Amavadin and homologues as promoters of technological ...

1

Amavadin and homologues as promoters of technological

applications

Lúcia Dias, José Ferreira e José A.L. da Silva

Centro de Química Estrutural, Complexo I, Instituto Superior Técnico, Universidade de Lisboa, Av. Rovisco Pais, 1049-001

Lisboa, Portugal

Abstract

Study of the reaction of two complexes homologues of amavadin with the ligands, HIDA and HIDPA with

the following oxidants: persulfate, iodate, bromate, chlorate, periodate, perchlorate, hypochlorite, nitrite, tellurate

and bismuthate, by UV-Vis spectroscopy in the range 350-850 nm. The reaction of the complex with HIDA was

faster than with the other complex, but always with decomposition. The light increases for the complex with

HIDPA the rate of reaction, but with decomposition. In contrast, in the absence of light the complex is

regenerated. A kinetic study was separately carried out for the oxidation and reduction of the complex with HIDPA

with all oxidants, except hypochlorite, perchlorate, tellurate and bismuthate. A second order reaction was

determined for all cases excluding persulfate, zero order for reduction, and chlorate for both reactions. The rate

constants showed reduction is slower than the oxidation. The values for the oxidation are 1,15X10-6

M/s (chlorate)

4,87X10-2

L/mol.s (iodate) and 6,14X10-2

L/mol.s (persulfate) and for the reduction 3,64x10-7

M/s (persulfate),

1,35X10-6

M/s (chlorate), 1,97X10-2

L/mol.s (nitrite), 1,99X10-2

L / mol.s (iodate), 2,38X10-2

L / mol.s (bromate)

and 6,69X10-2

L/mol.s (periodate).

Keywords: Amavadin, oxidants, HIDPA, HIDA, light, kinetic

1. Introduction

Amavadin is a natural complex ion of

vanadium, isolated in 1972 by Bayer and

Kneifel from the fungus Amanita muscaria that

contains a heterogeneous distribution of

vanadium with higher concentration in the

base[1]. The same authors proposed a

structure as a complex of S,S-N-

hydroxyiminodiproprionic acid ( ) with

a vanadyl ion in a metal:ligand proportion of

1:2, this was later proved wrong by X-ray

scattering [2]. An alternative connection was

suggested based on Wienghardt once

hydroxylamines could coordinate laterally to

vanadium[3]. It was then proposed by Bayer

and co-workers a new structure in which

hydroxylamines are ionized and coordinated to

the metal (figure 1)[4]. The new structure was

confirmed by crystallographic studies of an

amavadin model, , and by

crystallographic studies and NMR of oxidized

and reduced species of amavadin[2].

Amavadin's structure was known in 1999 by

precipitation with calcium ion[5]. It is a

vanadium complex dianionic with symmetry,

distorted dodecahedral geometry and five

chiral centers, one in the metal and the other

four in the carbon atoms bonded to a nitrogen

atom and also exhibiting S stereochemistry

with two possible stereoisomer in the

ligands[6]. Some of its physico-chemical

properties are presented in table 1.

2

Figure 1: Amavadin[2]

Table 1: Amavadin physico-chemical properties[6,7]

Soluble in: Water, DMSO, DMF and acetone

Insoluble in: 1-propanol, 2-propanol, 1-butanol, diethyl ether, nitrobenzene, THF e toluene

ν ( = 985 [misinterpreted as ν (V=O)]

Stability constant log =23(1)

Redox Couple , E/V vs NHE: 0,27 (DMSO), 0,81 (water, pH 7)

) :(1 0,5)X

Studies have shown that amavadin can

electrocatalitically oxidize tiols with a carboxylic

or ester group [6]. It exhibits a Michaelis-

Menten mechanism with a formation of an

intermediary with the substrate[6]. It acts as

catalase after studies with releasing

[6]. Later using Ce4+

as oxidant it was proved

that amavadin models mediates the oxidation

of water[8]. Amavadin is a complex with

applications in industry and organic

synthesis[6]. The complex and its models can

be used as catalysts for hydroxylation,

oxygenation, peroxidative halogenation of

alkanes and benzene, peroxidative

oxygenation of benzene and mesitylene and

carboxilation of linear and cyclo alkanes.

Moreover, amavadin can participate in

oxidation of alcohols and cyanide addition to

aldehydes[6]. It can also be used as catalyst in

a new process of conversion of methane in

acetic acid with as oxidant and yields

between 20%-30% in the absence of CO or

[9]. Its biological function is still unknown

but it is probably related to redox properties as

an electron mediator[10]. The focus of this

work is the biological function of amavadin and

possible synthetic applications by analyzing its

oxidation and reduction (leading to the water

oxidation) with several oxidants.

2. Experimental section

2.1. Synthesis of proligands

The proligands, HIDPA, as a racemic

mixture (compare with figure 2a), and HIDA

(figure 2b), were synthesized according to

Bayer and Kneiffel[11] having been the zinc

salts of both previously prepared using the

procedure of those authors[11]. 1.2 g were

weighed and dissolved in ca. 20 mL of distilled

water and 1 mL of concentrated HCl until the

solution became transparent. The solution was

introduced in a exchange column

containing and washed with distilled water

(to retrieve excess acid and prevent the

formation of a salt with the eluent that could

contaminate the product) and then eluted with

a 0.2 M NaOH solution; the effluent was

collected in test tubes (each one with ca. 20

mL), the first eight were rejected once they do

not contain product. The content of the

selected tubes was transferred to a flask of

100 mL. Evaporation of the solvent was

achieved with a rotary evaporator at less than

40 ℃ in order to avoid decomposition of the

compounds, to a final volume of about 1 mL. In

both cases a white product by evaporation of

the solvent in a vacuum line was obtained.

Figure 2: a) S,S-HIDPA b)HIDA[6]

3

2.2. Preparation of vanadium

complexes

The complexes were prepared in situ with

0.01 M concentration in a volume of 4 mL from

0.0101 g (4X10-5

mol) vanadyl sulfate

pentahydrated and 0.0142 g (8X10-5

mol) of

HIDPA or 0.0119 g (8X10-5

mol) of HIDA, then

a bluish purple solution for the first one and to

the second more bluish, were formed,

respectively. The initial pH of the solutions was

ca. 1,72.

2.3. Sample preparation

Samples were prepared with deoxygenated

water (prepared from distilled water, typically

after at least three consecutive cycles under

vacuum and nitrogen atmosphere) in order that

the presence of molecular oxygen would result

from the oxidation of water and not from

dissolved in the solution. After preparation of

the solution of the complex an oxidant was

added to the solution (persulfate, iodate,

bromate, chlorate, periodate, perchlorate,

hypochlorite, nitrite, tellurate or bismuthate)

with a concentration of 0,005 M and 0,03 M, for

the last two.

In some of the studies light radiation was

avoided by covering the samples with

aluminum foil (designated by “covered” in

contrast with the “transparent” that was

uncovered).

The samples were analyzed by UV-Vis

spectroscopy using a Perkin-Elmer

spectrophotometer in the wavelength range of

350 to 850 nm. The spectra were run to red,

purple and blue solutions, resulting from

oxidation for the first two and decomposition

for the last one. The oxidation of water was

checked by the oxygen meter in SG9-SevenGo

solution calibrated with a saturated solution of

distilled water in air. The pH measurements

were performed with a Metrohm 827 pH lab.

2.4. Tests

Some tests were carried out to identify the

mechanism of the reactions. A test to

determine by forming barium carbonate

precipitate from the reaction of carbon dioxide

in a solution of 1 g of barium nitrate in 20 mL of

deoxygenated water. In another flask with a

solution of vanadium complex (0.01M)

containing 0.029 g HIDA and 0.0249 g of

oxovanadium(IV) sulfate pentahydrate, in 1 mL

of deoxygenated water to which was added

0.013 g of sodium nitrite. Both flasks were

connected and vacuum was done and the

system closed before starting reaction.

The formation of NO was tested in complex

solutions with 0.029 g (1,95X10-4

mol) HIDA

and 0.0249 g (1,95X10-4

mol) oxovanadium(IV)

sulfate pentahydrate in 1 mL of deoxygenated

water and 0.013 g (1,95X10-4

mol) of nitrite

and 0.04395 g (1,04X10-1

mol) of

(if this complex reacts with

NO the color of the solution changes from

yellow to red) in 20 mL of deoxygenated water.

Both solutions were in different flasks

connected by a tube and vacuum was done

before starting reaction.

To confirm the formation of was tested

by the formation of a yellow insoluble salt,

. First it was added the

solutions of the complex with both proligands,

0.0079 g (4X10-5

mol) of HIDPA, 0.00596 g

(4X10-5

mol) of HIDA, separately and 0.00506

g (2X10-5

mol) vanadyl sulfate pentahydrated

to 0.0034 g (4X10-5

mol) of the oxidant

4

potassium nitrite. A few hours later it was

added cobalt nitrate,0.000378 g (1,3X10-5

mol).

The formation of or was tested

taking into account the pH variation resulting of

amavadin oxidation reactions (complex with

proligand HIDPA) at a solution of complex with

the concentration of 0.01 M by periodate,

iodate, persulfate, nitrite and cerium in

proportions metal: oxidant, 7: 1, 5: 1, 2: 1, 3: 1

and 1: 1, respectively.

Extraction with 2 mL of dichloromethane of

solutions of complexes using bromate, iodate

and periodate as oxidants were carried out.

The organic solution would be pink in the

presence of iodine[12] and yellow with

bromine[13].

2.5. Kinetics

The kinetics of the complex with HIDPA was

studied with, iodate, periodate, bromate,

chlorate, persulfate and nitrite by running 15

spectra at every 10 minutes interval for two

cases (eq. 1 and 2) with Perkin-Elmer

spectrophotometer in the range 350-850 nm.

Equation 1

Equation 2

The rate constants and order of reaction

were determined according table 2, and

considered[14,15],

Equation 3

3. Results and Discussion

In most cases, after addition of the oxidant,

solutions changed color to red due to oxidation

of vanadium complexes and later returned to

the blue color (with variants). These reactions

followed the pattern observed for [6] and

[8] meaning the return to the blue color

would lead to the water oxidation. In the cases

presented, there are some differences in the

cases already studied, and sometimes

reactions are very slow and some dependents

of light radiation.

Examining Table 3 with the ligand HIDPA by

various oxidants, it is clear that with the

exception of perchlorate (fig.3) they all reacted

with the complex. Although the latter being a

strong oxidizing agent, that is known to their

lack of reactivity[16].

Figure 3: Absorption spectra of [V(HIDPA)2 ]-2(0,01 M),

perchlorate(0,005M) covered(red) and transparent(blue).

Table 2: Absorbance vs time for the kinetics analyzed

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

Kinetic Concentration (M)[14] Absorbence[15]

0

1

2

5

With oxidants, iodate (fig.4), periodate (fig.5),

nitrite (fig.6) the solutions maintained under

light acquired a clear blue color probably

containing a vanadyl complex formed from

radical reactive species acting upon the

complex during the reduction of water.


iodate(0,005 M) covered(red) and transparent(blue).


periodate(0,005 M) covered(red) and transparent(blue).


nitrite(0,005 M) covered(red) and transparent(blue).

In a second study with nitrite with twice the

concentration and covered with yellow and red

filters, figure 7, there has been no different

behavior, having both solutions acquired the

initial color dark red and a final color light blue,

so it is possible that the concentration of nitrite

has no effect on the rate of the reaction.


nitrite(0,01 M) yellow and red filters.

In contrast, persulfate (fig.8) and

hypochlorite (fig.9) do not oxidize the complex

to red color (the five-vanadium oxidation state),

but only becomes purple in color, possibly a

reactive species. Additionally, the light does

not appear to have any effect on the rate of

reaction nor the final concentration of the

complex in solution.


persulfate(0,005 M) covered(red) and transparent(blue).


hypochlorite(0,005 M) covered(red) and transparent(blue).

For chlorate (fig.10) light has a different

effect. This was more prominent in the fact that

the solution alternatively changed color

between yellow and grey by the interaction

with light (in the presence and absence of light,

respectively).


chlorate covered(0,005 M)(red) and transparent(blue).

With tellurate (fig.10) and bismuthate (fig.11)

the results need additional studies, but a

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ (nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0,000

0,100

0,200

0,300

350 450 550 650 750 850

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 450 550 650 750 850

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

6

dilution of complex was observed after

variation of the color of the solution over time.


tellurate(0,01 M) (blue).


bismuthate (0,01M) (blue).

Table 3: Results for complex with HIDPA oxidation

Oxidants Covered Transparent

Bromate Initial color: Dark red Final Color: Violet blue In 6 days.

Initial color: Red Final Color: Violet blue In 2 hours.

Hipochlorite

It took 2 hours until purple. Initial color: Purple Final Color: Blue Violet In 6 days.

It took 2 hours to purple. Initial color: Purple Final Color: Blue Violet In 5 days or less.

Iodate

It took 1 day until dark red. Initial color: Dark red Final Color: Light blue In 21 days.

It took 1 day until dark red Initial color: Purple Final Color: Light blue In 3 days.

Perchlorate Reaction did not occur. Reaction did not occur.

Periodate Initial color: Dark red Final Color: Violet blue In 6 days.

Initial color: Dark red Final Color: Light blue In 1 day.

Nitrite Initial color: Dark red FinalColor: Violet blue In 7 days.

Initial color: Dark red Final Color: Light blue In 3 days.

Chlorate

It took 6 hours to purple and 1 day until red Initial color: Red Final Color: Light violet In 11 days.

It took 6 hours to purple and 1 day until red Initial color: Red Final Color: Light grey blue In 8 days.

Persulfate

It took 1 day until pink purple. Initial color:Pink Purple Final Color: Violet In 3 days.

It took 1 day until purple. Initial color: Purple Final Color: Violet In 3 days.

Telurate Initial color: Blue violet Final Color: Light violet In 6 days.

Bismuthate Initial color: Dark red Final Color: Blue In 7 days.

For the oxidation of the complex with ligand

HIDA, the results are summarized in table 4. It

can be seen that the reaction is much faster

compared to the previous ligand, (the

difference between them is just that the latter

has methyl groups in the carbon adjacent to

the carboxyl groups), and the radiation light

has no effect on the oxidation of the complex,

with decomposition of the vanadium complex

for all cases except hypochlorite,

persulfate(fig.13) and perchlorate (fig.14).

Figure 13: Absorption spectra of [V(HIDA)2 ]-2(0,01 M),

persulfate(0,005 M) covered(red) and transparent(blue).

Figure 14: Absorption spectra of [V(HIDA)2 ]-2(0,01M),

perchlorate(0,005 M) covered(red) and transparent(blue).

The absence of total regeneration of the

complex and a blue solution probably indicates

the formation of a vanadyl complex. Despite no

release was detected, the solution color

and the differences between the initial and

final spectrum, taking into account the results

obtained previously [6,8], the breakdown of the

complex should have occurred. The reaction

between the complex and oxidant nitrite

showed the higher decomposition of all cases

(fig.15).

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 450 550 650 750 850

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

7


nitrite(0,005 M) covered(red) and transparent(blue).

In respect of the complex with the ligand

HIDA, there were also no changes in the

behavior of the complex with yellow, red, green

and blue filters, taking solutions the initial color

clear purple and final color light blue at the end

of the same time, one day, Figure 16.


nitrite(0,01 M) yellow, red, green and blue filters.

The possible reduction of nitrite species are

the gases NO, , e , suggesting that

this oxidant originates , as final product,

given that the tests of NO and were

negative and the protons consumed on global

reaction in comparison with cerium(IV),

persulfate, iodate, and periodate. Nitrite is

used by some living organisms in the enzyme

nitrite redutase [17] in that the final product is

NO. If nitrite is used as a substrate by some

fungi Amanita, amavadina participates in its

reduction (metabolic pH is different from tested

in this work, which is very acidic, but in some

conditions protein medium can simulate low

pH) and leads to reduction, a new

contribution to the biogeochemical cycle of

nitrogen would be known. It also should be

noted that the production of by living

organisms is known [18], but involving other

substrate.

With bismuthate with the complex with HIDA

(fig.17) the results also need additional studies

to clarify the different behavior with the other

complex (compare with fig. 12), where the final

spectrum has higher absorbance at all

wavelength in contrast with the other cases.


bismutate (0,01M) (blue).

Table 4: Results for complex with HIDA oxidation

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 550 750

A

λ(nm)

0,000

0,100

0,200

0,300

350 550 750

A

λ(nm)

0 0,05

0,1 0,15

0,2 0,25

0,3 0,35

350 450 550 650 750 850

A

λ(nm)

Oxidants Covered Transparent

Bromate

Initial color: Purple Final Color: Light blue violet In 3 days.

Initial color: Purple Final Color: Light blue grey In 2 days.

Hipochlorite — Reaction did not occur

Iodate

Initial color: Light purple Final Color: Light blue grey In 3 days.

Initial color: Light purple Final Color: Light blue grey In 1 day.

Perchlorate Reaction did not occur.

Reaction did not occur.

Periodate

Initial color: Light orange Final Color: Light blue grey In 3 days.

Initial color: Light orange Final Color: Light blue grey In 1 day.

Nitrite

Initial color: Light purple Final Color: Blue In 1 day.

Initial color: Light purple Final Color: Almost colorless In 1 day.

Chlorate

Initial color: Blue violet Final Color: Light blue violet In 2 days.

Initial color: Blue violet Final Color: Light blue violet In 2 days.

Persulfate Reaction did not occur.

Reaction did not occur.

Telurate Reaction did not occur.

Bismuthate

Initial color: Blue Final Color: Light violet In 2 days.

8

The kinetics results are presented in table 5.

The rate constants showed reduction as the

slow reaction step. Considering all the data

iodate was the slowest reaction and especially

compared to periodate, although both are

reduced to iodine. The low value for the

reduction step using iodate is probably from

the strong character of its conjugate acid, iodic

acid, pKa = 0.75, being anionic form in initial

conditions (pH = 1.72). This species and the

complex are anions preventing the reaction

that occurs slowly. The reduction and the

oxidation step are both of second order

(fig.18), however persulfate showed zero order

for reduction (fig.19) and chlorate zero-order

for the oxidation (fig.20) and reduction, the only

example in cases studied and indicate a

concentration of the complex independent of

the rate reaction, and as such is the slowest

reaction after with iodate.

Figure 18: Kinetics for the reaction between complex with HIDPA and persulfate (0.005 M). R2 = 0.989.

Figure 19: Kinetics for the reaction between complex with HIDPA and persulfate (0.005 M) using the equation 6. R2 =

0.9997.

Figure 20: Kinetics for the reaction between complex with HIDPA and the chlorate (0.005 M). R2 = 0.979.

The isosbestic point was not well-defined

in the reduction of the complex for all oxidants

studied, example of chlorate (fig.21) unlike in

oxidation (fig.22), in which the point is between

675-680 nm. The absence of isosbestic point

may indicate that a secondary reaction occur

leading to the partial decomposition of the

complex.

Figure 21: Absorption spectrum in the visible region for the reaction with chlorate in 15 tests with 10-minute interval.

b = 1 cm.

Figure 22: Absorption spectrum in the visible region for the

reaction with chlorate in 15 tests with 10-minute interval.

b = 1 cm.

4

5

6

7

8

0 2000 4000 6000 8000 10000

1/A

t(s) 750 nm 775 nm 800 nm

0,3

0,31

0,32

0,33

0,34

0 5000 10000

A

t(s)

545 nm 554 nm 520 nm

0,15

0,17

0,19

0,21

0 2000 4000 6000

A

t(s)

750 nm 775 nm 800 nm

0,1

0,2

0,3

0,4

0,5

350 550 750

A

λ(nm)

0

0,2

0,4

0,6

0,8

1

1,2

350 550 750

A

λ(nm)

9

Table 5: Results for the kinetics

Oxidant Oxidation Reduction

Order k I.P. (nm) Order k I.P. (nm)

Iodate Second 4,87X10

-2

L.Mol-1

.s-1

680 Second

1,99X10-2

L.Mol

-1.s

-1

Indefined

Periodate Second 6,69X10

-2

L.Mol-1

.s-1

Indefined

Bromate Second 2,38X10

-2

L.Mol-1

.s-1

Indefined

Chlorate Zero 1,15X10

-6

M/s 675 Zero

1,35X10-6

M/s

Indefined

Persulfate Second 6,14X10

-2

L.Mol-1

.s-1

675 Zero

3,64X10-7

M/s

Indefined

Nitrite Second 1,97X10

-2

L.Mol-1

.s-1

Indefined

4. Conclusion

These results can inspire new

methodologies of synthesis, for example

halogenated organic compounds (which have

a high added value), as well as on amavadin

as mediator of water oxidation that has a

singular mechanism based on changing metal

oxidation in one unit and is a mononuclear

system[8]. Moreover, the behavior with respect

to chlorate with complex with HIDPA, may

allow specific sensors. It is noted that in our

research we found nothing comparable in the

literature for chlorate as oxidant.

Finally, these results reveal new information

about the biological role of amavadin despite

these experiments were carried out at different

pH than the biological ones. Hence its reaction

with nitrite (a biological metabolite) supports a

possible role to this vanadium

metallobiomolecule.

5. Acknowledgements

LD would like to thank family and friends.

6. Nomenclature

A-Absorbance

C-Concentration

Calc-Calculated

DMSO-Dimethylsulfoxide

DMF-Dimethylformamide

Exp-Experimental

HIDPA- N-hydroxyiminodipropionic acid

HIDA-N-hydroxyiminodiacetic acid

IV-Infra-Red

NHE-Normal Hydrogen Electrode

NMR-Nuclear Magnetic Ressonance

THF-Tetrahydrofuran

UV-Vis-Ultraviolet-Visible

7. References

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of vanadium, Springer, Vol. 89, 1999, 67-68.

[3] K. Wieghardt, V. Qulitzsch, B. Nuber and J.

Weiss, Angew. Chem. Int. Ed., 90, 1978, 385.

[4] E. Bayer, G. Koch and G. Anderegg

Angew. Chem. Int. Ed., 26, 1987, 545-546.

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Collison, S. Ertok, M. Heliweel and D. Garner,

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10

[6] J. Silva, J. R. Fraústo da Silva and A.

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[7] D. Rehder, Bioinorganic Vanadium

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[8] M, Domarus, M. Kutnesov, J. Marçalo, A. J.

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[9] A. J. L. Pombeiro, J. J. R. F. Silva, Y.

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[10] T. Hubregtse, Structural investigations of

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[11] E. Koch, H. Kneifel, and E. Bayer, Z.

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[12] G. Kerenskaya, I. U. Goldschleger, V. A.

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Phys. Chem. A , 2007, 111, 10969-10979.

[13] G. Kerenskaya, I. U. Goldschleger, V. A.

Apkarian and K. C. Janda, J. Phys. Chem A,

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[14] R. J. Silbey, R. A. Alberty and M. G.

Bawendi, Physical Chemistry, Wiley&Sons,

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[15] H. H. Perkampus, UV-Vis spectroscopy

and its applications, Springer, 1992, 3.

[16] G. M. Brown and B. Gu, The chemistry of

perchlorate in the environment, Springer,

2006, 17-47.

[17] R. R. Eady, S. Y. Antonyuk. and S. S. C.

Hasnain, C. Opinion in Chem. Bio., 2016, 31,

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