Introduction

Owing to the wide applications of core–shell nanostructures in different fields like catalysts, industrial and biomedical applications [1,2,3,4,5,6,7,8,9,10,11,12,13], they become having great attention in few decades. Also, the core–shell nanocomposites have core with different sizes, shapes, thicknesses and morphologies. The core–shell nanostructures can be taken different shapes such spherical, centric, eccentric, starlike or tubular in shape. The size, shape and properties of core–shell are different from one material to another. Also, core–shell nanoparticles showed great important applications in medical biotechnology such as molecular bioimaging, drug delivery system and cancer therapy. The properties of nanoparticles were improved by functional groups or molecules or coated with a thin layer of other materials with different constituents than the non-functionalized uncoated particles [14,15,16,17].

One of the fission by-products of most current nuclear power plants is cesium. The release of these radioactive isotopes like 134Cs+ and 137Cs+ can show significant biohazard with the long-lived radionuclides 2.1 and 30 years, respectively, that accumulating in the environment [18]. The migration of cesium through groundwater to the biosphere returns to its highly solubility, which has an effect on human health and also on the environment. Also, cesium incorporated in terrestrial and aquatic organisms because it is chemically similar to potassium [19]. When cesium ions taken by ingestion, it is adsorbed and distributed throughout the soft tissues of body. Thyroid cancer is one of the consequences of 137Cs adsorption in the contaminated food and water is [20].

The separation of cesium from wastewater is developed by using different technologies like extraction, ion exchange, chemical precipitation, adsorption and electrochemistry [21]. One of the above technologies that have recently attracted a lot of attentions is ion exchange process due to its convenience, the high efficiency and selectivity [22].

Hexacyanoferrates of transition metals (Fe, Cu, Zn, Ni, Co) are used as a good excellent ion exchanger for cesium ion removal from aquatic solution [23]. Among these, MHCFs is copper hexacyanoferrate (CHCF) [24], which gives good results in a wide range of pH like high adsorption capacity, selectivity and chemical stability. The unfeasibly separation of CHCF from aquatic solution after the adsorption is considered one of the disadvantages of using it to remove cesium radionuclides from environment. This disadvantage can be treated by immobilization the CHCF powders through different supporting materials to fabricate composite adsorbents. This process is achieved by implantation of inorganic into the wide range of organic materials during the polymerization process to obtain the composite ion exchangers [25]. The inorganic materials considered active components and organic materials are simply inert binders in composite.

From through chelating resin, the CHCF has been immobilized in them [26]; polyurethane foam [27], polyacrylonitrile (PAN) [28], polyether sulfone (PES) [29], polycarbonate track-etched member [30] and PVA are used to obtain the composite ion exchangers for the removal of cesium ions.

Ortho phenylenediamine (OPD) has a structure consisting of 1,4-substituted benzenoid and quinoid moieties and was proposed for the polymer obtained electrochemically and it exhibits high thermostability [31]. OPD is soluble electroactive in the polymer due to its dissolution in organic solvents. Due to the presence of two oxidizable groups (–NH2) attached to the aromatic ring in o-phenylenediamine, these sites are more reactive, so they play role in the polymerization, so that OPD has electrochemical behavior as anthranilic acid than anilines and phenols.

In recent years, synthetic polymer is used in adsorption experiments. Due to its macropores, synthetic polymer showed an excellent advantage as an adsorbent in continuous and batch adsorption process [32].

In our previous work, nanoparticles of copolymer of anthranilic acid with o-phenylenediamine poly(AA-co-OPD) were synthesized and used as chelating resin for the removal of copper ions from wastewater [33].

In this work, a new core–shell nanocomposite (CSNC) of copper hexacyanoferrate copolymer of anthranilic acid with o-phenylenediamine poly(AA-co-OPD) has been synthesized and used as nanocomposite ion exchanger for extracting cesium ions. The CHCF nanoparticles were implanted in poly(AA-co-OPD) during the copolymerization process. The synthesized CSNC was characterized by FT-IR, XRD and TGA. The surface morphology was also studied by TEM and SEM. The sorption of cesium from wastewater was investigated by batch technique as a function in pH of cesium solution, cesium ion concentration, temperature and contact time. Also, the regeneration of cesium was studied.

Synthesis and characterization techniques

Materials

Ammonium peroxydisulfate (APS) was obtained from Merck, India. Anthranilic acid was obtained from Rolex India Ltd, India, and o-phenylenediamine (RDH), copper sulfate (CuSO4·5H2O), potassium hexacyanoferrate (K4[Fe (CN)6]0.3H2O) and polyvinyl alcohol (PVA) were purchased from Sigma-Aldrich Co., USA. Non-radioactive cesium chloride (CsCl) was purchased from Alfa Aesar (China). Hydrochloric acid was obtained from (Merck, India). Deionized water was used to the preparation of the aqueous solutions.

Synthesis of the nanoparticles of copper hexacyanoferrate (CHCF)

Potassium hexacyanoferrate (10 wt%) was dispersed in 50 ml of distilled water and it was stirred for 1 h. 10 wt% of CuSO4·5H2O was dissolved in 50 ml distilled water, mixed with the potassium hexacyanoferrate solution drop by drop and stirred for 1 h. Upon the addition of CuSO4·5H2O to the K4[Fe (CN)6]0.3H2O solution, heavy reddish brown precipitate was observed with naked eye. Then, PVA (6 wt%) was dissolved in 100 ml distilled water and added drop by drop to the above-mentioned mixture of solutions. The reaction was left on magnetic stirrer for 24 h at room temperature. The reddish brown precipitate was filtered off next day and washed with distilled water for several times. Finally, the obtained reddish brown powder product was dried in air at room temperature and it was referred as CHCF.

Synthesis of core–shell nanocomposite (CSNC) of copper hexacyanoferrate copolymer of anthranilic acid with o-phenylenediamine (poly(AA-co-OPD))

In a typical synthesis, anthranilic acid (40 mmol) and o-phenylenediamine (10 mmol) were dissolved in 20 ml of 0.1 M HCl solution with magnetic stirring until they completely dissolved. 3 wt% of nanoparticles of CHCF was dispersed in 15 ml of 0.1 M HCl and added drop by drop to the above solution with stirring. A fresh solution of APS (1.2 M) in 20 ml of 0.1 M HCl solution was rapidly added to the above solution containing anthranilic acid, o-phenylenediamine and CHCF nanoparticles. The reaction was left on magnetic stirrer for 24 h. The black precipitate was filtered off next day, washed with 0.1 M of HCl followed by deionized water several times and dried in air at room temperature to give the black precipitate of CSNC.

Characterization techniques

Morphology properties of core–shell nanocomposite

Morphology and particle size of core–shell nanocomposite were determined using TEM (JEOL [JEM-1230 electron microscopy]), and the sample was prepared for measuring by the dispersion of the sample in water using ultra-sonication and then taken one drop of the dispersed sample on a carbon coated copper grid and evaporated and finally placed in the Phillips (CM/TEM). Also, SEM (SEM, FEI Inspect S, Oxford USA) with acceleration beam (25–30) k.v with vacuum pressure 60 Pa and spot size (5–6) using backscattering detector for Z-imaging and an X-ray diffractometer was used to determine the surface morphology and the size of core–shell nanocomposite particles.

FT-IR measurements

The FT-IR spectra for the prepared CHCF nanoparticles and the core–shell nanocomposite were measured using Thermofisher Nicolete IS10, USA between 400 and 4000 cm−1 spectrophotometer.

Thermal gravimetric analysis

Thermal analysis experiments including thermogravimetric analysis (TGA) for the core–shell nanocomposite were carried out using SDT Q600 V20.9 Build 20, USA, thermogravimetric analyzer. The measurement was taken in a dynamic atmosphere of nitrogen from room temperature to 1000 °C at heating rate of 10 °C min−1.

Surface area

Porous structure parameters of the core–shell nanocomposite were described by Brunauer–Emmett–Teller (BET) and BJH methods through N2 adsorption–desorption methods to inspect the porous properties of the CSNC using nitrogen as the adsorbent at 77.35 K. Using a model NOVA 3200 automated gas sorption system (Quantachrome, USA), the measurements were taken.

X-ray diffraction (XRD) of the prepared CHCF nanoparticles and CSNC

The diffraction of the CHCF nanoparticles and the core–shell nanocomposite (CSNC) was obtained from Rigaku Oxford XtaLAB pro.

Determination of the cesium ion concentration remaining after the adsorption process

Cesium ion concentration was determined by using Agilent Technologies 700 Series ICP-OES.

Removal of Cs+ ions using CSNC by batch technique

The adsorption experiment has been investigated by batch technique. 0.1 g of the prepared core–shell nanocomposite was shaken with 100 ml of cesium solution (3 mmol L−1) at 25 °C. At the end of time, the solution was fltered for phase separation and the concentration of supernatant solutions was measured by using Agilent Technologies 700 Series ICP-OES. Before adding the adsorbent, an aliquot of the bulk solution was withdrawn for the measurement of the initial concentration. For the evaluations of adsorption capacity (qe) of the prepared core–shell nanocomposite, Eq. (1) was applied.

To study the influence of pH on Cs+ uptake, 100 ml of Cs+ solution (3 mmol L−1) was mixed with 0.1 g of CSNC at different pH values in the range from 3 to 11. Solutions of 0.1 M HCl and 0.1 M NaOH were used to adjust desired pH of cesium solutions. At different concentrations of Cs+, the adsorption was studied. Adsorption experiments with different initial concentrations of Cs+ from 2 to 10 mmol L−1 were performed at optimum pH. The effect of shaking time on adsorption of Cs+ was investigated, in which the experiments were performed at different times (0.5–3.5 h.) at optimum pH and optimum concentration of Cs+. The change in temperature on the cesium uptake was obtained at different temperatures 25, 40 and 60 °C, at optimum pH and 10 mmol L−1 of Cs+.

Desorption process of cesium ions from CSNC

The regeneration of Cs ions that loaded onto CSNC was achieved using 2 M KCl at 25 °C. To study the reusability of CSNC, five adsorption–desorption cycles were performed.

Theory/calculation

Adsorption capacity

Using the following equation [34], the capacity qe (mmol g−1) of cesium onto CSNC can be calculated.

$$q = \frac{{\left( {C_{0} - C_{{\text{e}}} } \right)V}}{W}$$
(1)

where Co (mmol L−1) is the initial Cs+ concentrations, whereas Ce (mmol L−1) is the equilibrium Cs+ concentrations. V (L) is the volume of Cs+ solution taken for experiment, and W is the mass of dry CSNC in grams.

Adsorption isotherm models

To understand the mechanism of the removal of Cs+ using CSNC, four isotherm models were applied. The four isotherm models include Langmuir (Eq. 2) [35], Freundlich (Eq. 3) [36], Temkin (Eq. 4) [37] and Dubinin–Radushkevich (D–R) (Eq. 5) [38], which were examined for studying the distribution of Cs+ between the solution of cesium and CSNC. After applying four isotherm models, the Cs+ affinity, surface properties, sorption capacity and sorption mechanism can be explained.

$$\frac{{C_{{\text{e}}} }}{q} = \frac{{C_{{\text{e}}} }}{{Q_{\max } }} + \frac{1 }{{K Q_{\max } }}$$
(2)
$$\log \, q = N\log C_{{\text{e}}} + \log K_{{\text{F}}}$$
(3)
$$q = B\ln K_{{\text{T}}} + B\ln C_{{\text{e}}}$$
(4)
$${\text{Ln }}q_{{\text{e}}} = \ln Q_{\max } - K_{{{\text{DR}}}} \varepsilon^{2}$$
(5)

where Ce (mmol L−1) is the concentration of Cs+ after equilibrium; q and Qmax (mmol g−1) are amount of Cs+ adsorbed at equilibrium and maximum adsorption, respectively; K (L mmol−1) is the Langmuir affinity constant; and KF (mmol g−1) and N are the Freundlich constants. KT (Lg−1) and B (J mol−1) are constants related to equilibrium binding and heat of adsorption, respectively. KDR (mol2 kJ−2) is a constant for the mean free energy of adsorption, and ε (J mol−1) represents the Polanyi potential which can be calculated from the equilibrium concentration (Ce) as indicated in the following equation:

$$\varepsilon = {\text{RT}}\ln \left( { 1 + \frac{1}{{C_{{\text{e}}} }}} \right)$$
(6)

The free energy change when one mole of ion is transferred to the surface of the solid from infinity in the solution is pointed to the mean sorption energy E (kJ mol−1), which can be calculated according to the following equation [39]:

$$E = \frac{1}{{\sqrt {2K} }}$$
(7)

By studying the Langmuir sorption model, if the data were fitted with this model, the sorption of cesium is homogeneous adsorption, in which the solute molecules are adsorbed onto a homogenous and flat surface of adsorbent. The Langmuir adsorption model can be expressed by fundamental characteristics dimensionless constant (RL), which is used to explain the relation between the adsorbent and adsorbate. The RL value can be calculated from the following equation:

$$R_{{\text{L}}} = \frac{1}{{1 + {\text{KC}}_{{\text{o}}} }}$$
(8)

The RL values show the state of the isotherm process. Since the isotherm is unfavorable when RL > 1, whereas becomes linear when RL = 1, at condition of 0 < RL < 1 it becomes favorable and is irreversible when RL = 0 [40].

On the other hand, the heterogeneous adsorption can be confirmed by Freundlich adsorption model. The third isotherm model is Temkin isotherm model which assumes that the heat of sorption decreases linearly with coverage due to indirect adsorbate–adsorbate interactions. The adsorption mechanism with a Gaussian energy distribution onto a heterogeneous surface can be explained by the fourth isotherm model of D–R model.

Adsorption kinetic models

The kinetic process of adsorption of Cs+ by CSNC can be investigated by using four kinetic models, including the pseudo-first-order (Eq. 9) [41], pseudo-second-order (Eq. 10) [42], intra-particle diffusion model (Eq. 11) [43] and Elovich equation (Eq. 12) [44].

$$\log \left( {q_{{\text{e}}} - q_{{\text{t}}} } \right) = \log q - \left( {\frac{{K_{{{\text{ads}}}} }}{2.303}} \right)t$$
(9)
$$\frac{t}{{q_{{\text{t}}} }} = \frac{1}{{K_{2} q^{2} }} + \left( \frac{1}{q} \right)t$$
(10-1)
$$h = K_{2} q^{2}$$
(10-2)
$$q_{{\text{t}}} = K_{{{\text{id}}}} t^{0.5}$$
(11)
$$q_{{\text{t}}} = \frac{1}{\beta }\ln \left( {\alpha \beta } \right) + \frac{1}{\beta }\ln t$$
(12)

where qe and qt are capacities of Cs+ adsorbed onto CSNC (mmol g−1) at equilibrium and at any time t (min), respectively; Kads (min−1), K2 (g mmol−1 min−1) and Kid (mmol g−1 min−0.5) are the rate constants of the first-order, second-order and intra-particle diffusion models, respectively; h (mmol g−1 min−1) is the initial sorption rate constant; α (mmol g−1 min−1) is the initial adsorption rate; and β (g mmol−1) is the constant of the degree of the surface coverage and activation energy for chemisorption.

Thermodynamic parameters

The change of the standard Gibbs free energy ΔG°ads (kJ mol−1), the change of the standard enthalpy ΔH°ads (kJ mol−1) and the change of the standard entropy ΔS°ads (J mol−1) are thermodynamic parameters, which can be obtained using Eqs. (13) and (14):

$$\ln K_{{\text{d}}} = \frac{{\Delta S_{{{\text{ads}}}}^{^\circ } }}{R} - \frac{{\Delta H_{{{\text{ads}}}}^{^\circ } }}{RT}$$
(13)
$$\Delta G^{o}_{{{\text{ads}}}} = - RT\ln K_{{\text{d}}}$$
(14)

where T is the absolute temperature (K), R is the universal gas constant (8.314 J mol−1 K−1) and Kd is the equilibrium distribution constant (ml g−1). Kd was given by Eq. (15) [34].

$$K_{{\text{d}}} = \frac{{C_{{\text{o}}} - C_{{\text{e}}} }}{{C_{{\text{e}}} }} \times \frac{V}{W}$$
(15)

where Co (mmol L−1) is the initial Cs+ concentration, Ce (mmol L−1) is the equilibrium Cs+ concentration (mmol L−1), V is the total volume of the solution in (L) and W is the mass of the dry CSNC in gram.

Regeneration efficiency

From the uptake elution process, the regeneration efficiency can be calculated by the following equation [45]:

$${\text{Regeneration}}\;\;{\text{efficiency}} = \frac{{{\text{Uptake}}\;{\text{of}}\;{\text{metal}}\;{\text{ion}}\;{\text{in}} \;{\text{the}}\;{\text{second}}\;{\text{cycle}}}}{{{\text{Uptake}}\;{\text{of}}\;{\text{metal}}\;{\text{ion}}\;{\text{in}}\;{\text{the}}\;{\text{first}}\;{\text{cycle}}}} \times 100$$
(16)

Result and discussion

Synthesis of core–shell nanocomposite (CSNC) of CHCF-poly(AA-co-OPD)

A new CSNC was synthesized by the plantation of nanoparticles of CHCF during the copolymerization of anthranilic acid with o-phenylenediamine. CHCF was prepared in nanoscale by using PVA as stabilizer agent to control the average particle size. Scheme 1 illustrates the mechanism of copolymerization of anthranilic acid with o-phenylenediamine. The synthesized CSNC is used for the removal of Cs+ ions found in wastewater by ion exchange process, in which the potassium ions in CSNC were replaced by cesium ions as shown in Scheme 2. [31, 33].

Scheme 1
scheme 1

Mechanism of copolymerization of anthranilic acid and o-phenylenediamine

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Scheme 2
scheme 2

Ion exchange process of potassium ions in CSNC by cesium ions in aqueous solutions

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Characterization of the synthesized CSNC

Morphological properties of CHCF and CSNC of CHCF copolymer of anthranilic acid with o-phenylenediamine (poly(AA-co-OPD))

Figure 1 shows the morphology of synthesized CHCF by SEM and TEM, which was cuboid powder with the nanometer size of 25–50 nm. The core–shell structure of the nanocomposite was confirmed by TEM techniques as shown in Fig. 2. The SEM–EDX and mapping spectra of CSNC samples before and after cesium adsorption were studied and are shown in Figs. 3, 4. The SEM images in Fig. 3a were studied to determine the distribution of the CHCF nanoparticles through poly(AA-co-OPD). These images indicated that the CHCF nanoparticles were uniformly distributed throughout the sample. Following the cesium adsorption, it was shown that cesium was uniformly adsorbed and the surface becomes brighter which is thought to be cesium adsorption that was emerged at the after adsorption of cesium ions (Fig. 4a, as compared with Fig. 3a).

Fig. 1
figure 1

a SEM and b TEM of the synthesized CHCF

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Fig. 2
figure 2

TEM of the synthesized CSNC

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Fig. 3
figure 3

SEM–EDX and mapping spectra of CSNC sample before cesium adsorption

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Fig. 4
figure 4

SEM–EDX and mapping spectra of CSNC sample after cesium adsorption

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The adsorption of Cs+ ions by CSNC was studied by EDX elemental mapping spectra analysis with SEM images. The cesium ion peaks were clearly appeared onto the surface of immobilized CSNC in Fig. 4b compared with Fig. 3b. Also, the distribution of K+ in composite before and after the adsorption of Cs+ ions is shown in Figs. 3b, 4b, which showed that the concentration of K+ was decreased slightly after cesium adsorption. To observe the distribution of Cs+ and K+ ions before and after adsorption process (Figs. 3c, 4c), elemental mapping for CSNC samples before and after cesium adsorption was also studied. The images showed that homogeneous distribution of cesium ions with the decrease in the distribution of K+ due to adsorption process occurs with good efficiency.

FT-IR spectra of synthesized nanoparticles

Figure 5 shows the spectra of the synthesized CHCF nanoparticles and CSNC. In Fig. 5a, the spectrum of CHCF exhibits sharp peak at 2095 cm−1 due to the stretching vibration of cyanide group (C≡N) and the bands in the 480–592 cm−1 region are due to the (Fe–C) stretching. Figure 5b shows the spectrum of CSNC, which exhibits (–NH2) primary amino groups in the FT-IR spectrum which indicates that the composite was formed through the head-to-tail coupling of the two monomers via –NH– groups [31, 46]. The peak that observed at 3410 cm−1 is due to primary amine (–NH2) group [47]. At 1685 cm−1, a strong band appears for the C=O stretching of carboxylic group. The spectrum exhibits main band in the range 1611 cm−1 corresponding to the C=C stretching frequency of benzenoid and quinoid rings, respectively. The range of 1114 cm−1 is due to C=N stretching in the aromatic ring. Peaks at 856–922 cm−1 can be assigned to 1,2,4-tri-substituted benzene structure [48]. These data revealed to the emeraldine form for nanocomposite of poly(AA-co-OPD). Also, Fig. 5b shows broadband from 2500 to 3500 cm−1 of carboxylic OH and sharp peak at 2095 cm−1 due to the stretching vibration of cyanide group (C≡N) similar to that peak found in Fig. 5a for CHCF, so this observation proved the presence of CHCF inside poly(AA-co-OPD) to produce CSNC instead of coordination between copper and carboxylic group.

Fig. 5
figure 5

FT-IR spectra for a CHCF nanoparticles and b CSNC

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Surface area

The BET surface area, BJH pore volume and average pore diameter for CSNC were determined by Brunauer–Emmett–Teller (BET) and BJH methods through N2 adsorption–desorption methods at 77.35 K. The data illustrated that the CSNC has surface area 9.945 m2/g, and this leads to an efficient transfer of the metal ions to the internal adsorption sites [49].

X-ray diffraction (XRD)

The XRD patterns of CHCF nanoparticles and CSNC are shown in Fig. 6. Figure 6a shows the diffraction patterns of CHCF nanoparticles, which showed a characteristic peak at 2θ values of 17.5, 24.9, 35.9, 40.1, 44.2 and 51.1° which are due to Miller indexes of (200), (220), (400), (420), (424) and (440), respectively [50], of the cubic crystal structure of CHCF (with a cell constant of 9.99 A; JCPDS card no. 02-0383). Figure 6b shows the XRD patterns of CSNC and the diffraction peaks in 2θ from 10° to 80°, which suggested that CSNC existed in semicrystalline structure. Figure 6 shows that all peaks of XRD were sharp peaks and there were not any broad peaks which showed that CHCF and CSNC had semicrystalline structure. The sharp peaks were observed at 2θ = 15, 19, 20, 21, 22 and 25°. The peak at 2θ = 25° was due to van der Waals distances between stacks of phenylene rings (polyanthranilic acid and poly-o-phenylenediamine rings) [51, 52]. The crystallinity was very important in the polymers as the more the crystalline system, the more the metallic conductive state that may influence the anticorrosion performance [53].

Fig. 6
figure 6

XRD patterns of a CHCF nanoparticles and b CSNC

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Thermogravimetric analysis

By thermogravimetric analysis, the thermal stability of the prepared CSNC was evaluated and is shown in Fig. 7. The four-stage decomposition pattern was recorded. The first weight loss step started from range 29.25 to 160.78 °C with range of weight loss of 11.109% which corresponds to the loss of water molecules, free acids and volatile molecules in the CSNC. The second step was with weight loss of 25.167% at temperature range 160.78–294.21 °C which represents loss of dopant bound to the nanocomposite chain and evolution of CO2. The third step was with weight loss of 17.443% at temperature range 294.21–503.73 °C, which was because of loss of sublimation and removal of low molecular weight polymer from the CSNC matrix [33]. The fourth step was with weight loss of 1.792% at temperature range 504.77–799.06 °C which was attributed for the complete degradation and decomposition of the nanocomposite backbone. Above 799.06 °C, the results obtained were associated with the residues only.

Fig. 7
figure 7

TGA curve of the prepared CSNC

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Adsorption of cesium ions on core–shell nanocomposite (CSNC)

Effect of pH

One of the important factors affecting the adsorption performance of CSNC toward cesium ion is the pH of the aquatic solutions. The removal behavior was studied at pH in range of 3–11 to obtain the optimum pH, which recorded the highest capacity of Cs+. Figure 8 shows the capacity of Cs+ at different pH values. Figure 8 shows that the increase in the pH of aquatic solutions showed an increase in the adsorption capacity of Cs ions onto CSNC. The highest adsorption capacity was 0.993 mmol g−1 at pH 11; thus, all the following experiments in this study were performed at pH 11. Also, we observed that the relative low adsorption at pH 3, 4 and 5 may be due to one of the two reasons or both of them. The first reason is the competitive adsorption between Cs+ and the H+ ions from carboxyl groups [54]. The second reason is the dissolution of a fraction of CHCF/poly (AA-co-OPD) surface under acidic conditions [55]. Also, we observed that the adsorption of Cs+ onto CSNC was relatively higher than other resins in high acidic conditions.

Fig. 8
figure 8

Effect of pH on the uptake of Cs+ ions; 100 ml (3 mmol L−1), 0.1 g CSNC, contact time 2 h, shaking rate 250 rpm and 25 °C

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Effect of initial Cs+ concentration and equilibrium isotherm models

To study the influence of change in concentration of cesium on sorption process, the sorption processes were carried out using different concentrations of Cs+ from 2 to 10 mmol L−1 at 25 °C and pH = 11. The obtained data explained that the uptake of Cs+ increased with the increase in the starting concentration of Cs+ to reach maximum capacity at initial concentration equal to 10 mmol L−1 of 1.35 mmol g−1 (Fig. 9).

Fig. 9
figure 9

Effect of initial Cs+ concentration on the sorption process; 100 ml of Cs, 0.1 g CSNC, pH 11, contact time 2 h, shaking rate 250 rpm and 25 °C

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By studying the adsorption isotherm, we can illustrate the relationship between the amount of adsorbate on the adsorbent and the concentration of dissolved adsorbate in the aquatic solution at equilibrium. So, four models were studied to the mechanism of sorption process of Cs+ onto CSNC (Fig. 10). The first model is Langmuir, which can be expressed through Eq. (2), and the second model is Freundlich (Eq. 3). Temkin (Eq. 4) is the third model, and Dubinin–Radushkevich (D–R) (Eq. 5) is the last one.

Fig. 10
figure 10

Experimental adsorption isotherm a Langmuir adsorption isotherm, b Freundlich adsorption isotherm, c Temkin adsorption isotherm and d Dubinin–Radushkevich (D–R) adsorption isotherm of Cs+ ions on the CSNC at room temperature and pH 11

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The four isotherm models are listed in Table 1. Table 1 shows that the highest value of linear correlation coefficient (R2 = 0.9811) was obtained from Langmuir model which fitted with the collected data better than other isotherm models. So, we concluded the sorption of Cs+ occurs largely on homogenous surface by monolayer adsorption process. The obtained RL values lie between 0.17 and 0.50 (Table 1), which showed that the equilibrium adsorption is favorable at higher initial concentration. As shown in Table 1, the lowest value of R2 = 0.8003 for Dubinin–Radushkevich (D–R) isotherm indicated that this model not fitted with adsorption process in this work.

Table 1 Parameters of Langmuir, Freundlich, Temkin and D–R isotherms for the adsorption of cesium ions on the CSNC
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In Table 2, the other common adsorbents used for the removal of cesium (adsorption capacity) are listed and compared with the synthesized CSNC of this work. This table shows that the synthesized CSNC of this work recorded also the highest adsorption capacity as some literature is listed in Table 2. These collected data reveal that the CHCF/poly(AA-co-OPD) nanocomposite is an effective adsorbent for the removal of Cs+ from wastewater and the possibility of this material for treatment of large volumes of radioactive contaminated water. Moreover, CHCF/poly(AA-co-OPD) can be prepared by simple and inexpensive materials.

Table 2 Cesium adsorption capacities from other composites listed in recently published articles related to of metal hexacyanoferrates
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Effect of contact time and equilibrium kinetic models

To examine the kinetics of adsorption process, a series of experiments were carried out to determine the optimal shaking time for uptake of Cs+ onto CSNC in the shaking time range of 0.5–3.5 h. Figure 11 shows that the adsorption of Cs+ ions from aqueous solution using the nanocomposite adsorbent is continuously increased with the increase in time until an equilibrium is reached. Therefore, this obtained equilibrium time was selected for the next adsorption experiments. At the beginning of adsorption process, the rate of uptake was rapid, and the equilibration was achieved in 3 h. The highest initial adsorption rate toward Cs+ indicates that the adsorption process occurs mainly on the nanocomposite surface and the regular distribution of CHCF inside poly(AA-co-OPD), making the rapid Cs+ uptake at the beginning of the experiment. The fast removal process decreased the time needed for the removal of Cs+ from wastewater which reduced the costs for the operation. Pseudo-first-order (Eq. 9), Pseudo-second-order (Eq. 10), intra-particle diffusion model (Eq. 11) and Elovich equation (Eq. 12) are four models, which are applied in this work in order to evaluate the mechanism of the kinetics of Cs+ onto the CSNC. The kinetic parameters and correlation factor (R2) were estimated and are listed in Table 3. The results reveal that the pseudo-second-order model (R2 = 0.959) is more suitable than other models. The fitness of pseudo-second-order kinetic model may mean that Cs+ adsorbed on CSNC involving chemical reaction, which includes ion exchange, substitution or complexation [56] (Fig. 12).

Fig. 11
figure 11

Effect of contact time on the sorption process; 100 ml of Cs, 0.1 g CSNC, pH 11, 10 mol L−1, shaking rate 250 rpm and 25 °C

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Table 3 Parameters of pseudo-first-order, pseudo-second-order, intra-particle diffusion model and Elovich models for the adsorption of cesium ions on the CSNC
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Fig. 12
figure 12

a Pseudo-first-order, b pseudo-second-order, c intra-particle diffusion model and Elovich model

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Influence of temperature and thermodynamic parameters

The uptake of Cs+ was studied with a temperature from 25 to 60 °C. The results in Fig. 13 explained that the adsorption efficiency increases as the temperature was raised, which is due to the endothermic nature of the process. The increase in the temperature helps to bind Cs+ to adsorption sites. To better evaluate the adsorption process, three thermodynamic parameters including Gibbs free energy (ΔG), enthalpy change (ΔH) and entropy change (ΔS) were used and are listed in Table 4. The calculated value of ΔH is 0. 0.401 kJ mol−1 > 0 indicated that the adsorption process of Cs+ on CSNC is endothermic reaction process. That is, the increase in the solution temperature can promote the adsorption of Cs+. Also, the values of (ΔG◦ads) were negative, which indicates the spontaneity and feasibility of Cs+ adsorption. The positive value of standard entropy (ΔS°ads) indicates the affinity of the prepared nanocomposite for Cs+ and the increasing randomness at the nanocomposite/Cs+ solution interface during the uptake process. It can be observed that the removal efficiency was not significantly influenced due to the change of temperature. The marginal drop in Cs removal performance with increasing temperature may be resulted in the thermal destabilization which increased the mobility of Cs ions and inaugurated the desorption steps [57, 58].

Fig. 13
figure 13

Influence of the increase in the temperature on the distribution coefficient of Cs+ ions onto CSNC ion exchanger

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Table 4 Thermodynamic parameters for Cs+ sorption on the synthesized CSNC ion exchanger
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Desorption process of cesium ions from CSNC

From the uptake elution, we observed that the capacity of newly synthesized CSNC was hardly affected even after repeated five regeneration cycles and the adsorption capacity decreased from 100 to 99%, 98%, 97%, 95% and 93%, respectively, over five cycles as shown in Table 5.

Table 5 Desorption capacity on recycling CSNC
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Conclusion

A low-cost and facile route has been developed for the preparation of CSNC derived from CHCF and poly(AA-co-OPD). The synthesis route included (i) preparation of nanoparticles of CHCF and (ii) plantation of the nanoparticles in polymer matrix of poly(AA-co-OPD) during the copolymerization process. The structure of the prepared CSNC was confirmed via FT-IR, XRD, TGA, TEM and SEM–EDX mapping. The SEM and TEM confirmed the nanostructure of the prepared CHCF and CSNC. Also, XRD analysis confirmed the expansion of the CHCF into poly(AA-co-OPD) and the synthesized CSNC used for the sorption of cesium ions from aqueous solutions by ion exchange process between the potassium ions in CSNC replaced by cesium ions. FT-IR analysis confirmed the presence of CN bonds in the CSNC adsorbents, implying the successful plantation of CHCF on the copolymer, which was further supported by EDX mapping.

The sorption process of the synthesized CSNC toward cesium ions and the influence of parameters on the sorption of Cs+ from wastewater were investigated as a function in pH, metal ion concentration, shaking time and temperature. The obtained data showed that the maximum capacity of CSNC toward Cs+ ions was 1.35 mmol g−1 at pH 11, 10 mmol L−1 Cs+ and 25 °C. Four models including Langmuir, Freundlich, Temkin and D–R isotherms models were studied, in which the data were well fitted with Langmuir model, suggesting that the uptake of Cs+ was monolayer and homogeneous. Also, the sorption kinetics results were fitted well to pseudo-second-order model. Thermodynamic parameters were calculated in the temperature from 25 to 60 °C and the data revealed that Cs+ sorption was endothermic, spontaneous and more favorable at higher temperature. Up to 92% desorption of Cs+ was completed with 2 M KCl.