Prussian Blue Films Produced by Pentacyanidoferrate(II) and Their Application as Active Electrochemical Layers

Job/Unit: I42760 /KAP1 Date: 23-09-14 18:44:21 Pages: 9

FULL PAPER

DOI:10.1002/ejic.201402760

Prussian Blue Films Produced by Pentacyanidoferrate(II)and Their Application as Active Electrochemical Layers

Bruno Morandi Pires,[a] Sergio Augusto Venturinelli Jannuzzi,[a]

André Luiz Barboza Formiga,[a] and Juliano Alves Bonacin*[a]

Keywords: Electrochemistry / Redox chemistry / Sensors / UV/Vis spectroscopy / Iron

Iron complexes such as ferri/ferrocyanides are usually em-ployed as electrochemical mediators in portable devices forthe quick diagnosis of diseases. Stable complexes designedfor the electrochemical active layer may render devices withbetter performance and lifetimes than conventional devices.In this work we have synthesized and characterized spectro-scopically, electrochemically, and by DFT the complexNa4[Fe(CN)5(isn)] (isn = isonicotinate). This complex wasused as a single-source precursor of Prussian Blue (PB) in

Introduction

Point-of-care devices allow diagnosis by patients due totheir portability, fast response, and simplicity of handling.[1]

The most popular example of point-of-care devices is theglucose meter, which is a device created to measure the levelof glucose in blood. The advantage of using this approachis that the analysis can be performed by the owner, provid-ing early diagnosis. Such a device uses the specificity ofglucose oxidase to promote glucose oxidation to gluconicacid and uses ferrocyanide as a simple redox mediator.[2]

The change in role of complexes like ferri/ferrocyanideand their analogues from mediator to electroactive site re-quires the development of new types of complexes and al-ternative strategies to obtain modified electrodes. In ad-dition, there are several initiatives that use ferricyanide asthe electroactive material in the form of molecular com-plexes or structures like Prussian Blue (PB).[3–5]

In this context, pentacyanidoferrates ([Fe(CN)5L]n–) canbe highlighted due to their electronic structure, reactivity,and technological applications.[6–8] Such complexes have agreat affinity for N-heterocyclic compounds, amino acids,sulfoxides, thioethers, and thioamides.[9] Modification with

different ligands allows the modulation of their reactivitiesand makes possible their integration with other chemicalsystems. Recently, a pyridyl-bearing polymer was used to

[a] Institute of Chemistry, University of Campinas – UNICAMP,P. O. Box 6154, 13083-970, Campinas, SP, BrazilE-mail: [email protected]://www.lnanomol.iqm.unicamp.brSupporting information for this article is available on theWWW under http://dx.doi.org/10.1002/ejic.201402760.

Eur. J. Inorg. Chem. 0000, 0–0 © 0000 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim1

solution at pH 2.0 and 3.0, as confirmed by UV/Vis spec-troscopy. Potentiostatic deposition of PB onto glassy carbonelectrodes provided stable films at pH 3.0 and 5.0, with elec-trocatalytic activity for ascorbic acid oxidation with a linearresponse in the range 20–200 μM and an limit of detection of12 μM. The anodic peak potential of the modified electrodesvaried with solution pH with a minimum value of 0.350 Vversus NHE at pH 5.0.

bind pentacyanidoferrate in association with carbon nano-tubes to afford a superior electrocatalyst for cysteine detec-tion.[4]

Pentacyanidoferrate may also be used as a precursor toPrussian Blue, a well-known intervalence compound thathas been widely studied motivated by its magnetic, elec-trochromic, and electrochemical properties.[10–12] Further-more, this material has been employed in supercapacitors[13]

and energy storage devices.[14–16] PB is also recognized asan artificial peroxidase owing to its catalytic capacity in theconversion of H2O2 into H2O and O2. In addition, this elec-troactive compound has been used to modify electrodes toobtain electrochemical sensors and biosensors with low de-tection limits and selectivity. Electrodes modified by PBhave been used to quantify several analytes, such ashydrogen peroxide,[17] ascorbic acid,[11] and cysteine.[18]

Toma and co-workers[19] obtained PB films from penta-cyanidoferrates modified with nitrogen heterocycles but noelectrocatalytic properties of this type of film have been re-ported.

Vitamin C, or ascorbic acid (AA), is a well-known anti-oxidant and in addition exhibits anticarcinogenic activity aswell as acting against degenerative diseases and playing arole in cell division and genetic expression.[20,21] The quanti-fication of this analyte is of great interest in the pharmaceu-tical and food industries in which it is used as an antioxi-dant and nutraceutical ingredient. Furthermore, AA is aninterferent in glucose analysis performed by glucose me-ters.[22] The electrochemical oxidation of AA is well de-scribed in the literature[20,21,23] and can be improved byusing PB-modified electrodes by reducing the working po-tential.


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In this paper we report the synthesis and spectroscopicand electrochemical characterization of the pentacyanido-ferrate(II) complex Na4[Fe(CN)5(isn)] (isn = pyridine-4-carboxylate or isonicotinate) to determine whether the na-ture of the N-heterocyclic group and its charge influence1) the chemical stability of the pentacyanidoferrate(II) com-plex at various pH, 2) the ability to form PB films, and3) the electrochemical response of PB to AA.

Results and Discussion

Syntheses involving isonicotinic acid (isnH) and com-plexes with one labile site are a challenge. First, isonicotinicacid has basic and acid sites, so the formation of zwitterionspecies needs to be considered. Previous neutralization ofisnH to its sodium salt provides high yields of Na4[Fe(CN)5-(isn)] and avoids the formation of PB in this synthesis, asshall be discussed below.

Another question is how isn is coordinated in complex,through the nitrogen or the carboxylic group. To solve thisquestion, vibrational spectroscopy (FTIR) was used be-cause the vibrational modes of the pyridine ring are sensi-tive to the coordination of Fe(CN)5

3– group.[24] Assign-ments were performed based on computational tools andare presented in detail in the Supporting Information. Onthe basis of the results, it is evident that isn is coordinatedthrough the pyridyl group and the anionic form of the carb-oxylic group.

Another important point to consider is whether thecharge of the ligand (isonicotinate) changes the kinetics ofthe dissociation of isn. One hypothesis would be that thenegative charge can increase the dissociation rate due tothe electrostatic repulsion between [Fe(CN)5]3– and isn. Thekinetics of the reactions of the complexes can be investi-gated by monitoring their UV/Vis spectra, however, it isnecessary to understand the nature of the electronic transi-tion and its changes.

UV/Vis Spectroscopy

Figure 1 shows the electronic absorption spectrum of theprecursor complex Na3[Fe(CN)5NH3]. A weak band is ob-served at 400 nm (ε = 4.50� 102 m–1 cm–1), assigned to a d–d transition.[25,26] After reaction with isn, a band at 417 nmwith ε = 4.08� 103 m–1 cm–1 is observed. An absorption inthe 400–500 nm range with a molar absorption coefficientin the order of 103 is typical of a metal-to-ligand charge-transfer band (MLCT) for N-heterocyclic compounds coor-dinated to pentacyanidoferrate(II).[6] The data for theMLCT transition of the Na4[Fe(CN)5(isn)] complex andother pentacyanidoferrates are presented in Table S3 in theSupporting Information.


Figure 1. Molar absorption coefficient as a function of wavelengthfor pentacyanidoferrate(II) complexes in aqueous solution.

With the purpose of confirming the nature of this elec-tronic transition, theoretical studies were performed byusing time-dependent density functional theory (TDDFT).The calculations revealed an intense transition (oscillatorstrength: 0.139) between states, the major constituents ofwhich are the Kohn–Sham (KS) orbitals 78 and 80 (see Fig-ure S3 in the Supporting Information), which correspondto the HOMO–1 and LUMO. The composition of these KSorbitals in terms of atomic basis functions is presented inTable 1. One can see that the fundamental state, HOMO–1, is predominantly metal (81% Fe), whereas the LUMO ispredominantly ligand (81% isn).

Table 1. Compositions of the molecular orbitals involved in theMLCT calculated by TDDFT at the PBE0/def2-TZVP level.

KS Orbital Atom Basis function Composition [%]

78 Fe dyz 81.1(HOMO–1) Nax py 5.2

Neq pz 2.0

80 Fe dyz 1.6(LUMO) Nligand py 16.1

C(2) and C(6) py 4.5C(3) and C(5) py 8.4C(4) py 29.5C(7) py 3.4O py 3.3

The TDDFT calculations showed the MLCT transitionfor the complex at 394 nm, in good agreement with the ex-perimental value of 425 nm recorded for a freshly preparedsolution in dimethyl sulfoxide (see Figure S4). The devia-tion between the experimental and calculated values is com-patible with calculations performed at a high level oftheory.[27]

Substitution Kinetics in the Presence of Dimethyl Sulfoxide

Nitrogen heterocycles coordinated to pentacyanido-ferrate(II) are known to undergo substitution by dmso.[9]

The kinetics of isn substitution was investigated in order to



make comparison with literature data[28] and also to investi-gate the effect of a negative charge on the leaving ligand.

The replacement of isn by dmso is accompanied by adecrease in the absorbance of the MLCT band of the com-plex containing isonicotinate (λ = 417 nm), the dimethylsulfoxide complex not showing any transitions in the visiblerange.[29] Variations in the absorbance as a function of timeare shown in Figure 2.

Figure 2. Kinetic plots for the substitution of the pentacyanidofer-rate(II) complex in aqueous dmso solutions. Normalized ab-sorbances are shown. μ = 1.0 m (NaCl).

The mechanism of the ligand substitution in the penta-cyanidoferrate(II) complexes has been extensivelystudied[30,31] and can be rationalized by the dissociativemechanism described by Equations (1), (2) and (3). In thecomplex [Fe(CN)5L]n–, the rate constant for substitution ofligand L� is independent of L�, although the substitutionrate varies with different ligands.

(1)

(2)

(3)

It is assumed that when both L and L� are strongernucleophiles than water, the aqua species [Equation (2)] ispresent in insignificant concentrations and has no effect onthe mechanism. In our study, L = isn and L� = dmso. Fromthis information it is possible to obtain the observed rateconstant (kobs) by using Equation (4).

kobs =k–isnkdmso[dmso]

kisn[isn] + kdmso[dmso](4)

By using excess dmso to obtain pseudo-first-order condi-tions, then kdmso[dmso] �� kisn[isn] and the observed rateconstant will be equal to the rate constant for the dissoci-


ation of the isn ligand (k–isn). Considering the linear rangeof absorbance of the complex in the range of 0 to 600 s inFigure 2, different values of kobs as a function of tempera-ture were obtained. The activation energy of the processwas calculated by using the Eyring equation, the plot ofwhich is shown in Figure 3.

Figure 3. Plot of ln kobs/T vs. 1/T for the substitution of isn in[Fe(CN)5L]n– by dmso.

Table 2 presents the kinetic data obtained in this studyand for other related complexes reported in the literature.The relatively low value of the observed rate constant forthe dissociation of isn is in agreement with the value re-ported by Hoddenbagh and Macartney.[32] The enthalpy ofdissociation was calculated to be 22.9 kcalmol–1 and lies inthe same range as other pentacyanidoferrate(II) complexes.The entropy of dissociation was determined to be3 calmol–1 K–1 and the positive value is compatible with adissociative mechanism. Therefore, in spite of the fact thatisn is coordinated as a negative ligand, the kinetics of sub-stitution are similar to other N-heterocyclic ligands.

Table 2. Dissociation constants and activation parameters forpentacyanidoferrate(II) complexes.

Ligand Charge kd[a] ΔH ΔS Ref.

[104 s–1] [kcalmol–1] [calmol–1 K–1]

Isonicotinate –1 4.0 22.9 3 this workIsonicotinamide 0 7.3 26.0 14 [9]

dmso 0 0.75 26.5 11 [9,29]

Pyridine 0 11 24.8 11 [9]

Pyrazine 0 4.2 26.4 14 [9]

Methylpyrazinium +1 2.8 27.5 18 [9]

[a] At 25 °C.

Electrochemical Properties of the Complex [Fe(CN)5(isn)]4–

Pentacyanidoferrates(II) exhibit well-defined electro-chemical properties that are influenced by the nature of theligands. Cyclic voltammograms obtained for the penta-cyanidoferrate(II) complexes are presented in Figure 4.



Figure 4. Cyclic voltammograms of 5.0 �10–3 m aqueous solutionsof the pentacyanidoferrate complexes.

The voltammograms show a redox process at E°� =0.50 V versus NHE, which has been assigned to the[Fe(CN)5(isn)]3–/4– couple. All the potentials measured areconsistent with the values reported in the literature forpentacyanidoferrate(II) complexes containing N-hetero-cyclic ligands.[9,28] The difference between the cathodic andanodic peaks indicates that the redox process exhibitsquasi-reversible behavior (ΔEp = 97 mV at a scan rate of

Figure 5. UV/Vis spectra of the Na4[Fe(CN)5(isn)] complex in Britton–Robinson buffer solutions at pH 2.0 (A) and 3.0 (C) and variationof the MLCT and intervalence bands with time at pH 2.0 (B) and 3.0 (D).


25 mVs–1; see Figure S5A in the Supporting Information).The ratio between the cathodic and anodic current peaks

is close to unity for all the scan rates used, and a linearvariation of current peak with ν1/2 is observed (see Fig-ure S5B). This behavior is compatible with a diffusion-con-trolled process.[33] For comparison purposes, the cyclic vol-tammetry of the complex [Fe(CN)5NH3]3– was obtainedshowing E°� = 0.35 V versus NHE. The shift to more posi-tive potentials reflects the back-bonding ability of the N-heterocyclic ligand when exchanging the ligands NH3 andisn.[34]

Electrochemical Study at Different pH

The cyclic voltammograms obtained in Britton–Robin-son buffer solutions at different pH are presented in Fig-ure S6 in the Supporting Information. The redox couple of[Fe(CN)5(isn)]3–/4– shows quasi-reversible behavior at pHvalues lower than 7.0. At pH 2.0 it is possible to observe aredox couple at around 0.4 V at higher scan rates, whichwas assigned to the redox process of PB. As the pH in-creases above pH 7.0 the process becomes more irreversible,with higher values for the anodic and cathodic peak poten-tials. Little information is available in the literature concern-ing the redox process of cyanidoferrates at high pH. It hasbeen suggested that non-characterizable species form insolution with different diffusion coefficients that affect theshapes of the voltammograms. It is reasonable to consider



the formation of polynuclear complexes and the precipi-tation of FeOOH.[35] Another interesting feature is the vari-ation of the anodic peak potential with pH, which reachesa minimum in pH 5.0 (see Figure S7).

Study of the Influence of pH on [Fe(CN)5(isn)]4–

The UV/Vis spectra of solutions of [Fe(CN)5(isn)]4– atdifferent pH allowed the dissociation of the complex to bemonitored due to the decrease of the MLCT band at417 nm. This decrease is a result of ligand exchange be-tween isn and water to give the aquapentacyanidoferrate(II)complex (see Equation 5 in Scheme S1 in the SupportingInformation).

In comparison with neutral solutions, at pH 2.0 ahypsochromic shift of the MLCT band is observed, withthe maximum absorption wavelength changing to 405 nm(Figure 5). This can be ascribed to the protonation of thecyanide ligands, which stabilizes the dπ orbitals throughback-bonding donation.[36] At pH 3.0 a bathochromic shiftoccurs with the maximum absorbance wavelength appear-ing at 431 nm. After 1 h under both pH conditions the ab-sorbance maximum is shifted to 400 nm and a band ap-pears at around 730 nm (Figure 5), which has been assignedto an intervalence transfer band typical of PB.[37,38]

At pH � 7 the aquapentacyanidoferrate(II) decomposesinto FeII [see Scheme S1 in the Supporting Information,Equation (S2)],[39] which, in aqueous solution, is oxidizedto FeIII [see Scheme S1, Equation (S3)]. In solution, theseions can interact with the cyanidoferrates to form PB (seeSupporting Information, Equation 4 in Scheme S1). Thisstructure seems to contain the PB structure together withthe unreacted isn complex due to the stabilization of theMLCT band at the end of the experiment. Prussian Bluetype structures containing N-heterocyclic ligands have pre-viously been reported.[19] Another possible mechanism (seeSupporting Information, Equation 4 in Scheme S1) is thesubstitution of isn by cyanide (see Supporting Information,Equation 5 in Scheme S1), which would lead to PB without

Figure 6. A) Voltammograms of Prussian Blue deposited on CV electrodes in 0.1 m HCl and 0.1 m KCl at different scan rates. B) Relation-ship between peak currents and scan rates.


the isn ligand (see Supporting Information, Equation 6 inScheme S1).

At pH � 7, considerable dissociation occurs comparedwith under neutral solutions (see Figure S8). It is knownthat at high pH, there is strong competition between CN–

and OH– for iron, and mixed complexes can be formed.Also there is a possibility of the precipitation of FeOOH athigher concentrations. Even the stable complex [Fe(CN)6]3–

can undergo substitution reactions at higher pH values.[35]

Electrode Modification with PB Films Produced fromPentacyanidoferrate

The voltammograms of PB films in 0.1 m HCl and 0.1 m

KCl solution are presented in Figure 6. After film deposi-tion, the film is composed mainly of the insoluble form ofthis compound (Fe4

III[Fe(CN)6]3), but after several cyclesof the redox pair, partial conversion to the soluble form(KFeIII[Fe(CN)6]) occurs due to the insertion of potassiumions into the structure.[40] The redox pair with E°� ≈ 0.47 Vcan be rationalized by Equation (5) (insoluble form) andEquation (6) (soluble form)[41] and involves the conversionbetween Prussian Blue and its reduced form, known asPrussian White (Everitt’s Salt).

Fe4III[FeII(CN)6]3 + 4K+ + 4e– �K4Fe4

II[FeII(CN)6]3 (5)

KFeIII[FeII(CN)6] + K+ + e– �K2FeII[FeII(CN)6] (6)

To evaluate the stability of the films for application inelectrochemical sensors, the films were cycled in Britton–Robinson buffer solutions at different pH. The results wereas expected for a PB film, with increasing stability withdecreasing pH (see Figure S9 in the Supporting Infor-mation). At pH 3.0 and 5.0 the film shows well-defined andreversible processes. However, at pH 7.0 the film has lowstability after 50 cycles, with a considerable decrease in thepeak currents. The profiles of the PB films obtained bycyclic voltammetry at different pH are shown in Figure S10.The surface structures and thicknesses of the PB films mayalso have an effect on their stability.



Figure 7. Electrocatalytical oxidation of ascorbic acid with PB films by cyclic voltammetry and curves for the current response as afunction of concentration at pH 3.0 (A and B) and 5.0 (C and D).

Electrocatalytic Oxidation of Ascorbic Acid (AA)

The results obtained in the stability studies were used toselect the most suitable pH conditions for the electrocata-lytic studies of ascorbic acid oxidation using glassy carbon(GC)/PB electrodes. Figure 7 shows the results obtained forthe studies performed at pH 3.0 and 5.0. A linear incrementof the anodic peak current with increasing concentration ofAA can be observed under both conditions in the range of20 to 200 μm. An interesting feature of this material is thereduction of the anodic peak potential with increasing pH,which occurs at around 0.45 and 0.35 V versus NHE insolutions at pH 3.0 and 5.0, respectively (Figure 7). Thisfeature is desirable for applications of this material in sen-sors, in which low working potentials are required to avoidthe oxidation of interfering species. Thus, pH 5.0 would bethe optimal working condition for this system.

PB films obtained from [Fe(CN)5(isn)]4– showed loweroxidation potentials and lower limits of detection thanother reported electrodes used in AA analysis. This behav-ior can be explained by the presence of the isn ligand inthe PB structure. Our data are presented in Table S4 in theSupporting Information along with data obtained for othermodified electrodes using voltammetric techniques for com-parison. Pournaghi-Azar[42] observed a peak potential of0.44 V versus NHE in the CV analysis of AA with a linearresponse in the range of 5� 10–5 to 1.5�10–3 m and with a


detection limit of 20 μm. In our case, the peak potential was0.35 V versus NHE with a linear response observed overthe AA concentration range of 20–200 μm, and the limit ofdetection was 12 μm. Therefore the use of the [Fe(CN)5-(isn)]4– complex to obtain PB films is shown to be an inter-esting approach to producing stable sensors with low work-ing potentials.

Conclusions

Neutralization of isonicotinic acid prior to reaction withpentacyanidoferrate(II) is necessary to produce the complexNa4[Fe(CN)5(isn)] in high yields. The N-heterocyclic ligand(isn) is coordinated in this complex through the pyridinering and the deprotonated carboxy group. These findingswere supported by DFT calculations.

Kinetics studies of ligand substitution provided acti-vation parameters consistent with the results for analogouscomplexes, thereby confirming a dissociative mechanism.Despite the negative charge of the N-heterocyclic ligand, itsrate constant for dissociation is of the same magnitude asother reported ligands.

[Fe(CN)5(isn)]4– shows the greatest stability betweenpH 4 and 7; PB formation was observed below pH 4,whereas iron hydroxide may be formed above pH 7.



The complex was used to obtain PB-modified electrodesby using a simple potentiostatic method. The PB film pro-duced from this complex has potential in electrochemicalanalysis due to its low working potential (0.35 V vs. NHE),a linear response in the range studied (20–200 μm), and alow limit of detection (12 μm). The results suggest that PBfilms derived from pentacyanidoferrates are suitable for useas an electroactive layer in molecular sensing, for example,of ascorbic acid.

Experimental SectionSynthesis of Pentacyanidoferrates

Na3[Fe(CN)5NH3]·3H2O: The complex was obtained by adding28% ammonium hydroxide (NH4OH m/v, Sigma–Aldrich; 40 mL)to an 250 mL Erlenmeyer flask containing sodium nitroprusside(Na2[Fe(CN)5NO]·2H2O, 99%, Acros Organics; 6.0 g, 20 mmol).The flask was stirred until the complete solubilization of sodiumnitroprusside. Some bubbles appeared in the flask during this pro-cedure. Then the flask was covered with aluminium foil and thetop of the flask covered with cotton to allow the gas to exit. After3 h in the dark, the solution was a dark-yellow color. Sodium iod-ide (NaI 98% Merck; 6.0 g, 40 mmol) was added and a yellow pre-cipitate appeared. Ethanol (C2H6O 99.9%, Merck; 100 mL) wasslowly added to ensure complete precipitation. The solid was fil-tered, washed with ethanol, and dried on a vacuum line until con-stant weight, yield 93%. C5H9FeN6Na3O3 (325.98): calcd. C 18.42,H 2.78, N 25.78; found C 18.50, H 2.07, N 26.06.

Na4[Fe(CN)5(isn)]·5H2O (isn = pyridine-4-carboxylate): Isonicotinicacid (C6H5NO2, Sigma–Aldrich; 0.12 g, 0.97 mmol) was neutral-ized by adding an aqueous solution of sodium hydroxide (NaOH,Synth; 0.0327 g, 0.82 mmol, 20 mL) to produce sodium isonicotin-ate. The solvent was removed in a rotary evaporator at 40 °C andthen the salt was solubilized again in distilled water (20 mL). Asample of this solution (10 mL) was slowly added to an aqueoussolution of Na3[Fe(CN)5NH3]·3H2O (0.159 g, 0.5 mmol, 10 mL)whilst stirring. An orange color was observed in the solution imme-diately after the addition of the first droplets. The flask was stirredfor 30 min and the solvent removed in a rotary evaporator at 25 °Cover 1 h. The orange precipitate obtained was dried on a vacuumline until constant weight, yield 82%. C11H14FeN6Na4O7 (490.07):calcd. C 26.96, H 2.88, N 17.15; found C 26.77, H 2.87, N 17.92.

Characterization of the Pentacyanidoferrates

IR Spectroscopy: IR spectra of the iron complexes and isonicotinicacid were recorded as KBr pellets in an MB100 Bomem Spectrome-ter at a resolution of 2 cm–1 in the range 4000–400 cm–1.

Electronic Spectroscopy: Electronic spectra in the range 190–1100 nm were acquired by using a 1 cm quartz cuvette in a diodearray HP8453 UV/Visible absorption spectrophotometer equippedwith a HP89090A Peltier.

Molecular Modeling: The reader is referred to the Supporting In-formation for details.

Influence of pH on the Stability of [Fe(CN)5(isn)]4–: The stability ofthe complex [Fe(CN)5(isn)]4– at different pH was investigated byrecording the UV/Vis spectra of 1.0 �10–4 m solutions of the com-plex in Britton–Robinson buffer solutions[43] every 30 s for 1 h at25 °C.

Electrochemical Properties: Cyclic voltammograms were obtainedwith an Autolab EcoChemie PGSTAT20 instrument by using 0.1 m


potassium chloride (KCl, Synth) as supporting electrolyte in allmeasurements. A glassy carbon electrode was used as the workingelectrode, an Ag/AgCl/KCl saturated electrode as the reference elec-trode, and platinum electrode as the auxiliary electrode. For eachmeasurement, 5.0 �10–3 m solutions of the complexes potassiumferrocyanide, sodium aminopentacyanidoferrate(II), andNa4[Fe(CN)5(isn)]·5H2O were used. All the solutions were deaer-ated by bubbling nitrogen before the measurements. The cathodicand anodic potentials chosen were –0.2 and 1.0 V versus Ag/AgCl,respectively, and the scan rates were 25, 50, 100, and 200 mV s–1.

The effect of pH on the electrochemical behavior was investigatedby preparing 2.0 �10–3 m solutions of the complex in Britton–Rob-inson buffer solutions at pHs ranging from 2 to 12 with 0.1 m KClas the supporting electrolyte. The potential range used was from–0.1 to 0.8 V and the scan rates were 5, 10, 25, 50, 100, and200 mVs–1.

Substitution Kinetics of the N-Heterocyclic Ligand in the Presenceof dmso: The substitution kinetics of the pentacyanidoferrate(II)containing the N-heterocyclic ligand was studied by UV/Vis spec-troscopy of aqueous solutions of the complex containing excessdimethyl sulfoxide (Synth). This ligand was chosen on the basis ofthe higher stability of the complex obtained.[9,29] The absorbanceof aqueous solutions of the complex [Fe(CN)5(isn)]4– containing0.5 m dimethyl sulfoxide was measured over 1 h with stirring undera controlled temperature. The ionic strength was controlled byusing 1.0 m sodium chloride.

Electrode Modification with Cyanidoferrate Films: Before modifica-tion, glassy carbon electrodes were polished in an alumina slurry,rinsed with deionized water, washed ultrasonically, and finallycleaned with deionized water. Films of the complex [Fe(CN)5-(isn)]4– were deposited onto the surfaces of glassy carbon electrodesby electrochemical deposition of a PB analogue. A solution con-taining 1.0 �10–3 m [Fe(CN)5(isn)]4–, 1.0�10–2 m KCl, and1.0�10–2 m HCl was prepared in an electrochemical cell and a po-tassium peroxydisulfate solution (Merck) was added to obtain afinal concentration of 1.0 �10–2 m of this compound. The solutionwas stirred for 10 min and then ferric chloride was added to obtaina final concentration of 1.0 �10–3 m. Films were deposited on aglassy carbon electrode by applying a potential equal to 0.61 V for600 s. The stability of the films was studied by cycling the modifiedelectrodes in Britton–Robinson buffer solutions at pH 3.0, 5.0, and7.0 containing 0.1 m KCl in the range 0.11–0.61 V.

Electrocatalytic Oxidation of Ascorbic Acid: The electrocatalyticproperties of the GC/PB electrode were tested in the oxidation ofAA by cyclic voltammetry. Stock solutions of this analyte wereadded to an electrochemical cell containing Britton–Robinsonbuffer solution containing 0.1 m KCl to study the response of themodified electrode with increasing AA concentration.

Supporting Information (see footnote on the first page of this arti-cle): Molecular modeling methodology and the results, descriptionof vibrational spectroscopy (IR), comparison between theoreticaland experimental spectroscopic data, cyclic voltammograms at dif-ferent pHs, plausible mechanisms for the formation of PB, UV/Visspectra of [Fe(CN)5(isn)]4– in pH range of 4.0 to 12.0, stability stud-ies of the PB films, effect of pH on GC/PB electrodes.

Acknowledgments

The authors acknowledge the Conselho Nacional de Desenvolvi-mento Científico e Tecnológico (CNPq), the Fundo de Apoio aoEnsino, à Pesquisa e à Extensão - Universidade Estadual de Cam-



pinas (FAEPEX-UNICAMP) (1321/12), and Fundação de Am-paro a Pesquisa do Estado de São Paulo (FAPESP) for financialsupport and the National Center for High Performance Computingin São Paulo (CENAPAD, Project 557/2013) for computational re-sources.

[1] C. H. Ahn, J.-W. Choi, G. Beaucage, J. Nevin, J.-B. Lee, A.Puntambekar, R. J. Y. Lee, Proc. IEEE 2004, 92, 154–173.

[2] J. Wang, Chem. Rev. 2008, 108, 814–825.[3] A. Zloczewska, A. Celebanska, K. Szot, D. Tomaszewska, M.

Opallo, M. Jönsson-Niedziolka, Biosens. Bioelectron. 2014, 54,455–61.

[4] C. C. Corrêa, S. A. V. Jannuzzi, M. Santhiago, R. A. Timm,A. L. B. Formiga, L. T. Kubota, Electrochim. Acta 2013, 113,332–339.

[5] E. Nossol, A. J. Gorgatti Zarbin, J. Mater. Chem. 2012, 22,1824–1833.

[6] L. M. Baraldo, P. Forlano, A. R. Parise, L. D. Slep, J. A. Olabe,Coord. Chem. Rev. 2001, 221, 881–921.

[7] H. E. Toma, J. M. Malin, Inorg. Chem. 1973, 12, 1039–1045.[8] H. E. Toma, J. M. Malin, Inorg. Chem. 1973, 12, 2080–2083.[9] S. H. Toma, J. A. Bonacin, K. Araki, H. E. Toma, Eur. J. Inorg.

Chem. 2007, 3356–3364.[10] M. Ohba, H. Okawa, Coord. Chem. Rev. 2000, 198, 313–328.[11] D. M. DeLongchamp, P. T. Hammond, Adv. Funct. Mater.

2004, 14, 224–232.[12] A. A. Karyakin, Electroanalysis 2001, 13, 813–819.[13] J. Chen, K. Huang, S. Liu, X. Hu, J. Power Sources 2009, 186,

565–569.[14] C. D. Wessells, S. V. Peddada, R. A. Huggins, Y. Cui, Nano

Lett. 2011, 11, 5421–5425.[15] S. Yagi, M. Fukuda, R. Makiura, T. Ichitsubo, E. Matsubara,

J. Mater. Chem. A 2014, 2, 8041–8047.[16] P. Nie, L. Shen, H. Luo, B. Ding, G. Xu, J. Wang, X. Zhang,

J. Mater. Chem. A 2014, 2, 5852–5857.[17] E. Jin, X. Lu, L. Cui, D. Chao, C. Wang, Electrochim. Acta

2010, 55, 7230–7234.[18] S. S. L. Castro, V. R. Balbo, P. J. S. Barbeira, N. R. Stradiotto,

Talanta 2001, 55, 249–254.[19] F. M. Matsumoto, M. L. A. Temperini, H. E. Toma, Electro-

chim. Acta 1994, 39, 385–391.


[20] B. N. Ames, M. K. Shigenaga, T. M. Hagen, Proc. Natl. Acad.Sci. USA 1993, 90, 7915–7922.

[21] L. Yang, D. Liu, J. Huang, T. You, Sens. Actuators B 2014,193, 166–172.

[22] F. Martinello, E. Luiz da Silva, Clin. Chim. Acta 2006, 373,108–116.

[23] I. G. Casella, M. R. Guascito, Electroanalysis 1997, 9, 1381–1386.

[24] S. A. V. Jannuzzi, B. Martins, M. I. Felisberti, A. L. B. For-miga, J. Phys. Chem. B 2012, 116, 14933–14942.

[25] E. J. Baran, A. Muller, Z. Anorg. Allg. Chem. 1969, 368, 144–154.

[26] H. B. Gray, C. J. Ballhausen, J. Chem. Phys. 1962, 36, 1151–1153.

[27] A. L. B. Formiga, S. Vancoillie, K. Pierloot, Inorg. Chem. 2013,52, 10653–10663.

[28] F. S. Nunes, L. S. Bonifácio, K. Araki, H. E. Toma, Inorg.Chem. 2006, 45, 94–101.

[29] H. Toma, J. Malin, E. Giesbrecht, Inorg. Chem. 1973, 12, 2084–2089.

[30] S. Alshehri, J. Burgess, R. van Eldik, C. D. Hubbard, Inorg.Chim. Acta 1995, 240, 305–311.

[31] K. Bal Reddy, R. van Eldik, Inorg. Chem. 1991, 30, 596–598.[32] J. Hoddenbagh, D. Macartney, Inorg. Chem. 1986, 25, 2099–

2101.[33] A. Bard, L. Faulkner, Electrochemical Methods, Wiley, New

York, 2012.[34] H. E. Toma, C. Creutzl, Inorg. Chem. 1977, 16, 545–550.[35] W. N. Perera, G. Hefter, Inorg. Chem. 2003, 42, 1230–1234.[36] H. E. Toma, J. Inorg. Nucl. Chem. 1976, 38, 431–434.[37] D. Ellis, M. Eckhoff, V. D. Neff, J. Phys. Chem. 1981, 96, 1225–

1231.[38] K. Itaya, T. Ataka, S. Toshimaf, J. Am. Chem. Soc. 1982, 4767–

4772.[39] S. Asperger, I. Murati, D. Pavlovic, J. Chem. Soc. 1953, 730–

736.[40] Z. Wang, H. Yang, B. Gao, Y. Tong, X. Zhang, L. Su, Analyst

2014, 139, 1127–1133.[41] F. Ricci, G. Palleschi, Biosens. Bioelectron. 2005, 21, 389–407.[42] M. Pournaghi-Azar, Talanta 1995, 42, 1839–1848.[43] H. T. S. Britton, R. A. Robinson, J. Chem. Soc. 1931, 1456–

1462.Received: August 6, 2014

Published Online: �



Prussian Blue Sensor

B. Morandi Pires,S. A. Venturinelli Jannuzzi,A. L. Barboza Formiga,J. Alves Bonacin* .............................. 1–9

Prussian Blue Films Produced by Penta-cyanidoferrate(II) and Their Application asActive Electrochemical Layers

Prussian Blue is widely used in electro- tivity, selectivity, and specificity of these de-Keywords: Electrochemistry / Redox chem-chemical sensors due to its ability to cata- vices can be improved by using Prussianistry / Sensors / UV/Vis spectroscopy / Ironlyze H2O2 reduction. Glucose meters are Blue produced by pentacyanidoferrate(II)common point-of-care devices that employ as an electroactive species rather than a re-ferricyanide as a redox mediator. The sensi- dox mediator.


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Prussian Blue Films Produced by Pentacyanidoferrate(II) and Their Application as Active Electrochemical Layers