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CHAPTER FOUR

Implementation of

Copper (II)-ethylenediaminetetraacetate

Absorbing System for Indirect

Spectrophotometric

Determination of Iron (III)

Determination of Iron by using Copper (II)-ethylenediaminetetraacetate Absorbing System

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CHAPTER FOUR

Implementation of

Copper (II)-ethylenediaminetetraacetate Absorbing

System for Indirect Spectrophotometric

Determination of Iron (III)

Absorbance quenching capacity of Fe(III) towards the copper(II)-

ethylenediaminetetraacetate absorbing system, is implemented for measurement

of permittance in spectrophotometric determination of iron. As the cupric ion is

one of the serious interfering ions in the quantitative determination of Fe(III)

using ethylenediaminetetraacetic acid reagent, therefore accuracy in this

measurement is increased by employing the [Cu(II)-EDTA]2- absorbing system.

The use of this absorbing system allows the determination of Fe(III) in presence

of Cu(II). Similar to earlier study (of determination of bismuth), the stability

constant (log Kf) difference of [Cu(II)-EDTA]2- and [Fe(III)-EDTA]1- chelates is

also considered valuable for quantitative determination of iron. Stability of the

[Fe(III)-EDTA]1-chelate is more than [Cu(II)-EDTA]2- chelate, consequently the

surplus EDTA in the Fe(III) solution is determined by adding excess of Cu(II).

In this experiment, the solution of Fe(III) buffered at pH=1.5, was

initially treated with fixed and excess of EDTA reagent; after quantitative

chelation of Fe(III) as [Fe(III)-EDTA]1- the surplus EDTA was utilized for

generation of [Cu(II)-EDTA]2- absorbing system via adding fixed quantity of

Cu(II) solution. The sample solution with greater concentration of Fe(III) left the

smaller amount of surplus EDTA for generation of absorbing system and vice a

versa. As a result, the color intensity of [Cu(II)-EDTA]2- chelate was observed

inversely proportional and permittance directly proportional to the concentration

of Fe(III). Validity of this method was tested at 722 nm by measuring clearance

of the [Cu(II)-EDTA]2- system and the linearity of logarithm of clearance

(permittance) versus concentration of Fe(III) was used for determination of iron.

The effects of some important variables (such as pH, wavelength of

measurement, concentration of reagent, dilution of test solutions, role of foreign

ions, etc.) on permittance and permittance coefficients of the absorbing system


115

were studied and the average value of permittance coefficient was used for

determination of iron. The method was successfully applied for quantitative

determination of iron in synthetic mixture of cations and pharmaceutical

preparations.

4.1 Introduction and Review of Literature

Iron, the most abundant element in the earth’s crust (5.6 % by mass), is

immensely important both in human civilization and in living systems [1]. Iron is

so widely diffused in nature, in both divalent and trivalent oxidation states

combined as ferrous and ferric compounds [2]. The ferrous iron has light green

color while the ferric iron is yellow in color, but ferric iron produces the red-

colored complex with thiocynate solution while ferrous iron yields no coloration

[3]. Iron plays an important role in biological processes. In living systems iron is

an essential constituent of numerous biomolecules. The body of a healthy human

adult contains about 4 to 5 g of iron, 65% of which is present in hemoglobin and

muscle hemoglobin [1, 4] that is myoglobin [5]. Ferrous ion is the central

structural unit of hemoglobin and myoglobin, performing the function of binding

of molecular oxygen through transferring an electron by self oxidation to ferric

ion [5]. For the deficiency of iron, the ferrous ion (generally in the form of

ferrous ascorbate or ferrous fumarate) being used for medical purposes and can

be taken internally without danger.

Iron compounds found many applications in medicines and also used for

preparation of analytical regents for spectrophotometric [6-11] analyses.

Chemically modified working electrode of iron compounds finds extensive

applications in voltammetric [12-18] analysis (Table 4.1). As the uses of iron in

medicine increases, it has spread in the environment and the chance of exposure

of organisms to iron has increased. Therefore, determination of iron at ultra-trace

levels in environmental and biological samples is important.

Several methods have been developed for determination of iron (Table

4.2). These include direct spectrophotometric [19-24] assay using specific

reagent, simultaneous spectrophotometric determination with other metal [25-26]

analyte, derivative spectrophotometric methods [27-29], photometric titrations

[30-31], normal flow injection spectrophotometric [32-34] methods for

determination of iron, flow injection simultaneous [35-36] spectrophotometric

determination of Fe(III) along with other analyte, flow injection simultaneous


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[37-42] spectrophotometric determination of Fe(II) and Fe(III), reverse flow

injection [43-44] spectrophotometric methods, flow injection luminescence

spectrometry [45], voltammetric methods of determination of iron [46-49] and

total iron[50-53], potentiometric methods [54-55], electrothermal inductively

coupled plasma-optical emission spectrometry [56], titrimetric methods [57] and

indirect spectrophotometric method [58]. However, most of these above

mentioned methods are highly sensitive but are economically unsound since

those require costly sophisticated analytical instruments. Some of the methods

(listed in Table 4.2) need costly and poisonous complexing reagents for detection

and determination of iron in the sample.

The disodium salt of ethylenediaminetetraacetic acid, EDTA has found

considerable use as a standard reagent for determination of numerous metals by

photometric titration. Sweeter and Bricker have reported good photometric

titration of ferric iron with EDTA by using salicylic acid as indicator for ferric

ions [30]. The information of stability constant (log Kf) of [Fe(III)-EDTA]1- and

[Cu(II)-EDTA]2- chelates and their absorption spectra are the parameters utilized

by Underwood for simultaneous determination of iron and copper in a single

photometric titration [31]. Although detection of the exact end point by graphical

means is tedious and time consuming route yet photometric titration methods are

consistently used, since the presence of other substances absorbing at the same

wavelength does not necessarily cause the interference, in as much as only the

change in absorbance is significant [59].

The log Kf value [60-61] of EDTA complexes of iron and copper

are reported as 24.23 and 18.70 respectively, these values are sufficiently larger

indicates both chelates have satisfactorily stability; however, [Fe(III)-EDTA]1- is

more stable than [Cu(II)-EDTA]2-. The wide difference in these stability

constants permits for iron to react with EDTA first in presence of copper

consequently, copper ions functioning as an indicator [59] in iron titration as well

as allows for simultaneous [59] determination of both metals in a single

photometric titration. The same concept of difference in log Kf values of [Fe(III)-

EDTA]1- and [Cu(II)-EDTA]2- complex was exercised for spectrophotometric

determination of iron through taking the advantage of copper ion as an indicator

for determination of surplus EDTA.


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TABLE 4.1

Analytical applications of iron and its compounds in spectrophotometric and

voltammetric studies

Characteristic of the method Ref.

Spectrophotometric determination perphenazine is carried out using

potassium ferricyanide-Fe(III) system. In the presence of potassium

ferricyanide, Fe(III) is reduced to Fe(II) by perphenazine at pH 4.0. The

absorbance of this Prussian blue thus formed is measured at 735 nm.

[6]

The ability of negatively charged phosphatidates to form complexes

with Fe(III) ions is used for spectrophotometric assay of phosphatidic

acid, phosphatidic acid extracts Fe(III) ions from purple colored Fe(III)–

salicylate and decreases the absorbance measured at 490 nm.

[7]

Fe(III) is reduced to Fe(II) by captopril, then the in situ formed Fe(II)

reacts with potassium ferricyanide to give the soluble prussian blue at

pH 4, which is measured at 735 nm.

[8]

The magnetic iron oxide nanoparticles used as an adsorbent for

extraction, preconcentration and determination of traces of fluoride ions.

The determination method is based on the discoloration of Fe(III)–SCN

complex with extracted fluoride ions which was subsequently measured

at 458 nm.

[9]

Ascorbic acid reduces Fe(III) to Fe(II), the reaction of the produced

Fe(II) with 1,10-phenanthroline in a weak acidic medium to form a

colored complex.

[10]

The Fe(III) plus 1,10-phenanthroline reagent was used for determination

of organic compounds, via its reaction with phenolic compounds. By a

redox reaction the reagent forms the chelate [Fe(phen)]+2, which is

extractable as ion-association complexes.

[11]


118

An electrochemical approach was applied for detection of the choline

using prussian blue modified iron phosphate nanostructures. A catalytic

property the nanostructure towards electro-reduction of H2O2 used an

amperometric choline biosensor was used for determination of choline.

[12]

Chemically modified carbon paste electrodes based on Fe(III)-schiff

base containing graphite pastes were prepared and evaluated as

electrochemical sensor for the voltammetric determination of

mefenamic acid and indomethacin in aqueous solutions by differential

pulse voltammetry in at pH 3.5.

[13]

A thiolated 19-mercapture probe was attached to gold coated ferric

oxide nanoparticles and a square wave voltammetric procedure was

applied for the detection of specific hybridization processes, using a

disposable magnetic DNA sensor.

[14]

Fe(II) phthalocyanine modified multi-wall carbon nanotubes paste

electrodes were used as voltammetric sensors to selectively detection of

dopamine in the presence of serotonin (5-HT).

[15]

A carbon-paste electrode modified with iron (II)-phthalocyanine was

used for the sensitive voltammetric determination of epinephrine. The

electrochemical response characteristics of the modified electrode

toward epinephrine, ascorbic acid and uric acid were investigated by

cyclic and differential pulse voltammetry.

[16]

Chemically modified electrode is constructed by doping Fe(III) on

zeolite modified carbon paste, which was used for determination of

tryptophan, uric acid and ascorbic acid at pH 3.5 by application of the

differential pulse voltammetry.

[17]

Titrimetric, potentiostatic coulometric and electrodeposition methods

are employed for determination of total iron, Fe(II), Fe(III) in presence

of copper and copper in the slag sample.

[18]


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TABLE 4.2

Analytical methods of quantitative determination of iron in different types of

samples by various analytical techniques

Characteristic of the method Ref.

A nucleophile formed from Fe(III) oxidized catechol in 0.1 M HCl

couple with o-tolidine, system 1/p-toluidine, system 2 to produce dye

product, measured at 520 nm.

[19]

Fe(II) after complexation with o-phenanthroline and coupled with

picrate anion and extracted by dispersive liquid-liquid microextraction

for spectrophotometric determination of iron in water samples.

[20]

The 2, 4, 6-Trimorpholino 1, 3, 5-triazine gives complex with Fe(III)

which absorbs in the ultraviolet region. The Fe(III) alone or with

cationic surfactants is titrated with this reagent.

[21]

The N, N’-bis (2-hydroxy-5-bromo-benzyl) 1, 2 diaminopropane reacts

with Fe(III) at pH 3–6 to produce a red complex soluble in chloroform

is measured for determination of iron.

[22]

A selective spectrophotometric determination of both Fe(II) and Fe(III),

has been carried out using a melamine-formaldehyde resin for

concentration of Fe(III) at pH 5. The Fe(II) was not retained, which is

oxidised to Fe(III) concentrated on a second resin column. The iron in

both columns was eluted with 1 M HCI solution and separately analyzed

by the 1, 10-phenanthroline-citrate spectrophotometric method.

[23]

Tiron (1, 2-dihydroxybenzene p-3, 5 disulfonic acid) chromophoric

reagent used for determination of Fe(III).

[24]

A multivariate calibration model was developed for the simultaneous

spectrophotometric determination of Al(III) and Fe(III) in

posthemodialysis fluids with pyrocathecol violet as chromogenic

reagent at 580 nm.

[25]


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Metal ions such as Co(II), Ni(II), Cu(II), Fe(III) and Cr(III), present in

electroplating baths, were analysed simultaneously by a spectrophoto-

metric method by the ethylenediaminetetraacetate (EDTA) solution as a

chromogenic reagent.

[26]

The formation of the binary complexes of iron and ruthenium with 4, 7-

diphenyl-1, 10-phenanthroline (bathophenanthroline) in the presence of

ethyleneglycol and ruthenium and iron in mixture were determined by

second derivative spectrophotometric method.

[27]

The complexes of Al(III) and Fe(III) ions with chrome azurol S are

formed at pH 5.5 and these metals are determined by first-derivative

spectra at 675 and 623.5 nm, respectively.

[28]

Copper and iron in mixtures separated by liquid-liquid extraction as

picrate ion pairs. Iron-picrate, reacts in the organic phase with 5-phenyl-

3-(4-phenyl-2-pyridinyl)-l, 2, 4-triazine. The copper-picrate reacts with

2, 9-dimethyl-4, 7-diphenyl-l, 10-phenantroline. The extracts were

evaluated directly by second derivative spectrophotometry.

[29]

Photometric titration of ferric iron with EDTA has been carried out by

using salicylic acid as indicator for ferric ions.

[30]

The difference in the stability constant of [Fe(III)-EDTA]1- and

[Cu(II)-EDTA]2- chelates and their absorption spectra are the parameters

utilized for simultaneous determination of iron and copper in a single

photometric titration.

[31]

The catalytic effect of Fe(III) on the oxidative coupling of p-anisidine

with N, N-dimethylaniline to form a blue compound measured at 735nm

in the presence of H2O2 and 1,10-phenanthroline as an activator, and the

inhibitory effect of EDTA on the Fe(III)-catalyzed reaction.

[32]

Sodium 3-(2-pyridyl)-5, 6-diphenyl-1, 2, 4-triazine-4%, 4-disulphonate

(ferrozine) used for the determination of iron by normal FIA and

reversed FIA.

[33]


121

The production of I3- from the Fe(II) or Fe(III)-induced perbromate-

iodide reaction is monitored at 353nm, by a FIA spectrophotometry.

[34]

Simultaneous determination of iron and zinc in the human hair with 2-

(5-bromo-2-pyridylazo)-5-diethylaminophenol (5-Br-PADAP) using a

pH gradient technique by FIA method.

[35]

The ferroin complex of iron and azomethine-H reaction product used for

the spectrophotometric determination of available iron and boron in soil

extracts, by computer-controlled multi-syringe flow injection systems.

[36]

The simultaneous spectrophotometric determination of Fe(II) and

Fe(III) at 530 nm has been carried out by sulfosalicylic acid and 1,10-

phenanthroline as indicators for ions respectively. EDTA solution is

injected into a carrier stream (HNO3) and EDTA replaces sulfosalicylic

acid, forming a more stable colourless complex with Fe(III), whereas

Fe(II) remains in a complex with 1,10-phenenthroline.

[37]

The flow injection gradient titration in a system of reversed flow is used

for simultaneous spectrophotometric determination of Fe(III) and Fe(II),

at wavelength at 530nm when radiation is absorbed by both complexes.

After injection EDTA replaces sulfosalicylic acid and forms with Fe(III)

more stable colourless complex.

[38]

The iron in tap and natural water samples is acidified to decompose iron

hydroxide and iron-complexes into free iron, Fe(III) and Fe(II). The

amounts of free iron were detected using a catalytic action of Fe(III) and

Fe(II) on the oxidation of N, N-dimethyl-p-phenylenediamine in the

presence of H2O2 at 514 nm.

[39]

The different kinetic-catalytic behaviour of Fe(II) and Fe(III) in the

redox reaction between leucomalacmte green and peroxodisulphate with

and without the presence of the activator l, l0-phenanthroline, are used

for simultaneous spectrophotometric determination Fe(II) and Fe(III).

[40]


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The direct spectrophotometric determination of Fe(III) and Fe(II) by

FIA with acetohydroxamic acid and 1, 10-phenanthroline as reagents.

[41]

Fe(III) is determined in a FIA by passage through reduction before

spectrophotometric detection with l, 10-phenanthroline in citrate buffer,

pH 5.0 and then Fe(II) and total iron are determined by splitting the

injected sample so that a portion passes through the reductor column.

[42]

A simple reversed FIA colourimetric procedure for determination of

Fe(III) in which, the reaction between Fe(III) with chlortetracycline,

resulting in an intense yellow complex measured at 435 nm.

[43]

The reaction between Fe(III) with norfloxacin in ammonium sulfate

solution, gives an intense yellow complex measured at 435 nm by a

reversed FIA colorimetric method.

[44]

A FIA technique used for determination of Fe(II) and total-Fe in water

samples via preconcentration on the Amberlite XAD-4 resin,

functionalized by N-hydroxyethylethylenediamine groups, and

chemiluminescence detection, using brilliant sulfoflavine plus H2O2.

[45]

Differential-pulse cathodic stripping voltammetry with a carbon paste

electrode have been used for determination of iron in 5-aminoiso-

phthalic acid.

[46]

Chemically modified carbon paste electrodes containing 1, 10-

phenanthroline and Nafion as modifiers were used for preconcentration

Fe(II) and determined by differential pulse voltammetry.

[47]

Cathodic stripping voltammetric determination of total Fe has been

carried out by a chitosan modified glassy carbon electrode.

[48]

Brilliant green exhibits a sensitive differential polarographic wave in

dilute H2SO4. Fe(III) has a strong catalytic effect on the slow photo-

chemical reaction of brilliant green under ultraviolet irradiation in the

[49]


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presence of oxalic acid as activator. The changes in concentration of

brilliant green can be monitored by polarography.

Pre-reduction of Fe(III) in the presence of hexacyanoferrate (II) and 1,

10-phenanthroline, to form Fe(II)-1, 10-phenanthroline, on the

unmodified carbon electrode surface and total iron was then determined

voltammetrically by oxidation of the Fe(II)-1,10-phenanthroline.

[50]

The voltammetric characteristics of Fe(III) oxinate at a mercury

electrode, in the presence of 0.2M tributylam monium perchlorate and

0.2M tributylamine as the supporting electrolyte in chloroform have

been used with a two electron quasi-reversible process for the reduction

of Fe(III)-oxinate

[51]

Fe(II) and Fe(III) complexes of pyrophosphate determined simultane-

ously at pH 9 by differential pulse anodic stripping voltammetry.

[52]

A Z-type flow cell is operated simultaneously as an optical flow cell for

the spectrophotometric determination of Fe(III) by monitoring the

absorbance of the sulfosalicylic acid-Fe(III) complex at 505 nm and as

an electrochemical flow cell for amperometric monitoring of the current

due to the oxidation of Fe(II) to Fe(III).

[53]

The potentiometric and voltammetric measurements of Fe(II) and

Fe(III), are performed at ionomer coated glassy carbon electrode.

[54]

Ferroin with a PVC membrane sensor containing ferroin-TPB as an

electroactive component plasticized with 2-nitrophenyl phenyl ether is

used for both batch and FIA of Fe(II) or Fe(III).

[55]

A solid phase extraction technique for the speciation of trace dissolved

Fe(II) and Fe(III) in environmental water samples was used by coupling

micro-column packed with N-benzoyl-N-phenylhydroxylamine loaded

on microcrystalline naphthalene to electrothermal vaporization

inductively coupled plasma-optical emission spectrometry.

[56]


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The metallic iron, ferrous oxide, and ferric oxide are estimated when

present together. The sample is treated with bromine dissolved in

ethanol, and filtered. Iron in the filtrate is titrated iodometrically, and

corresponds to the metallic iron present in the mixture. The oxide

residue is dissolved in HCl in CO2 atmosphere. The Fe(II) formed,

equivalent to FeO present, is titrated with a standard vanadate solution,

and the total iron(III) (= FeO + Fe2O3) in the titrated solution is then

estimated iodometrically.

[57]

An indirect spectrophotometric continuous-flow method was used for

the determination of trace amounts of complexing agents such as

EDTA, DTPA, CyDTA, EDTA-OH, NTA, citrate and pyrophosphate

[58]

In this method, the sample solution of Fe(III) buffered with chloroacetic

acid was treated first with measured and excess of EDTA reagent; after

quantitative chelation of Fe(III) as [Fe(III)-EDTA]1- the surplus EDTA was

utilized for generation of [Cu(II)-EDTA]2- absorbing system via adding measured

quantity of Cu(II) solution. The solution with greater concentration of Fe(III) left

the smaller amount of surplus EDTA (for generation of absorbing system) and

vice a versa. As a result, the color intensity of [Cu(II)-EDTA]2- chelate was

observed inversely proportional to the concentration of iron. Consequently,

permittance [62] of the absorbing system was found directly proportional to the

concentration of iron. Thus, the absorbance quenching [62] action of iron analyte

on to the [Cu(II)-EDTA]2- absorbing system was worked out in alternative

manner for determination of iron.

4.2 Experimental

4.2.1 Instruments and accessories

1. Shimadzu double beam UV-Visible spectrophotometer with the quartz

rectangular cuvettes was used for measurement of %T.

2. Equiptronics pH-meter (EQ-610) was used to check the pH of solutions.

4.2.2 Reagents and chemicals:

All chemicals used were of analytical reagent grade and were used without

further purification. One liter each of the following reagents were prepared.


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(1) 0.0179M Fe(III) solution [viz. 1.0 mg ml-1 of iron] was prepared by

dissolving 8.635 g. of FeNH4(SO4)2.12H2O in minimum quantity of conc.

HNO3, followed by dilution with distilled water containing sufficient HNO3

to make the final solution in 0.5M HNO3. This solution of Fe(III) was

standardized against standard EDTA solution using chloroacetate buffer and

salicylic acid in acetone as an indicator.

(2) 0.05M EDTA solution was prepared by dissolving 18.612 g. of

Na2H2EDTA.2H2O in distilled water.

(3) 0.05M Cu(NO3)2 solution was prepared by dissolving 12.080 g. of

Cu(NO3)2.3H2O in distilled water. The solutions of EDTA and Cu(NO3)2 are

termed here as absorbing system’s reagents.

(4) 1.0 M CH2ClCOOH solution was prepared by dissolving 94.500 g. of

monochloroacetic acid in distilled water. The solution of chloroacetate was

used as buffer.

(5) 0.5M HNO3 solution was prepared by diluting 32.0 ml of conc. HNO3 with

distilled water.

(6) The iron tablets, Livogen-z (contains 50mg of elemental iron) manufactured

by Merck India Ltd., Ferium-xt and Orofer-xT (both tablet contains 100mg of

elemental iron) manufactured by Emcure Pharmaceuticals Ltd. were used for

determination of iron.

3. Analytical Method for Determination of Iron

The method proposed for determination of iron was tested primarily with

the standard solution of Fe(III). The test solutions (TS) of iron in the range of

0.001g to 0.010 g were prepared by adding 1.0 ml, 2.0 ml, …, to 10.0 ml aliquots

of 1.0 mg ml-1 Fe(III) solution sequentially into 50ml graduated flasks each

containing 5.0 ml of 1.0 M CH2ClCOOH as a buffer solution. To equalize the

proton ions concentration, 9.0 ml, 8.0 ml,…, to 0.0 ml aliquots of 0.5M HNO3

were also added in descending order into these volumetric flasks numbered as 2

to 11. After addition of 5.0 ml of 0.05M EDTA solution reaction mixture in the

flask were shaken thoroughly for quantitative chelation of Fe(III) with EDTA and

5.0ml of 0.05M Cu(NO3)2 solution was added for utilization of the surplus EDTA

and generation of [Cu(II)-EDTA]2- absorbing system. The reaction mixtures were

further diluted up to the mark with distilled water. Excluding only Fe(III)


126

solution, the reagent blank (RB) solution was prepared in flask No.1 with 10.0ml

0.5M HNO3. The true blank (TB) or reference solution was also prepared in the

similar way using 14.0 ml of 1.0 mg ml-1 Fe(III) solution. The added 5.0 ml of

0.05M EDTA was completed utilized for chelation of the iron present in this

flask, therefore, [Cu(II)-EDTA]2- absorbing system was not generated in TB

solution, so act as a reference. The visible absorption spectrum of [Cu(II)-

EDTA]2- absorbing system (viz. RB solution) was obtained against the TB

solution, and a wavelength 722 nm was selected for measurement. The %T of RB

as well as each TS was measured at 722 nm against TB as a reference. The %T of

the RB was used for obtaining the clearance value of test solutions. The graph of

logarithm of clearance viz. permittance against the concentration of Fe(III) was

used for determination of iron. An analogous method was employed for

quantitation of iron in Livogen-z, Ferium-xt and Orofer-xt tablets.

4.4 Results and Discussion

The hexadentate chelating reagent, ethylenediaminetetraacetic acid gives

remarkably stable chelates with Fe(III) and Cu(II) in one-to-one stoichiometry

[63]. Reilley and Schmid [64] showed the minimum pH needed for satisfactorily

chelation of various cations with EDTA. The strongly acidic pH (in the range of

1-2) can be accepted by the trivalent metal cation for complexation with EDTA.

Therefore, the quantitative chelation of Fe(III) with EDTA can be attainable at

pH 1-1.5. Wagreich and Harrow [65] also showed that, EDTA has enormous

nucleophilic capability to chelate the cupric ion over a broad pH range and the

stability of [Cu(II)-EDTA]2- chelate is not much affected by the pH of solution.

Even at acidic pH the reaction of Cu(II) with EDTA results with generation of

intensely blue colored [Cu(II)-EDTA]2- chelate; which was employed as the

absorbing system at 745 nm for spectrophotometric determination of copper [30].

The earlier study of quantitative determination of bismuth was carried out at pH

1.0-1.25 via generating the [Cu(II)-EDTA]2- absorbing system and the same

absorbing system was exploited here for spectrophotometric determination of

iron.

4.4.1 Determination of Surplus EDTA

The formation constant (log Kf) value [60-61] of [Fe(III)-EDTA]1- and

[Cu(II)-EDTA]2- chelates are reported 24.23 and 18.70 respectively, which


127

indicates that both chelates have enough stability but iron chelate is much more

stable than copper chelate. The sufficient difference (about 5.53) in these log Kf

values permits to utilized surplus EDTA in the test solution for generation of

[Cu(II)-EDTA]2- absorbing system, without disturbing the stability of iron

chelate. Furthermore, copper chelate exhibits maximum absorption in the visible

region where other species in the test solution exhibits nil absorption.

4.4.2 Effect of [H+] on the Linearity

The stability of [Cu(II)-EDTA]2- absorbing system is not much affected

by small change in pH of the solution [66]. But variable concentration of proton

in test solutions affects the ionization of EDTA that reflects an adverse effect on

the analytical performance of the method. Because variable aliquots (1.0 ml to

10.0 ml) of the standard solution of iron (having concentration 1.0mg ml-1

prepared in 0.5M HNO3) used for preparation test solutions, those carry the

different concentration of proton. Consequently, the solution of 0.5M HNO3 was

added in descending order (from 9.0 ml to 0.0 ml) for compensation of effect of

the H+ ions on the ionization of EDTA. The proposed method involves the

measurement of permittance of test solution (TS) in comparison with permittance

of the reagent blank (RB) solution, so excluding only the sample, the composition

of both solutions was kept essentially identical. Moreover, addition of buffer

species ensures quantitative chelation of both metals through nullification of the

effect of protons released from EDTA during complexation reactions.

4.4.3 Selection of the Wavelength for Measurement

In the proposed method, the quantitative determination of iron was carried

through measurement of permittance of [Cu(II)-EDTA]2- absorbing system. The

optical density of absorbing system indicates the concentration of surplus EDTA

in test solution. Therefore, for attending the greater sensitivity, it was necessary

to carry out measurement at the wavelength to which absorbing system shows

absorption maxima (max). For this purpose, the visible absorption curve of

[Cu(II)-EDTA]2- chelate (viz. RB solution) was obtained against the TB solution

as a reference, both were prepared as described in the method. The spectrum

study showed that, [Cu(II)-EDTA]2- chelate in chloroacetic acid and nitric acid

medium exhibits absorption maxima at 722 nm. Therefore, measurement of %T


128

was carried out at 722 nm to which all other species are transparent, except the

free Cu(II) ions have little absorbancy at this wavelength. Therefore, equivalent

amount of it also added in TB or reference solution for compensating the

background absorbance of unused Cu(II) ions. Though the concentration of

unused/free Cu(II) ions was not identical in all test solutions, but that does not

much affect the analytical linearity of permittance versus concentration of iron.

Only Cu(II) ions have slight absorbancy at 722 nm, therefore, other blank

solution (copper blank) was also prepared simply by diluting with distilled water

(to 50.0 ml) the experimental volume of cupric nitrate, nitric acid and

chloroacetic acid solutions. The absorption spectrum of the RB solution when

obtained against copper blank, then also absorbing system proved its absorption

maxima at 722 nm. The results obtained against true blank do not differ

significantly from those obtained against copper blank. So, all the measurements

were carried out using free copper ions solution as a blank or reference.

4.4.4 Analytical Performance of the Method

The test solutions of iron prepared as mentioned in method confirms that,

at the measured and excess concentration of EDTA, the optical density of

[Cu(II)-EDTA]2- system was directly proportional to the concentration of surplus

EDTA and inversely proportional to the concentration of iron. Because, the

extent of chelation reaction that occurred between Fe(III) and excess of EDTA is

governed by only the concentration of Fe(III). Therefore, the concentration of

surplus EDTA left after complexation of Fe(III) was inversely proportional to the

concentration of iron and hence permittance of test solutions was observed

directly proportional to the concentration of iron. After measurement of %T at

722 nm of [Cu(II)-EDTA]2- absorbing system (each TS and RB) against TB as a

reference, the clearance of test solution was calculated and the logarithm of

clearance (permittance) of the absorbing system was determined by using

following eq.(1). It was observed that, permittance of the test solution is linear

function of concentration of iron and was dependent only on the concentration of

iron. Therefore, the linearity of permittance against the concentration of iron was

used for construction of calibration curve. With 1.0 cm optical path of absorbing

system the proportionality constant determined for determination of iron.


129

4.4.5 Permittance and Permittance Coefficient

The permittance (Pr) of the test solution was calculated by using following

equation (1)

%TTSPr = log --------- --- (1)

%TRB

In this equation, %TTS and %TRB designates the percent transmittance of the test

solution and reagent blank solution respectively, measured against the same

reference solution. The relationship between permittance (Pr) of test solution and

concentration (c) of quencher analyte in it was reported in earlier study and was

used here for determination of proportionality constant/permittance coefficient

(a') of 1.0cm path length (b) absorbing system.

Pra' = ------- (concentration of quencher analyte in g L-1) --- (2)

b c

In this experiment, the quantitative determination of iron in the range of

1.0mg to 10.0mg was carried out at 722 nm by using different volumes (5.0ml,

6.0ml, 7.0ml and 8.0ml) of the absorbing system’s reagents. The copper blank

solution used as a reference for these measurements. Permittance value of the test

solutions was observed directly proportional to the concentration of iron (Table

4.3) and even if the concentration of reagents was altered, the permittance is

constant for a fixed concentration of iron. Similarly to earlier the study in this

experiment it is observed that, at a fixed wavelength, permittance is dependent

only on concentration of the analyte and is independent on the concentration of

absorbing system’s reagents. At 722 nm, the permittance coefficient at every

different concentration of Fe(III) as well as the volume of reagents is also nearly

constant.

The values of permittance (Table 4.3) are little bit more at higher

concentration of iron because, the yellow color intensity of ferric-EDTA chelate

fades the blue color intensity copper-EDTA chelate (have low concentration at

higher concentration of iron) through generating the green color to test solution.

This effect also increases value of the permittance coefficient corresponding

more, so the average value of permittance coefficient was considered for

determination of analyte. In these determinations (Table 4.3) the average value of


130

permittance coefficient was found equal to 0.5168 lit.g-1cm-1 and which was for

determination concentration of the stock solution of iron (1.0 mg ml-1) from

which the test solutions were prepared.

4.4.6 Magnitude of Permittance and Permittance Coefficient

The concentration/volume of reagents does not affect the permittance as

well as permittance coefficient (Table 4.3) but dilution (of the test solutions)

factor shows pronounced effect on permittance. This study was carried out

through quantitative determination of iron in the range of 1.0 mg to 10.0 mg

(with the 5.0ml volume of the reagents) at the 25 ml, 50 ml and 100 ml final

dilution of test solutions. The result of these assay (is summarized in (Table 4.4),

elucidates that, the value of proportionality constant [calculated with eq. (2)] was

observed decreasing with increase in final dilution volume of test solutions. This

is because, the number of absorbing species per unit path length in the solution

decreases with dilution. For result reported in Table 4.2, the average

proportionality constants are 20.656 at 25.0 ml, 10.338 at 50 ml and 5.171 at 100

ml final dilution. From this it is clear that, when dilution volume of test solutions

is doubled, the proportionality constant was decrease nearly equal to half but at

one liter dilution is permittance coefficient (measured in lit.g.-1cm-1) was

observed unchanged.

The second important factor which affects the magnitude of permittance

and permittance coefficient is the wavelength selected for measurement. The

wavelength determines the extent of absorption and hence the readings of %T;

consequently, execute the marked effect on the values of these two parameters.

The [Cu(II)-EDTA]2- absorbing system in chloroacetic acid and nitric acid

medium was showed the max at 722nm. Along with max wavelength, when

same test solutions were measured at other wavelengths that generates the

different values for permittance and permittance coefficient. That means the

magnitude of permittance and also permittance coefficient (a') is absolutely

administrated by the wavelength selected for analysis. The magnitude of both of

these parameters was observed maximum at system’s max wavelength.

Therefore, for achieving the greater sensitivity for the method, measurements

were carried out at the absorption maxima (722 nm) of the absorbing system.


131

TABLE 4.3

Effect of concentration of reagents on permittance and permittance coefficient;

results obtained in quantitative determination of iron at 50 ml dilution with

different volume of absorbing system’s reagents

5.0ml 6.0mlIron(g)

Pr Pr. Coeff. Pr Pr. Coeff.

0.000 0.0000 --- 0.0000 ---

0.001 0.0103 0.5153 0.0102 0.5114

0.002 0.0204 0.5108 0.0205 0.5121

0.003 0.0307 0.5125 0.0310 0.5160

0.004 0.0415 0.5183 0.0409 0.5116

0.005 0.0515 0.5152 0.0518 0.5179

0.006 0.0614 0.5117 0.0614 0.5119

0.007 0.0722 0.5154 0.0727 0.5191

0.008 0.0836 0.5225 0.0833 0.5206

0.009 0.0939 0.5216 0.0941 0.5225

0.010 0.1048 0.5241 0.1048 0.5239

Average value: 0.51674 -- 0.51669

TABLE 4.3 Continues…

7.0ml 8.0mlIron(g)

Pr Pr. Coeff. Pr Pr. Coeff.

0.000 0.0000 --- 0.0000 ---

0.001 0.0102 0.5117 0.0103 0.5136

0.002 0.0206 0.5148 0.0204 0.5096

0.003 0.0307 0.5118 0.0309 0.5143

0.004 0.0414 0.5177 0.0409 0.5117

0.005 0.0515 0.5146 0.0517 0.5171

0.006 0.0617 0.5145 0.0624 0.5202

0.007 0.0724 0.5172 0.0722 0.5158

0.008 0.0835 0.5220 0.0836 0.5222

0.009 0.0941 0.5227 0.0939 0.5217

0.010 0.1045 0.5227 0.1044 0.5222

Average value: 0.51697 -- 0.51684


132

TABLE 4.4

Effect of final dilution of test solutions; results obtained in quantitative

determination of iron in the range of 1.0 mg to 10.0 mg using 5.0 ml of 0.05M

EDTA and 5.0 ml of 0.05M Cu(NO3)2 reagents at different dilutions

25ml dilution 50ml dilutionIron(g)

Pr. Pro.Const. Pr.Coeff. Pr. Pro.Const. Pr.Coeff.0.000 0.000 -- -- 0.000 -- --

0.001 0.0206 20.62 0.5155 0.0103 10.31 0.5153

0.002 0.0413 20.65 0.5163 0.0205 10.25 0.5124

0.003 0.0619 20.64 0.5160 0.0309 10.31 0.5154

0.004 0.0826 20.65 0.5161 0.0415 10.37 0.5183

0.005 0.1030 20.59 0.5148 0.0513 10.26 0.5129

0.006 0.1237 20.61 0.5152 0.0615 10.25 0.5127

0.007 0.1445 20.64 0.5160 0.0722 10.31 0.5154

0.008 0.1653 20.67 0.5166 0.0831 10.39 0.5196

0.009 0.1863 20.70 0.5175 0.0941 10.46 0.5230

0.010 0.2081 20.81 0.5202 0.1048 10.48 0.5241

Average value: 20.656 0.51641 -- 10.338 0.51691


100ml dilutionIron(g)

Pr. Pro. Const. Pr. Coeff.0.000 0.000 -- 0.000

0.001 0.0052 5.17 0.5172

0.002 0.0103 5.14 0.5142

0.003 0.0155 5.16 0.5163

0.004 0.0207 5.17 0.5171

0.005 0.0258 5.16 0.5163

0.006 0.0309 5.15 0.5148

0.007 0.0362 5.18 0.5178

0.008 0.0415 5.19 0.5192

0.009 0.0466 5.18 0.5180

0.010 0.0520 5.20 0.5202

Average value: 5.171 0.51712


133

4.4.7 Determination of Iron in Iron Tablets

The proposed method was applied for determination of iron in Livogen-z,

Ferium-xt and Orofer-xT tablets. The sample (two tablets of Livogen-z, one

tablet of Ferium-xt, or one tablet of Orofer-xT) was heated with 15 ml of conc.

HCl followed by addition of 5 ml of conc. HNO3. The organic matter was

destroyed by treatment with 5-6 ml of 70% perchloric acid. (WARNING: Boiling

perchloric acid can result in serious explosions). The solution was slowly heated

in fuming hood for about 20-25 minutes; at this stage maximum of the acid fumes

were ceased. With the addition of 10ml of distilled water, the sample solution

was again boiled for 10 minutes. The solution was cooled and diluted nearly to

90ml with distilled water. With drop wise addition of conc. HNO3, the pH of

sample solution was adjusted equal to pH (=0.48) of standard solution containing

1.0mg ml-1 iron. The sample solution was filtered after dilution to 100ml and the

amount of iron was determined by the recommended method. The results of this

determination are represented in Table 4.5.

The iron tablets contain ferrous iron generally in the form of ferrous

ascorbate or ferrous fumarate. When the tablets were digested with these acids,

the ferrous iron gets oxidized to ferric iron. The concentration of iron in the

sample solutions of tablet (two tablets of Livogen-z, one tablet of Ferium-xt, or

one tablet of Orofer-xT) at 100ml dilution is 1.0mg ml-1. Therefore, 1.0 ml to

10.0 ml of sample aliquots were tested for determination of iron in the range of

1.0 mg to10.0 mg and the value of the permittance coefficient thus obtained was

used for determination of concentration of iron in stock solution of iron tablets.

That is, the accuracy of method was studied with permittance coefficient; values

obtained in the sample analysis are compared with those obtained in the analysis

of standard Fe(III) solution of same strength and same pH.

4.4.8 Interferences in the Determination of Iron

The type of interference can be predicted form the formation constant of

the EDTA chelates of the metal cations. Therefore, the interfering cations can be

classified into two groups. The first group is composed of those cations whose

EDTA chelates in acidic medium are sufficiently stable comparative to iron

chelate. The EDTA chelate of Bi(III) is nearly stable as Fe(III) chelate,

accordingly Bi(III) is the serious interfering cation in determination of iron.


134

TABLE 4.5

Results obtained at 722 nm in the quantitative determination of iron in Livogen-

Z, Ferium-xt and Orofer-xT tablets, permittance coefficient values are compared

for determination concentration of sample

Volume of

0.05M EDTA and

0.05M Cu(NO3)2

Iron sample

analyzed

Obs. Pr.

coefficient

(lit. g-1cm-1)

Conc. of

iron found

(mg ml-1)

Error

in %

1.0mg ml-1 Fe(III) 0.51674 1.0000 0.00

2 Livogen-Z 0.51673 0.9999 0.01

1 Ferium-xt 0.51697 1.0005 0.05

5.0 ml each

1 Orofer-xT 0.51666 0.9998 0.02

1.0mg ml-1 Fe(III) 0.51669 1.0000 0.00

2 Livogen-Z 0.51691 1.0004 0.04

1 Ferium-xt 0.51673 1.0001 0.01

6.0 ml each

1 Orofer-xT 0.51696 1.0005 0.05

1.0mg ml-1 Fe(III) 0.51697 1.0000 0.00

2 Livogen-Z 0.51697 1.0000 0.00

1 Ferium-xt 0.51683 0.9999 0.01

7.0 ml each

1 Orofer-xT 0.51678 0.9996 0.04

The second type of interfering ions includes those cations which do not

compete with iron but with copper for the EDTA. These are the cations which

forms more stable chelate than copper chelate. At such strongly acidic pH 1.15,

no any divalent metal forms more stable chelate than copper chelate. The cations

whose EDTA chelate is less stable than copper chelate, particularly the

aluminum, barium, calcium, cadmium, lead, magnesium, manganese, zinc and as

well the copper does not interferes in iron determination. The interference study

of copper was not carried out in this experiment because that was added in excess

for generation of absorbing system. The interference study was carried out by

adding 10 mg, 20 mg, 30 mg or 40 mg of these metal cations separately in to the

test solutions containing 1.0 mg to 10.0 mg of iron. The interfering cation of

same concentration was also added in true blank solution but not in the reagent

blank solution. The result of the interference study is reported in Table 4.6, in the


135

form of permittance coefficient. The values of permittance coefficient in absence

(only with standard 1.0 mg ml-1 Fe(III)) and in presence (with standard 1.0 mg

ml-1 Fe(III) plus the added cation) of interference are observed nearly same. At

pH 1.15, all of these metal cations up to 40 mg do not interfere in the

determination of iron. When the added metal cations interfere in the

determination which increases the %T reading of test solutions, consequently that

raises the values of permittance as well permittance coefficient. The

concentration of added cation when reaches in excess, that increases the optical

density of the test solutions with respect to regent blank solution. Hence, the %T

readings were observed decreased and that decreases the values of permittance

and permittance coefficient. In the previous literature the inferences study was

also completed.

TABLE 4.6

Effect of cations on determination of iron, results obtained in the determination of

iron from 1.0 mg to 10.0 mg, the value permittance coefficient was determined at

722 nm in absence and in presence of different concentration of cations

Cation

added as

interference

Amount

added

(mg)

Obs. Pr.

coefficient

(lit.g-1cm-1)

Conc. of iron

solution found

(mg ml-1)

Error

in %

Standard 0.00 0.51690 1.0000 0.00

Aluminum 10.0 0.51696 1.0002 0.02

20.0 0.51695 1.0001 0.01

30.0 0.51690 1.0000 0.00

40.0 0.51697 0.9994 0.06

Barium 10.0 0.51679 0.9998 0.02

20.0 0.51705 1.0003 0.03

30.0 0.51684 0.9999 0.01

40.0 0.51683 0.9998 0.02

Calcium 10.0 0.51683 0.9998 0.02

20.0 0.51696 1.0002 0.02

30.0 0.51663 0.9995 0.05

40.0 0.51700 1.0002 0.02


136


Cadmium 10.0 0.51676 0.9997 0.03

20.0 0.51704 1.0003 0.03

30.0 0.51683 0.9998 0.02

40.0 0.51688 0.9999 0.01

Lead 10.0 0.51689 0.9999 0.01

20.0 0.51695 1.0001 0.01

30.0 0.51677 0.9997 0.03

40.0 0.51704 1.0003 0.03

Magnesium 10.0 0.51676 0.9997 0.03

20.0 0.51704 1.0003 0.03

30.0 0.51683 0.9998 0.02

40.0 0.51688 0.9999 0.01

Manganese 10.0 0.51696 1.0002 0.02

20.0 0.51695 1.0001 0.01

30.0 0.51690 1.0000 0.00

40.0 0.51677 0.9997 0.03

Zinc 10.0 0.51683 0.9998 0.02

20.0 0.51680 0.9998 0.02

30.0 0.51690 1.0001 0.01

40.0 0.51687 0.9999 0.01

4.5 Conclusions

The method described here for determination iron is based on the

measurement of concentration of surplus EDTA through generating the [Cu(II)-

EDTA]-2 absorbing system. The absorbing system was found excellent (for

determination of iron) because of its formation and stability at strongly acidic pH

1.15 to which quantitatively chelation of trivalent Fe(III) was attained. At this

pH, many divalent metal cations do complexed by EDTA, this makes the process

selective. The sufficient difference in the stability constant of [Fe(III)-EDTA]-1

chelate and [Cu(II)-EDTA]-2 chelate, the stability of iron-EDTA chelate is not

affected because of the addition of relative larger amount of Cu(II)ions. For the


137

same reason that, copper ions does not interferes in the determination of iron. The

linearity between permittance and the concentration of iron was maintained even

at low concentration (1.0 mg) of iron. The lower concentration of iron (less than

1.0 mg) is not determined in this experiment, but it is possible through increasing

the sensitivity of the method by decreasing the volume/concentration of reagents;

since permittance and permittance coefficient are not preside over the

concentration of reagents. The proposed method is very simple, easy to execute

and which proved to be a better method as compared to other methods involve

the step of preconcentration of analyte. The sensitive of the method practiced

here was found up to 1.0mg of iron, when determined with 5.0ml of 0.05M

EDTA and 5.0ml of 0.1M Cu(NO3)2 and produces the reproducible results with

good accuracy. The method also found excellent over the photometric titration of

iron with EDTA since, the graphical method of determination of exact end point

is tedious and time consuming process. When the proportionality constant at

specific dilution (or permittance coefficient) is determined for standard solutions

of known concentration, in this method the concentration analyte can be

determined directly with eq.(1) reported in earlier study. The proposed method

also neglects a step of standardization of reagents, since it involves the

measurement of permittance of TS with respect to RB. Excluding only the

analyte, the composition of the reagent blank must be identical in every respect to

test solutions. This is the only care have to be taken for the good linearity. The

results reported in the Table 4.5 showed that the method is excellent for the

determination of iron in iron tablets.

References Cited

[1] J. McMurry, R. C. Fay, Chemistry, 4th Ed. Pearson Education Singapore, 874

(2004).

[2] N. H Furman, ‘Standard Methods of Chemical Analysis’, 6th Ed., vol. 1, Ch.

7, E. Robert Krieger Publishing Comp. Malabar, Florida, 191-208 (1962).

[3] H. E. Bent, C. L. French, J. Am. Chem. Soc. 63, 568 (1947).

[4] H. Varley, ‘Practical Clinical Biochemistry’,4th Ed., Heinemann Professional

Publishing Ltd., London, 871 (1969).


138

[5] D. F. Shriver, P. W. Atkins, Inorganic Chemistry’, 3rd Ed. Oxford University

Press, New York, 649 (1999).

[6] L. Guo, Y. Zhang, Q. Li, Spectrochimica Acta Part A, 74 , 307-311(2009).

[7] M. Dippe, R. Ulbrich-Hofmann, Anal. Biochemistry, 392, 169-173 (2009).

[8] S. L. Wang, M. Wang, Q. M. Li, Chinese Chemical Letters, 20, 88-91 (2009).

[9] H. Parham, N. Rahbar, Talanta, 80, 664–669(2009).

[10] M. Zenki, A. Tanishita, T. Yokoyama, Talanta, 64, 1273-1277(2004).

[11] S. Koch, G. Ackermann, P. Lindner, Talanta, 39, 693-696, (1992)

[12] H. Zhang, Y. Yin, P. Wu, C. Cai, Biosensors and Bioelectronics, 31, 244-

250 (2012).

[13] M. Hasanzadeh, N. Shadjou, L. Saghatforoush, J. E. Nazhad Dolatabadi,

Colloids and Surfaces B: Biointerfaces 92, 91-97 (2012).

[14] O. A. Loaiza, E. Jubete, E. Ochoteco, G. Cabanero, H. Grande, J.

Rodríguez, Biosensors and Bioelectronics, 26, 2194-2200 (2011).

[15] D. Patrascu, I. David, V. David, C. Mihailciuc, I. Stamatin, J. Ciurea, L.

Nagy, G. Nagy, A. A. Ciucu, Sensors and Actuators B, 156, 731-736 (2011).

[16] S. Shahrokhian, M. Ghalkhani, M. K. Amini, Sensors and Actuators B, 137,

669-675 (2009).

[17] A. Babaei, M. Zendehdel, B. Khalilzadeh, A. Taheri, Colloids and Surfaces

B: Biointerfaces, 66, 226-232 (2008).

[18] R. L. Altman, Anal. Chim. Acta, 63 129-138 (1973).

[19] B. Shyla, C. V. Bhaskar, G. Nagendrappa, Spectrochimica Acta Part A,

86,152-158 (2012).

[20] A. B. Tabrizi, J. Hazardous Materials, 183, 688-693 (2010).

[21] N. Meddourene, T. Douadi, S. Chafaa, M. Khan, G. Bouet, C. R. Chimie 7,

1113–1118 (2004).

[22] D. Karaa, M. Alkanb, Talanta, 55, 415-423 (2001).

[23] H. Filik, B. D. Ozttirk, M. Douta, G. Gumas, R. Apak, Talanta, 44, 877-884

(1997).

[24] W. A. Bashir, Microchemical Journal, 26, 477-480 (1981).

[25] P. C. Nascimento, C. L. Jost, M. V. Guterres, L. D. Del’ Fabro, L. M. de

Carvalho, D. Bohrer, Talanta, 70, 540-545 (2006).

[26] Y. Ni, S. Chen, S. Kokot, Anal. Chim Acta, 463, 305-316 (2002).

[27] M. I. Torala, P. Richtera, A. E. Tapiab, J. Hernandezb, Talanta, 50, 183-191

(1999).


139

[28] N. Aguerssif, M. Benamor, M. Kachbi, M. T. Draa, J. Trace Elements in

Medicine and Biology 22,175-182 (2008).

[29] M. I. Toral, P. Richter, C. Rodriguez, Talanta, 45, 147-153 (1997).

[30] P. B. Sweeter, C. E. Bricker, Anal. Chem., 25, 253-255(1953).

[31] A. L. Underwood, Anal. Chem., 25, 1910 (1953).

[32] T. Watanabe, N. Teshima, S. Nakano, T. Kawashima, Anal. Chim. Acta,

374, 303-307(1998).

[33] M. I. Pascual-Reguera, I. Ortega-Carmona, A. Molina-Diaz, Talanta, 44,

1793-1801 (1997).

[34] T. D. Yerian, T. P. Hadjiioannou, G. D. Christian, Talanta, 33, 547 (1986).

[35] S. Zhao, X. Xia, G. Yu, B. Yang, Talanta, 46, 845–850 (1998).

[36] D. M. C. Gomes, M. A. Segundo, J. L. F. C. Lima, A. O. S. S. Rangel,

Talanta, 66, 703-711 (2005).

[37] J. Kozak, N. Jodłowska, M. Kozak, P. Koscielniak, Anal. Chim Acta, 702,

213-217 (2011).

[38] J. Kozak, J. Gutowski, M. Kozak, M. Wieczorek, P. Koscielniak, Anal.

Chim Acta, 668, 8-12 (2010).

[39] S. Lunvongsa, M. Oshima, S. Motomizu, Talanta, 68 969-973 (2006).

[40] H. Muller, V. Muller, Anal. Chim. Acta, 230, 113-123 (1990).

[41] A. T. Senior, J. D. Glennon, Anal. Chim. Acta, 196, 333-336 (1987)

[42] A. T. Faizullah, A. Townshend, Anal. Chim. Acta, 167, 225-231 (1986).

[43] W. Ruengsitagoon, Talanta, 74 1236-1241 (2008)

[44] T. Pojanagaroon, S. Watanesk, V. Rattanaphani, S. Liawrungrath, Talanta,

58, 1293-1300 (2002).

[45] S. Hirataa, H. Yoshiharab, M. Aiharab, Talanta, 49, 1059-1067 (1999).

[46] A. Komersova, M. Bartos, K. Kalcher, K. Vytras, J. of Pharma. Biomedical

Anal., 16 (1998) 1373–1379.

[47] Z. Gao, P. Li, G. Wang, Z. Zhao, Anal. Chim. Acta, 241, 137-146 (1990).

[48] G. Lu, X. Yao, X. Wu, T. Zhan, Microchemical Journal, 69, 81-87 (2001).

[49] Z. Jiang, X. Liu, M. Zhao, W. Mo, Anal Chim Acta, 354, 359-363 (1997).

[50] J. Jezek, J. W. Dilleen , B. G. D. Haggett, A. G. Fogg, B. J. Birch, Talanta,

71, 202-207(2007).

[51] M. H. Pournaghi-Azar, B. M. Fatemi, Microchemical Journal, 65, 199-207

(2000).

[52] J. F. van Staden, M. C. Matoetoe, Anal. Chim. Acta, 376, 325-330 (1998).


140

[53] B. Haghighi, A. Safavi, Anal. Chim. Acta, 354, 43-50 (1997).

[54] P. Ugo, L.M. Moretto, A. De Boni, P. Scopece, G. A. Mazzocchin, Anal

Chim Acta, 474, 147-160 (2002)

[55] Saad, S. M. Hassan, Sayed A. M. Marzouk, Talanta, 41, 891-899 (1994).

[56] C. Xiong, Z. Jiang, B. Hu, Anal Chim Acta, 559, 113-119 (2006).

[57] B. R. Sant, T. P. Prasad, Talanta, 15, 1483 to 1486 (1968)

[58] H. Itabashl, K. Umetsu, N. Teshlma, K. Satoh, T. Kawashlma, Anal. Chim.

Acta, 261, 213-218 (1992).

[59] R. F. Goddu, D. N. Hume, Anal. Chem., 26, 1740 (1954).

[60] J. G. Speight, ‘Lange’s Handbook of Chemistry’ 16th Ed., McGraw-Hill,

New York, 4.152 (2005).

[61] P. Patnaik, ‘Dean’s Analytical Chemistry’, 2nd Ed., McGraw-Hill, New

York, 2.6-2.7 (2004).

[62] S. R. Labhade, V. B. Gaikwad, Asian J. Chemistry, 21, 5285-5294 (2009)

[63] D. A. Skoog, D. M. West, F. J. Holler, S. R. Crouch, ‘Fundamentals of

analytical Chemistry’ 8th Ed. Thomson Learning 458-464 (2004).

[64] C. N. Reilley, R. W. Schmid, Anal. Chem., 30, 947 (1958).

[65] H. Wagreich, B. Harrow, Anal. Chem., 25, 1925 (1953).

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