Preparation of Monodisperse Cobalt(II) Hexacyanoferrate(III

Article pubs.acs.org/JPCC

Preparation of Monodisperse Cobalt(II) Hexacyanoferrate(III) Nanoparticles Using Cobalt Ions Released from a Citrate Complex Fumiyuki Shiba,*,† Ryosuke Fujishiro,† Takashi Kojima,‡ and Yusuke Okawa† †

Graduate School of Advanced Integration Science, and ‡Graduate School of Engineering, Chiba University 1-33 Yayoicho, Inageku, Chiba 263-8522, Japan S Supporting Information *

ABSTRACT: A procedure to obtain monodisperse nanoparticles of cobalt(II) hexacyanoferrate(III) was investigated. To prevent a reaction with [Fe(CN)6]3−, Co2+ was stored in the form of a citrate complex at pH 7.8. By lowering the pH to 3.4 by the addition of acid, part of the Co2+ was released and reacted with [Fe(CN)6]3− to form nuclei, followed by their growth via mass-transfer of Co2+ from the citrate complex to the solid. Here, gelatin was employed as a protective colloid. The particles obtained were cubic and fairly monodisperse. The mean particle size was typically 160 nm and varied from 30 to 450 nm, depending on the reactant contents. The formation mechanism associated with equilibrium shifts was supported by a calculated estimation of free Co2+ on the basis of the equilibrium relationships involved in the preparation system.

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3Co2 + + 2[Fe(CN)6 ]3 − → Co3[Fe(CN)6 ]2

INTRODUCTION Cobalt(II) hexacyanoferrate(III) (Co3[Fe(CN)6]2 or CoHCF) is one of the Prussian blue analogues (PBAs). The original Prussian blue, Fe4[Fe(CN)6]3 (FeHCF) is a pigment that contains both divalent and trivalent irons, and its mixed valence state is responsible for its blue color via intervalence charge transfer absorption.1 Because of the reversible and electrochemically applicable redox property associated with cation insertion for charge compensation, FeHCF has been studied for use in a number of applications such as electrochromic,2 sensor,3 and Li+ storage4 materials. On PBAs, including CoHCF, a part of, or sometimes the entire iron in FeHCF is replaced by other transit metals, and CoHCF has also been studied for sensor5 and electrochromic6 applications, in addition to being of interest owing to its magnetic properties.7 Nanoparticles are an interesting form of the materials used for these objectives because of their large specific surface area, applicability to printing processes, etc. The properties of particulate materials generally depend on their size. Thus, monodisperse nanoparticles that have a narrow distribution of size and properties are ideal for use. However, in the literature, there are a limited number of reports that focus on the formation of monodipserse CoHCF nanoparticles in media such as water-in-oil emulsion systems,8,9 despite there being many studies on CoHCF synthesis itself.10−16 The preparation of monodisperse particles generally requires the clear separation of the nucleation period from the growth period that follows, and this is done by controlling the supersaturation ratio, as suggested by the LaMer mechanism.17 To achieve this, the issue of the storage and supply of the reaction resources is of great practical importance.18 The formation reaction of PBA, for instance, © 2012 American Chemical Society

(1)

for CoHCF, occurs instantaneously, forming polydisperse particles when reactant solutions such as Co(NO3)2 and K3[Fe(CN)6] are simply mixed. Thus, a way to preserve the reactants should be introduced to satisfy the LaMer mechanism for this material. Recently, Shiba19 reported the synthesis of monodisperse FeHCF nanoparticles via a reduction reaction by citric acid, where one of the reactants, Fe2+, was stored in an oxidized form (Fe3+) in the reacting solution and was supplied by a moderate reduction from Fe3+ to Fe2+ to react with [Fe(CN)6]3−. In the present study, we propose an alternative procedure for the preparation of monodisperse PBA nanoparticles using a citrate complex of metal ions as a reservoir in the CoHCF system. Dissociated forms of citrate ions can form complexes with Co2+,20 and they are unreactive with [Fe(CN)6]3−. Since the capacity of Co2+ storage decreases at a lower pH because there are smaller association constants for more protonated forms of citrate ions, a controlled Co2+ release that gives monodisperse CoHCF nanoparticles can be expected from the appropriate lowering of the pH. This means of preparation makes it possible to change the particle size in the range 30−450 nm simply by adjusting the preparation conditions.

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EXPERIMENTAL SECTION Monodisperse CoHCF nanoparticles are typically prepared as follows. Maintained at 20 °C in a water bath, 5.0 mL of Received: November 8, 2011 Revised: December 22, 2011 Published: January 9, 2012 3394

dx.doi.org/10.1021/jp210707y | J. Phys. Chem. C 2012, 116, 3394−3399

The Journal of Physical Chemistry C

Article

were cubic and fairly monodisperse. The mean size ± one standard deviation, observed from the TEM images, was estimated to be 160 nm ±16 nm. The XRD pattern of the CoHCF nanoparticles is shown in Figure 2. Concordant with

Co(NO3)2 and 5.0 mL of K3[Fe(CN)6] aqueous solutions (0.075 mol L−1 and 0.050 mol L−1, respectively) were added in this order to 40 mL of 0.25 mol L−1 trisodium citrate aqueous solution, which also contained 0.10 g of gelatin as a protective colloid (pH 7.8). Next, 10 mL of HNO3 (2.0 mol L−1) was added to lower the pH to 3.4 by a transfer pipet under stirring (addition time 15 s). A summary of the content of each reagent in the reacting solution is given in Table 1. After 60 min passed, Table 1. Composition of Reagents in Reacting Solution for Preparation of Monodisperse CoHCF Nanoparticles (Total Volume 60 mL) reagent

content

gelatin trisodium citrate Co(NO3)2 K3[Fe(CN)6] HNO3

0.17% 0.17 mol L−1 6.3 × 10−3 mol L−1 4.2 × 10−3 mol L−1 0.33 mol L−1

Figure 2. XRD pattern of the CoHCF nanoparticles shown in Figure 1.

the dispersion was centrifuged at 10000 rpm for 30 min, and the supernatant solution was removed. The precipitate was redispersed in distilled water and centrifuged again (this was repeated twice). The shape of the particles was observed with a scanning electron microscope (SEM, JEOL JSM-6700F) at 15 kV. A transmission electron microscope (TEM, Hitachi H-7650) was used at 100 kV to evaluate the size distribution, and at 120 kV to observe high-resolution TEM (HR-TEM) images. The atomic ratio of Co to Fe in the particles was confirmed by energy-dispersive X-ray spectroscopy (EDX) using an EDAX MX2T system (detector: rTEM) installed on the TEM. The crystal structure was identified by X-ray diffraction analysis (XRD) using a Mac Science M18X-HF-SRA (CuKα radiation, λ = 1.5406 Å). Reagents of cobalt(II) nitrate hexahydrate, potassium hexacyanoferrate(III), and trisodium citrate dihydrate were purchased from Wako Pure Chemicals and were used as received. The inert deionized gelatin used was of photographic grade and was manufactured from bovine bone via an alkaline process (DGF-Stoess, #69827).

the literature,10,21 all peaks were indexed as being of facecentered cubic (fcc) structure (lattice parameter = 10.28 Å). The HR-TEM image shown in Figure 3, where additional

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RESULTS AND DISCUSSION Figure 1 shows an SEM image of the CoHCF nanoparticles prepared under the conditions listed in Table 1. The particles

Figure 3. HR-TEM image of a CoHCF nanoparticle. A lattice spacing of 5.1 Å for (200) faces is observed.

treatment with a 0.4% pepsin in a 0.01 mol L−1 HCl solution was applied in order to diminish the amount of gelatin from the particle surface. The lattice spacing of 5.1 Å for the (200) faces was observed. The CoHCF particle seems to be single crystalline basically, but the lattice image that slightly bends suggests the existence of some defects. The atomic ratio Co/Fe = 59/41 that was estimated by TEM−EDX analysis was in good agreement with the stoichiometric ratio of CoHCF (Co/ Fe = 3/2). A small amount of potassium, about 5% of the Co content, was also detected, suggesting the partial formation of K2Co[Fe(CN)6] species6 caused by the possible coexistence of [Fe(CN)6]4− that was contaminated by a reduction of [Fe(CN)6]3−. The starting solution, which was composed of trisodium citrate, cobalt(II) nitrate, potassium hexacyanoferrate(III), and gelatin, was an orange-colored transparent solution. This transparent state was the result of there being a sufficient

Figure 1. SEM image of CoHCF nanoparticles obtained under the conditions listed in Table 1. 3395



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supersaturation ratio, and nuclei of CoHCF were formed. Once the solid phase appears in the liquid phase, the solid−liquid equilibrium must be considered, in addition to the equilibria of the citrate species. Accordingly, the nuclei consumed the Co2+ existing in a liquid phase in order to satisfy the precipitation equilibrium defined by [Co2+] and [HCF3−], where [Co2+] and [HCF3−] were the concentrations of free Co2+ and [Fe(CN)6]3− ions, respectively. The decrease in [Co2+], however, was compensated for by the release of Co2+ from the citrate complexes via reconstruction of the complex equilibrium. This series of processes acted as a process for the supply of Co2+ for growth via spontaneous mass-transfer, and continued until a new state of equilibrium was established. In other words, the nuclei pulled Co2+ out of the complexes, as has been previously reported for silver halide systems,23,24 and the self-growth mechanism played a key role in the separation of the growth period from the nucleation period. This consideration could be confirmed by estimating the [Co2+] at each step of preparation on the basis of the equilibrium relationships. Citric acid (HOC(CH2COOH)2COOH) has three stepwise dissociation constants for H+ from each carboxyl group: log Ka1 = −2.87, log Ka2 = −4.35, and log Ka3 = −5.68.20 The dissociated citrate ions can form a complex with Co2+, and their association constants are log KCoH2L = 1.25, log KCoHL = 3.02, and log KCoL = 5.00 for the CoH2L+, CoHL, and CoL− complexes, respectively,20 where L3− represents a trivalent citrate ion. The solubility product of CoHCF, Ksp (≡ [Co2+]3[HCF3−]2) = 6.7 × 10−22 mol5 L−5, is also assumed (see Supporting Information, SI). These equilibrium constants are applied to the estimation of [Co2+] under the preparation conditions given in Table 1 by numerically solving eqs S11 and S16, as described in the SI. Before the addition of acid (pH 7.8), [Co2+] = 4.0 × 10−7 mol L−1 is obtained, which increases to 2.7 × 10−4 mol L−1 at pH 3.4 in the absence of CoHCF precipitates. The latter concentration corresponds to only 4.3% of the total Co2+ content in the system. After the formation of the precipitate, [Co2+] in equilibrium is expected to be low as 1.9 × 10−5 mol L−1, and 93% of Co2+ would be contained in the solid phase. While these values are just estimations, they are still sufficient to support the formation mechanism described above. In the present system, the main role of gelatin is to act as a protective colloid to prevent coagulation among particles. This was adequately achieved when its concentration was 0.1% or above. Moreover, binding with Co2+ contributed to the increased stability of the starting solution, as mentioned above. Hence, the effect of the Co2+−gelatin complex on particle formation needs to be evaluated. Deprotonated carboxyl groups in side chains of aspartic acid and glutamic acid residues are the binding sites for Co2+ around the formation pH of CoHCF (pH 3.4).22,25 The number of aspartic acid and glutamic acid residues per 1000 residues are reported as 42 and 73, respectively, in bovine bone gelatin.26 These values correspond to 4.7 × 10−4 and 8.1 × 10−4 mol/gram of gelatin; thus, 2.1 × 10−3 mol L−1 of the total Co2+ binding sites are estimated for a 0.17% gelatin solution. On the other hand, the dissociation constant of the hydroxyl group in a gelatin molecule has been reported as pKaGel = 4.2527−4.7,28 so a deprotonated site is expected to be only 1.0−2.6 × 10−4 mol L−1. Despite the fact that the association constant of the −COO− group of protein with Co2+ (pKCoGel = 3.029) was larger than it was for the CoH2L+ complex (the dominant component at pH 3.4, pKCoH2L = 1.2520), the fact that there

concentration of citrate, with the assistance of gelatin. That is, in the absence of gelatin and using Co(NO3)2 and K3[Fe(CN)6] contents that were twice those in Table 1 (this condition will also be used below), CoHCF was formed in the starting solution without the addition of acid, although the formation rate was quite slow. In the presence of gelatin, on the other hand, such unexpected precipitation was prevented owing to the additional binding of Co2+ to gelatin molecules.22 Hence, gelatin was important not only as a protective colloid, but also as a complexing agent of Co2+. By adding acid to shift the pH from 7.8 to 3.4, the equilibrium state was shifted, and Co2+ was released. In spite of the rapid shift in pH, however, an induction period of about 25 s was observed before a state of turbidity was initiated. This turbidity then increased gradually. Since the intrinsic rate of the reaction between Co2+ and [Fe(CN)6]3− is rapid, as mentioned above, the existence of an induction period implies a controlled release of Co2+. The formation of monodisperse CoHCF particles also suggests that satisfaction of the LaMer mechanism was achieved, which involves a short nucleation period followed by a relatively long growth period. In fact, the increase in turbidity, which reflects the progress of growth, was not so rapid, as is shown in Figure 4 (the inset shows the spectra

Figure 4. Time evolution of optical density at 700 nm during the formation of CoHCF nanoparticles shown in Figure 1. The inset shows the change of spectrum in the process. The optical path length is 2 mm.

during the reaction), where turbidity is indicated by optical density at 700 nm (2 mm of the optical path length) as measured by a photodiode-array spectrophotometer (Shimadzu Multispec-1500). A slower rate of addition (addition time of 200 s) worsened the monodispersity, owing to a prolonged shift in the equilibrium that produced a nucleation stage of longer duration. A rapid rate of addition (1.5 s) also widened the size distribution, probably owing to the formation of a nonuniform, extremely lower pH region around the point of addition, which could cause the uncontrolled release of Co2+. Thus, an appropriate addition rate is required for the formation of monodisperse particles. The shift in pH was a prompt response to the progressive addition of acid, which suggests the quick protonation of citrate, and thus, a probably rapid rate of release of Co2+. Taking into account the intrinsically fast reaction of Co2+ with [Fe(CN)6]3−, the following scheme of formation is most probable. Namely, only a part of the Co2+ was released by shifts in the acid dissociation and in the complex equilibria of citrate when the acid was added. At this moment, most of the Co2+ still remained in the form of a citrate complex. The amount of released Co2+, however, was sufficient to exceed the critical 3396



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Table 2. Contents of Co(NO3)2 and K3[Fe(CN)6] in Reacting Solution and Estimated Value of Concentration Product α for 3 2 CoHCF Preparation, where α ≡ [Co2+]calc × [HCF3−]ini at pH 3.4 reactant content/mol L−1 Co(NO3)2 condition condition condition condition condition a

1a 2 3 4 5

6.3 × 10−3 6.3 × 10−3 9.4 × 10−3 9.4 × 10−3 1.25 × 10−2

α/mol5 L−5

K3[Fe(CN)6] 4.2 8.3 8.3 6.3 8.3

× × × × ×

HNO3 addition

10−3 10−3 10−3 10−3 10−3

3.55 1.40 4.86 2.75 1.22

× × × × ×

10−16 10−15 10−15 10−15 10−14

H2SO4 addition 1.93 7.59 2.59 1.46 6.34

× × × × ×

10−17 10−17 10−16 10−16 10−16

The same condition as in Table 1

The adsorption of SO42− on the CoHCF surface, acting as a growth modifier,31 might be another explanation of the size change, but this seems less likely in this case, because reduced υ yields smaller particles as a result of a larger n, as is apparent from eq 2. Although adsorbed gelatin molecules, rather than SO42−, would decrease υ from the intrinsic rate, a constant υ is expected, at least with the same concentration of gelatin. Hence, it is reasonable to omit the adsorption effects in order to simplify the discussion. To consider the formation process further, the Co(NO3)2 and K3[Fe(CN)6] contents were varied in the addition of both HNO3 and H2SO4, as listed in Table 2, where the other conditions of the citrate and gelatin contents, the amount of H+ added, and the temperature were the same as in Table 1. Under all these conditions, the initial and final pH values were practically same. In Figure 5, the particle size is plotted with

were much fewer binding sites of gelatin suggests a small contribution of gelatin to Co2+ storage, 3%−9% of it by citrate complexes, on the particle precipitation process at pH 3.4. In addition, the behavior of gelatin on the storage and release of Co2+ by a pH shift is basically the same as that for citrate because it depends on the dissociation of functional groups of amino acid residues. Therefore, in considering the fundamental formation mechanism, the contribution of gelatin is not so significant. With the addition of 10 mL of H2SO4 (1.0 mol L−1) instead of HNO3, larger particles were obtained, and the final pH was the same as it was with the addition of HNO3. The particle size was 440 nm ±46 nm, and the shape was also cubic. In this case, the effect of the formation of a soluble CoSO4 complex needs to be introduced into the [Co2+] calculation, in which the association constant is log KCoSO4 = 2.36.30 At [SO42‑] = 0.17 mol L−1 in the reacting solution, [Co2+] = 1.0 × 10−4 mol L−1 is estimated at pH 3.4 before nucleation, which is about one-third the value with the addition of HNO3. This decrease in [Co2+] is possibly the reason for the size change. That is, under the condition of a consistent reactant content and yield, the particle size of a monodisperse system is determined by the number of formed particles, n, which can be expressed by a theoretical formula for closed systems as

n = QVm/υ

(2)

where Q, Vm, and υ are the supply rate of monomers, the molar volume of a solid, and the rate of volume increase of the nucleus in the nucleation period, respectively.23 In the present system, the monomer is a Co3[Fe(CN)6]2 molecule, and thus Q must be expressed in terms of [Co2+]3 and [HCF3−]2. Because of the identical [HCF3−] and acid addition time in the HNO3 and H2SO4 systems, the ratio between Q in them may be denoted by the ratio in [Co2+]3 for free precipitates at pH 3.4. The ratio, QNA/QSA, is estimated as 19.7 from the values for [Co2+] given above, and thus nNA/nSA = 19.7 is expected, under the assumption of consistent υ, where the subscripts NA and SA represent the addition of HNO3 and H2SO4, respectively. The particle size, L, is proportional to n−1/3, and LNA/LSA corresponds to 0.37. This ratio is in good agreement with the experimental result of 160 nm/440 nm = 0.36. The concentration of free Co2+ gradually increases with the progressive addition of acid. When the supersaturation ratio, S, exceeds the critical supersaturation ratio, S*, nucleation occurs immediately. The supplied monomers are consumed by the formation, and the growth of nuclei follows in this stage, so that an increase of S would be suppressed. Hence, the assumption of constant υ may be reasonable, in which υ is a function of S. In other words, the difference between nNA and nSA, or between LNA and LSA, is caused by the difference in Q.

Figure 5. Effect of concentration product α on the mean size of particles, where α is defined by eq 3.

respect to the product of the concentrations in the nucleation period, α, which is defined in the same dimension as Ksp by

α ≡ [Co2 +]3calc × [HCF3 −]2ini 2+

(3)

2+

where [Co ]calc is the Co concentration calculated from eq S11 in the SI at pH 3.4 in the absence of CoHCF precipitates, and [HCF3−]ini is the initial concentration of [Fe(CN)6]3−, which corresponds to the K3[Fe(CN)6] content in Table 2. The particle size was obviously reduced with higher α. Moreover, it should be noted that the change in size was continuous among all the tested conditions regardless of which acid was used. In addition, a shortened induction period and a faster increase in turbidity were observed at higher α. Figure 6 shows the relationship between α and n, where the latter is calculated from the Co2+ content and the lattice 3397



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nucleation from spontaneous mass-transfer for the growth step caused by a shift in equilibria.

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ASSOCIATED CONTENT

S Supporting Information *

Information regarding the derivation of equations for the calculation of [Co2+] and the experimental estimation of the solubility product for CoHCF is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone and fax: +81-43-290-3492; e-mail: [email protected].

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Figure 6. Relationship between concentration product α and particle number. The slope is 1.39.

ACKNOWLEDGMENTS This research was partially supported by Grant-in-Aid for Scientific Research No. 23550223.

parameter of the fcc structure (4/3 molecules per unit cell).32 The double logarithmic plot shows a linear relationship, which suggests that the nucleation process is mainly determined by α. The slope in Figure 6 is 1.39, although unity of the slope is expected if α directly reflects Q. The somewhat larger slope of the experimental result implies an imperfect proportionality of Q to α, which is probably due to the consecutive change in the equilibrium state during the addition of acid in a certain period. When both the Co(NO3)2 and K3[Fe(CN)6] contents were decreased to half of that in condition 1 in Table 2 for the addition of H2SO4, no precipitate was generated within at least 60 min of the reaction. Thus, S* may exist between condition 1 and half of the values of this condition. The concentration of Co2+ in equilibrium with the CoHCF solid at the moment of the initiation of nucleation, [Co2+]o, is estimated from the values of Ksp and the K3[Fe(CN)6] content as [Co2+]o = 3.4 × 10−6 mol L−1 for condition 1, while [Co2+]o increased to [Co2+]o = 5.3 × 10−6 mol L−1 when the K3[Fe(CN)6] content is decreased by half. On the other hand, the values of [Co2+]calc are estimated as 1.0 × 10−4 mol L−1 and 5.2 × 10−5 mol L−1 for condition 1, and for half of the values of this condition, respectively. By employing these values, S*(≡[Co2+]calc/ [Co2+]o) is estimated as being in the range of 9.8−29. The large difference in the experimental condition caused the candidate S* to spread broadly. However, the estimation seems within a reasonable range of S* as compared with other homogeneous nucleation systems, such as 2.9−4.0 for silver halides,23,24 30 for MgF2,33 and 106 for PbCO3.33 This consideration also supports the proposed formation mechanism of the present monodisperse CoHCF system.

REFERENCES

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CONCLUSION Monodisperse cubic cobalt(II) hexacyanoferrate(III) (CoHCF) nanoparticles were prepared using a cobalt−citrate complex as a reservoir of Co2+ to regulate the formation reaction properly. By adding acid, a small part of Co2+ was released and reacted with [Fe(CN)6]3− to form CoHCF nuclei, which subsequently pulled out almost all Co2+ from the citrate complex in order to satisfy the precipitation equilibrium. This mechanism is supported by calculations and evaluation of the Co 2+ concentration on the basis of equilibrium relationships, including the acid dissociation of citric acid, the complex formations of free Co2+ with citrate and sulfate ions, and the solubility product of CoHCF. It is concluded that the present system consists of an ideally separated stage of enforced 3398



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Preparation of Monodisperse Cobalt(II) Hexacyanoferrate(III

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