers were used as flat

bottom reaction vessels.

[B] Test tubes : Elongated test tubes with wide mouth were used.

[C] Burettes : Corning burettes of 5 x 0.02 ml; 5 x 0.05 ml, 10 x 0.05 ml, 25 x 0.1

ml and 50 x 0.1 ml were used.

[D] Pipetts : Special pipettes of 0.5, 1.0, 2.0 ml were used for withdrawing aliquots

from reaction mixture.

[E] Piplip : Piplip is a device for pipetting liquids and used for transferring corrosive

and toxic liquids.

2.2 PREPRATION AND CHARACTERIZATION OF DIHYDROXAMIC ACIDS.

Three dihydroxamic acids [HOHNCO-(CH2) 0-CONHOH; n=O, oxalo

(i); n=1, malono (ii) and n=2 for succino dihydroxamic acids) have been

synthesized and characterized by m.p., elemental analysis, UV and IR spectral

data.

(41)

Ex eflmental

2.21 Prepratlon of Dlhydroxamic Acids :

All the three d1hydroxam1c acids, oxalodihydroxamic acid (ODHA).

malonod1hydroxamic acid (MDHA) and succinodihydroxamic acid (SDHA) were

prepared [1-3] by the dropwise addition of diethyloxalate (0.1 mole). diethylmalonate

(0.1 mole) and diethyl succinate (0.1 mole) to an ammonical solution containing of

hydroxylamine hydrochloride (0.2 mole) with vigorous stirring at ooc.

cooc,H, I

(rH,)" + NH,OH.HCI

COOC,H,

Diethylester Hydroxylamine

hydrochloride

NH3 ~

CONHOH I

(fH,), +

CONHOH

Dihydroxamic

Acid

2C,H,OH

[n=O, Oxalodihydroxamic Acid; n=l, Malonodihydroxamic Acid and n=2 for

Succinodihydroxamic Acid]

The white precipitate obtained filtered and recrystallized twice from distilled

water.

Trihydroxamic acid Desferal (DFB) was obtained from Novaratis. Monohydroxamic

acid, i.e., acelohydroxamic acid (AHA) was procured from Fluka whereas

benzohydroxamic acid (BHA) was prepared by standard method [4].

2.22 Characterization of Dihydroxamic Acids

Prepared dihydroxamic acids were characterized by m.p. (ODHA = 160"C, MDHA

= 155°C and SDHA = 180"C), elemental analysis, UV and IR spectral data.

Table 2.01 lists the formula, molecular weight and results of the elemental

(42)

Ex onmental

Table 2.01 Characterization data of prepared dlhydroxamic acids

Dihydroxamic Molecular MP Found (Calcd) %

Acid Formula 'C c H N

CONHOH C,H,N,O, 160 20.78 3.89 24.06

tONHOH (20) (3.33) (23.33}

(ODHA)

CONHOH I CH, C3H,N,04 155 25.98 4.82 20.56 I CONHOH (26.86) (4.48) (20.89)

(MDHA)

CONHOH

(tH,), C,H,N,04 180 31.94 4.87 19.42

tONHOH (32.43) (5.41) (18.92}

(SDHA)

Calculated values are given in parentheses.

(43)

rxpomnontal

analysts of the prepared dihydroxamic acids. Oxalodihydroxamic ac1d is less soluble 111

cold water but readily soluble in hot water whereas Succinodihydroxamic is highly

soluble in cold water. Thus, the solubility of dihydroxamic acid in cold water tncreases

from oxalo to succino dihydroxamic acid.

2.23 IR spectral study I Analysis

Infrared spectra of prepared dihydroxamic acids are shown in Figs. 2.01 to 2.03

and their representative data are shown in Table 2.02.

The IR spectra of hydroxamic acids and their complexes are generally very

complex though some characteristic bonds for different bonds I groups were

suggested. For example, the band around frequency 3200 em-• was assigned as the

NH valence frequency, while those observed in the 3080-3060 em-• region were

attributed to the OH vibrations and the CO valence vibrations [5-7]. The broad band

around 1610-1585 em-• observed is assigned to the ketonic carboxyl vibrations.

The most characteristic bands associated with the hydroxamic acid functional

grouping are due to C=O and 0-H stretching vibrations and these can be assigned

rather unambiguously. The N-0 and C-N stretching and 0-H deformation vibrations

were assigned with less certainty because of the nonavailability of systematic data on

the assignment of these bands in the infrared spectra of hydroxamic acids.

0-H Stretching Vibrations

In the dihydroxamic acids examined here the band due to 0-H stretching

vibrations has been assigned in the region between 3104 em-•. It is well known that the

absorption bands due to free 0-H stretching vibrations appear around

3650 em-•. The marked lowering of these bands implies the presence of strong

hydrogen bonding in these hydroxamic acids [8-12].

Vibrations in the 0--H stretching vibrations in different hydroxamic acids are

mainly due to the ability of acidic hydrogen of the hydroxyl group to form "hydrogen

bond" with electron rich atom.

(44)

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Ex onmontal

Table 2.02 IR Absorption Maxima In em-• of Dlhydroxamlc Acids.

OOHA Mode of MDHA Mode of SDHA Mode of

vibration vibration vibration

3262 N-H Stretch1ng 3199 N-H Stretching 3220 N-H Stretching

3104 0-H Stretching 3030 0-H Stretching 3059 0-H Stretching

2858 C-H Stretching 2894 C--H Stretching 2873 C-H Stretching

2374 C-H Stretching 2849 C--H Stretching 1641 C=O Stretching

2345 C-H Stretching 2361 C-H Stretching 1552 C-C Stretching

2128 C-H Stretching 2340 C-H Stretching 1468 C-N Stretching

1874 C-H Stretching 2166 C-H Stretching 1432 C-H Stretching

1655 C=O Stretching 1638 C=O Stretching 1402 C-H Bending

1562 C-C Stretching 1537 C-C Stretching of CH, group

1526 C-C Stretching 1459 C--N Stretching 1342 C-H Bending

1440 C-N Stretching 1426 C--H Stretching of CH2 group

1403 C-H Bending of CH, group 1239 C-H Stre. Unsymmetrical

1350 C-H Bending 1403 C--H Bending of CH2 group

1235 C-H Str. Unsymmetrical 1329 C--H Bending 1180 C-N Str. unconjugated

1096 C-0 Stretching 1257 C-H Str. Unsymmetrical 1085 N-H Bending (Wagging)

1067 N-H Bending (Wagging) ofCH2

group 1060 N-H Bending (Wagging)

1038 N-H bending (Out of plane) 1171 C-N Str. unconjugated 981 N-H Bending

949 N-H bending (Out of plane) 1121 C-0 Stretching (Out of plane)

848 C-H Rocking 1073 N-H Bending (Wagging) 720 C-H Rocking

812 C-H Rocking 990 N-H bending (Out of plane) of CH2 group

731 C-HTwisting 904 N-H bending (Out of plane)665 C-H Rocking

677 C-H Twisting 840 C-H Rocking of CH2

group of CH, group

616 C-HTwisting 696 C-H Rocking of CH, group

604 C-H Twisting 602 C-H Twisting

517 C-HTwisting 552 C--H Twisting

453 C-HTwisting

(48)

tx onmenlal

Thus the formation of strong hydrogen bonding causes a large shift of the 0-H

absorption band to lower frequencies, and may be ascribed to the resonance

stabilisation. A consequence of this resonance stabilisation should be to lower the force

constant of the carbon-oxygen band and is increase the contribution of single bond

form, thereby causing a fall in the frequency due to C=O stretching vibrations.

C=O Stretching Vibrations

In the N-arylhydroxamic acids examined here the C=O stretching bands are

assigned in the region between 1655 em-• and 1638. This assignment is made with

reference to the spectra of analogous compounds.

2.3 PREPRATION AND CHARACTERIZATION OF METAL COMPLEXES OF

DIHYDROXAMIC ACIDS

Metal complexes of dihydroxamic acids in solution as well as solid complexes

have been prepared and solid metal complexes were characterized by elemental

analysis and I R spectral data.

2.31 Preparation of Metal Complexes of Dihydroxamic Acids

Metal complexes of dihydroxamic acids in solution state as well as in solid state

have been studied.

To study the effect of pH on metal complexes of different hydroxamic acids, the

stock solutions of iron (Ill) ions were prepared from FeCI3.6H,O (Qualigenes) and a

known amount of 0.1 M HCI were added to prevent hydrolysis of metal ions. The final

concentrations of iron (Ill) and ligand (dihydroxamic acids) were about 5 x 10 .... M and.

2.5 x 1Q-3 M respectively. All the solutions were prepared in triple distilled water. pH

measurements were made by using Systronics digital pH meter 335. The pH of metal

complex solutions were adjusted with the help of NaOH and HCI solutions.

Spectrophotometric measurements of effect of pH on iron (Ill) complexes of different

hydroxamic acids were made with the help of Unicam UV 2 300 spectrophotometer.

(49)

Expomnontal

Sol1d metal complexes of 1ron (Ill), copper (II), nickel (II) and cobalt (II) of

dihydroxamic acids have also been prepared by the addition of 0.01 mole of

dihydroxamic acid (dissolved in 25 ml of H20) to the 0.005 mole of ferric chloride

(dissolved in 25 ml of 0.1 N HCI). A dark purple coloured complexes are obtained

whose pH were 1 to 1.6. The pH of mixed solutions were increased up to 5-7 by the

addition of 1% NaHC03

(Merck mol. wt. 84.01) solution. Mixture were stirred 2-3 hours

by magnetic stirrer and filtered, dark purple coloured complexes were obtained.

Similarly the Cu(ll), Ni(ll), Co(ll) complexes of dihydroxamic acids have been prepared.

The colour of the chelates are as follows :

Iron (Ill) - Dark purple

Copper (II) - Green

Nickel (II) - Sky blue

Cobalt (II) - Pink

All solid complexes were obtained are insoluble in water as well as in many other

organic solvents (like, DMF, Acetonitrile, CCI4 , CH,OH, acetone etc.) and they possess

very high thermal stability. These type of nature of solid metal complexes of

dihydroxamics acids can be explained by their polymeric structure [13].

2.32 Characterization of Metal Complexes of Dihydroxamic Acids

Prepared metal complexes of dihydroxamic acids were characterized by

elemental analysis, UV and IR spectral data.

Elemental analysis of some metal complexes of dihydroxamic acids are shown in

Table 2.03.

Infrared spectra of metal complexes of some dihydroxamic acids are shown in

Figs. 2.04 to 2.06 and their representative data are shown in Table 2.04.

It has been observed that the bands due to N-0 and C=O groups shift on

chelation with metal ions. The most characteristic bonds associated with the

(50)

Ex onment.1J

Table 2.03 Elemental analysis of Fe(lll) and Nl(ll) complexes of

Oxalodihydroxamlc acid.

Dihydroxamic Molecular Formula Found (Calcd) %

Acid of metal complex c H N

CONHOH Fe,C6H6N60 12 14.78 1.76 16.83

tONHOH (15.45) (1.28) (18.02)

Oxalodihydroxamic Ni,C6H6 N60 12 14.29 1.62 15.97

Acid (ODHA] (15.27) (1.27) (17.81)

Calculated values are given in parentheses.

Table 2.04 IR Absorption Maxima in em-' of some metal complexes of

Oxalodihydroxamic Acids.

Iron Mode of Nickel Mode of Copper Mode of

Complex vibration Complex vibration Complex vibration

3187 0-H Stretching 3250 0-H Stretching 3444 0-H Stretching

1670 C=O Stretching 2855 0-H Stretching 1675 C=O Stretching

1404 G-N Stretching 2375 C-H Stretching 1617 C=O Stretching

1281 G-H Stretching 2270 C-H Stretching 1487 C-N Stretching

1047 N-H Bending 1629 C=O Stretching 1412 C-H Bending

896 G-H Rocking 1452 C-H Bending 1347 C-H Bending

822 G-H Rocking 1281 C-H Bending 1300 G-H Bending

496 G-H Twisting 999 C-H Rockinig 1026 N-H Bending

866 G-H Rockinig 869 G-H Rocking

701 G-H Twisting 594 G-H Twisting

560 G-H Twisting

488 C-H Twisting

(51)

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hydroxamic acid functional grouping are due to C=O and 0-H stretching vibration and

these can be assigned rather unambiguously. The N-0 and C-N stretching and 0-H

deformation Vibrations were assigned with less certainty because of overlapping by

several other modes of vibrations and because of the nonavailability of systematic data

on the assignment of these bands in the infrared spectra of hydroxamic acids.

Theoretically the bands due to 0-H group in free ligand disappear while bands due to

N-0, C=O and C-N shift on complex formation. Thus the infrared spectroscopic study

will help in elucidating the structure of the ligands as well as of metal complexes.

2.4 KINETIC MEASUREMENTS

For each kinetic run, two reaction vessels were used. One of these contained

appropriate volumes of acid (catalyst) and water; the other one contained the

hydroxamic acid. After thermostating for about 30 minutes the acid solution was

transferred to the reaction vessel containing hydroxamic acid. After the content of the

reaction vessel was shaken, an aliquot of the reaction mixture was withdrawn into a

10·ml volumetric flask containing 2 ml of lron(lll) chloride. A double purpose, quenching

of the reaction and colour development, was thus served. The volume of the coloured

solution was made up to 10 ml and its absorbance was measured at 500-520 nm using

a reference solution containing 2 ml of the same lron(lll) chloride in 10 ml of water. The

kinetic runs were studied generally up to two half-lives. For measuring absorbance, a

Systronics UV-VIS Spectrophotometer type 108 was used. For protonation studies a

Unicam UV2 300 spectrophotometer was used.

For a pseudo first-order reaction, a plot of log absorbance vs. twill give a straight

line of slope, - kw/ 2.303 (where kw is the pseudo first-order rate constant for the

reaction). The rates of hydrolysis were determined spectrophotometrically by following

the decrease in the characteristic absorption of the hydroxamic acid-ferric chloride

complex. As Beer's law is applicable to all the lron(lll) hydroxamic acid complexes, the

concentration of reacting species is proportional to the absorbance A. [log A ~ log (a­

x)]. To obtain the rate constant kw• log (a-x) was plotted against time t; from the slope

of the plot, k, was determined. The fact that a straight line was obtained for all the plots

of (log a-x) vs. t measured in this investigation is, in itself, an indication that the

(55)

tx enmontal

reactions are all first-order in substrate. The initial concentration of the hydroxamic acid

in reaction mixture 1s about 7.0 x 10-' M. The experimental errors in the respective runs

were generally less than 1.0% and the reproducibility of the rate constants was within

±1.5%.

Simple and multivariate linear regression analyses were carried our by the least

square method on a personal computer. The programme for the multivariate linear

regression enabled us to evaluate regression coefficient of the variables, coefficients

of multiple determination and standard deviation etc.

2.5 pK8". AND pK, MEASUREMENTS

An attempt have been made to measure the pK8 ". and pK. values, for study the

acid-base equilibria of hydroxamic acids.

2.51 pK8". Measurements

A weighted amount of compound was prepared in water to prepare the stock

solution of substrate, from which 2.0 ml aliquots was pipetted out into 10 ml volumetric

flasks and make upto mark with suitable acid water mixtures, so as to give solutions

of desired acid concentrations (the final concentration of substrate were around

3 x 1 o-s mol dm-a). All the solutions were cooled in ice-bath before and during the

addition. The resulting cold was then thermostated to 25 ± 0.1 oc. The solutions were

then transferred in to 10 mm glass cells and the UV spectra in the 190-300 nm range

were recorded in spectrophotometer. A reference cell containing acid of the same

strength as that in the reaction cell was used in each case. For protonation studies

Unicam UV2 300 spectrophotometer were used.

2.52 pK. Measurements

pK. of hydroxamic were determined by pH metric titration method and

spectrophotometric method. Procedure for the pH metric titration for determining

ionisation constants (PK.) of hydroxamic acids is basically the same as described be

(56)

Ex onmontal

Albert and Sergeant [14]. Generally 0.01M solution of hydroxamic acids were titrated

against 0.1 M NaOH. The solution was thermostated. Titration was carried out by

adding 0.5 ml of 0.1 M NaOH from the microburette, to the titre solution and the steady

pH of the resultant solution were recorded each time.

The ionization ratio of monohydroxamic acid (BHA) and dihydroxamic acid

(ODHA) in aqueous sodium hydroxide has been determined using UV

spectrophotometric method [15-16]. The spectrum changes appreciably with increase

of sodium hydroxide concentrations. The ionization ratio is alkaline solutions of different

concentration was calculated by using the relation.

[A-] EHA-E

[HA] E-EA_

Ionization constant (PK.) is calculated by the use of following equation.

[A-] pK. = pH - log [HA]

2.6 PRODUCT IDENTIFICATION

Product studies were carried out in solutions identical with those used in

the kinetic measurements, except that the concentrations of the substrates employed

were higher. In moderately concentrated acids, 1.5 g of the substrate was dissolved

in 100 ml of acid and heated at 85°C using a water bath. After reactions had reached

at least 90% completion, aliquots were removed and chilled for product identification.

Hydrolysis products, i.e. oxalic, malonic and succinic acids and hydroxylamine

hydrochloride, were identified qualitatively for dihydroxamic acids. For desferal succinic

acid, acetic acid and 5-amino pentyl hydroxylamine were identified qualitatively by

usual organic tests. These products were separated by fractional crystallization and

purified by recrystallization. The UV spectra of isolated products were compared with

those of authentic samples.

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Ex oomontal

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