All chemicals were analytical grade, purchased and used as received without further purification: Magnesium nitrate Mg(NO₃)₂∙6H₂O Merck; Ferric nitrate Fe(NO₃)₂∙6H₂O, Merck; Urea NH₂CONH₂; Fluka. Oxalic acid C₂H₂O₄∙2H₂O; Sigma-Aldrich; Citric acid HOC(COOH)(CH₂COOH)₂ Sigma-Aldrich; Ammonium hydroxide 25% NH₃ in H₂O, Sigma-Aldrich and T 99 cO 4 − radioactive waste was extracted from previously used ^99mTc generator /columns collected from the nuclear medicine centers using diluted nitric acid.

Magnesium ferrite, MgFe₂O₄, nanostructures were synthesized via sol-gel auto combustion method [21] [22] [23] in which an aqueous solution of magnesium nitrate (Mg(NO₃)₂∙6H₂O) was added to a stirring aqueous solution of ferric nitrate (Fe(NO₃)₂∙6H₂O) and the reaction was heated up at 60˚C, and magnetically stirred for 10 min to the hot stirring solution an aqueous solution of Urea was added with molarity ratios between metal nitrate and the fuels (1:2). The solution was heated at 80˚C and stirred for one hour until the gel obtained. The produced solution was dried at 200˚C for 2 h. The produced foamy powder was ground and then, calcined at various temperatures 600˚C and 800˚C to give MgFe₂O₄ nanostructures referred as a₆₀₀ and A₈₀₀, respectively. The produced magnesium ferrite samples (B and C) were prepared by applying similar conditions using Oxalic and Citric acids as a fuel, and the calcined products at 600˚C and 800˚C referred as (b₆₀₀, B₈₀₀) and (c₆₀₀, C₈₀₀) respectively. The synthesized samples were characterized for crystal phase identification by Powder X-Ray Diffraction (XRD) spectroscopy. The morphology was determined by the Field Emission Scanning Electron Microscopy (FE-SEM) and Fourier Transform Infrared (FTIR) spectroscopy are recorded using KBr pellets on a BRUKER® Tensor 27 FT-IR spectrometer, in the range of 4000 - 400 cm⁻¹.

Batch experiments were carried out to examine the sorption of pertechnetate, T 99 cO 4 − from aqueous solution by the synthesized Nano MgFe₂O₄. For the sorption determination, a certain weight (50 mg) of the investigated material (MgFe₂O₄) contacted with 10 ml of solution contain T 99 cO 4 − at different activity concentration from 2.13 to 15.58 Bq/vial for 24 hours. The sample suspensions were shaken at room temperature (25˚C), then centrifuged at 4000 rpm for 10 minutes. The activity concentration of the obtained solution was determined by Liquid Scintillation Counter (LSC) (type: Packard, USA, model 3720) by mixing 5 ml of the supernatant with 5 ml of ultima gold liquid scintillation cocktail and feeding the measuring vials to the Liquid Scintillation Counter (LSC). The adsorption capacity of the adsorbent (q_t, Bq/g); was calculated from Equation (1) [24], where, C₀ is the initial T 99 cO 4 − concentration, C_t is the T 99 cO 4 − concentration in solution at a pre-defined time t, V (L) is the volume of the T 99 cO 4 − solution, and m (g) is the mass of MgFe₂O₄ adsorbent. Moreover, Equation (2) [24] was employed to estimate the removal percent (%R) of the T 99 cO 4 − .

q t = ( C 0 − C t ) × V m (1)

T 99 cO 4 − R % = C 0 − C t C 0 × 100 (2)

Various factors have been investigated such as contact time (20 - 120 min), pH (2 - 9), competing cation concentration KCl (0.05 - 0.55 g), temperature (25˚C - 55˚C), and initial T 99 cO 4 − activity concentration of 951 CPM/vial (3.17 Bq/ml). The equilibrium adsorption capacity of MgFe₂O₄ adsorbent, q_e (Bq/g) can be determined from Equation (3).

q e = ( C 0 − C e ) × V m (3)

where: C_e (Bq/l) is the T 99 cO 4 − concentration at equilibrium in the supernatant after separation of the adsorbent and V, m, and C₀ have the same meaning as in Equations (1) and (2).

X-Ray Diffraction

The XRD patterns of magnesium ferrite samples produced by calcination of the combusted precursors at 600 and 800˚C are shown in Figures 1(a)-(c) and Figures 2(a)-(c), respectively. It is clearly observed that the calcination temperature 600˚C was not enough to produce a pure phase of MgFe₂O₄ with samples (a) and (b) as in Figure 1 which prepared using Urea and Oxalic acid as fuels respectively. The sample (a) showed a characteristic peak of hematite at 2θ = 35.48 corresponding to (JCPDS File No. 79-1741). The pattern sample (b) has some intermediate products of Magnesium oxide and hematite that evidenced from their reflection patterns at (2θ = 78.74, 2θ = 74.70) indexed to (JCPDS File No. 87-0652) and at (2θ = 33.11, 2θ = 49.41) indexed to (JCPDS File No. 89-0598), respectively. A similar approach is used for sample (c) Figure 1 revealed the production of single phase of magnesium ferrite with a relatively low amount of impurities.

The diffraction peaks of samples (a) and (c) were consistent very well with the cubic MgFe₂O₄ with cell parameters; a = 8.386 Å and V_cell = 589.81 ų (JCPDS File No. 088-1937). But for sample (b), when oxalic acid is used as a fuel, the diffraction peaks of sample (b) was indexed well with the cubic MgFe₂O₄ with cell parameters; a = 8.380 Å and V_cell = 588.61 ų (JCPDS File No. 088-1935). The same results of getting impure phase with urea fuel and obtaining pure phase with citric acid fuel was reported by Druc et al. in 2013 [25]. On the other hand, the calcination temperature for the three samples was raised to 800˚C to investigate its effect on the phase composition of the combustion products of the three fuels. Figures 2(a)-(c) shows the XRD patterns of the produced magnesium ferrite samples by calcination of the ignited products at 800˚C, it is obvious from Figure 2 that the calcination temperature at 800˚C was enough to produce a crystalline single product with the three fuels and all the diffraction peaks in Figure 2 are consistent well the cubic MgFe₂O₄ (JCPDS File No. 088-1937). Moreover, the crystallite size (D, nm) of magnesium ferrite products was calculated using the Debye-Scherer Equation (4) [26].

D = 0.9 λ β cos θ B (4)

where, λ (nm) is the X-ray radiation wavelength, β is the diffraction peak full width at half maximum (FWHM), and θB is the Bragg diffraction angle. The average crystallite size of the prepared MgFe₂O₄ nanoparticles calcined at 600˚C was about 10.41, 15.15 and 13.53 nm for samples a₆₀₀, b₆₀₀ and c₆₀₀ respectively. Additionally, the crystallite size was increased to 18.88, 18.56 and 15.22 nm on increasing the calcination temperature to 800˚C for samples A₈₀₀, B₈₀₀ and C₈₀₀ respectively. It can be concluded that, changing of the used fuel has a remarkable effect on both the particle size and the phase of the synthesized Nano-MgFe₂O₄, since using the citric acid fuel produced pure MgFe₂O₄ nanostructure with the smallest crystallite size at lower temperature (600˚C). Therefore, Citric acid is the optimum fuel in the present work.

FT-IR Analysis Nanoparticles

FT-IR spectroscopy is one of the most important and widely used analytical methods for obtaining information about the presence of certain functional groups in the structure of the magnetic nanomaterials. Figure 3 shows FT-IR spectra (A-C) for samples A₈₀₀, B₈₀₀, and C₈₀₀ prepared by Urea, Oxalic acid and Citric acid as a fuel respectively and annealed at 800˚C. The IR spectra of the samples showed a strong absorption band at 567, 563 and 568 cm⁻¹ for sample A, B, and C, respectively, which is characteristic and indicates the formation of MgFe₂O₄.

In ferrites crystal structure the metal ions have two characteristic different sub

lattices, known as tetrahedral and octahedral according to the geometrical configuration of the oxygen nearest neighbors. The metal-oxygen band at 567, 563 and 568 cm⁻¹ corresponds to the intrinsic stretching vibrations of the metal at the tetrahedral site [27]. Sample B₈₀₀ shows a weak absorption band at 1336 cm⁻¹ that is related to carbonate ions [27]. The band at wavenumber 1425 and 1421 for samples B₈₀₀ and C₈₀₀ respectively is a characteristic absorption peak for CO 3 2 − . However, the IR spectra also showed abroad vibration band at 3421 and 3434 cm⁻¹ for the B₈₀₀, and C₈₀₀ samples, respectively, that can be assigned to the OH stretching vibrations of water molecules (physical adsorbed molecular water) while their bending mode appeared at 1641and 1635 cm⁻¹ for the B₈₀₀, and C₈₀₀ respectively [28].

Field Emission Scanning Electron Microscopy (FE-SEM) of MgFe₂O₄ Nanoparticles

Magnesium ferrite MgFe₂O₄ nano-structures were observed using a scanning electron microscope are shown in Figures 4(a)-(c), which denoted as a₈₀₀, b₈₀₀ and c₈₀₀ referred to the used fuels Urea, Oxalic acid and Citric acid respectively in the preparation process. It is observed that the micrograph of sample (a₈₀₀) in Figure 4 shows aggregates of wool or cotton of the spherical with average crystallite size about 20 μm. On the other hand, Figure 4(b) & Figure 4(c) revealed that the products (b₈₀₀ and c₈₀₀) are composed of small single spheres, bulk agglomerates of sphere like structures and foamy structures with average size of 20 μm. The image illustrated that the spheres in the micrograph of sample b₈₀₀ are relatively smaller than that in sample c₈₀₀. The aggregates in the SEM images of the three samples could be attributed to the interaction between the nanoparticles due to the magnetic nature of the MgFe₂O₄ samples. In cases of nanocrystalline

spinel ferrites, it has been observed that there is a tendency for the nanoparticles to agglomerate as reported by Naseri et al., in 2014 [29]. From obtained results it can be concluded that the fuel type in the combustion process could affect the morphology of the final product.

Effect of solution pH

The influence of the initial pH of the solution on the removal percent of the investigated isotope ( T 99 cO 4 − ) was investigated in the pH range from 2 to 9 (Figure 5). The experimental data revealed that the removal percentage of T 99 cO 4 − is high at low pH values, it reaches the maximum removal% at pH ~2 (%R = 99.02%) followed by a plateaued over a wide pH range from pH 3 to 7 (%R = 98.69% - 98.59%), Afterward, the removal percentage reduces sharply for T 99 cO 4 − on increasing the pH values to 9. The behavior can be explained by considering the surface charge of the MgFe₂O₄ nanoparticles and the T 99 cO 4 − anions at different pH values. At lower pH values, the MgFe₂O₄ adsorbent surfaces (hydrous oxide (MOH)) will be probably covered by protons forming positively charged particles MOH⁺². However, at higher pH values, hydroxide ions are available in high concentrations and may react with the hydrous oxide forming negatively charged deprotonated oxide (MO⁻) [25] [30]. The adsorption process may be controlled by the electrostatic interactions between the negatively charged pertechnetate anions and the positively charged MgFe₂O₄ adsorbent. This behavior enhancing the adsorption process at lower pH values, but the adsorption efficiency percent decreased at higher pH values due to the competition between the high concentrations of OH⁻ ions and the negatively charged T 99 cO 4 − anions [31].

Effect of contact time

The influence of contact time on the adsorption efficiency of T 99 cO 4 − anions by MgFe₂O₄ nano-adsorbent was studied using 0.05 g MgFe₂O₄ Nano-adsorbent and initial concentration of 128 CPM/vial to 951 CPM/vial (0.43 Bq/ml to 3.17 Bq/ml) for T 99 cO 4 − . Figure 6 represents the T 99 cO 4 − removal percentage against shaking time ranged from 20 to 120 min. The sorption of the investigated isotope T 99 cO 4 − increases with time to reach a saturation level 98.84% for 100 min. depending on the nature of the isotope and selectivity of the prepared Nano adsorbent. The instantaneous removal percent of the investigated isotope by MgFe₂O₄ indicates that the mechanism of the sorption may be ion exchange [32].

Effect of temperature

The effect of different temperature (25˚C, 35˚C, 45˚C and 55˚C) on the adsorption of T 99 cO 4 − was investigated at initial concentration of 951 CPM/vial (3.17 Bq/ml) for T 99 cO 4 − .and 0.05 g from the nano-MgFe₂O₄ Figure 7, shows a relatively decrease in the removal percentages of T 99 cO 4 − on MgFe₂O₄ this may be attributed to the increasing mobility of T 99 cO 4 − ions as the temperature increases, causing small amounts of the ions escapes from the solid phase to the liquid phase [27]. This in turn, means that the reaction may be has an exothermic nature and the adsorption process is favored at low temperatures.

Effect of competing ion concentration

The effect of ionic strength on the adsorption of T 99 cO 4 − on MgFe₂O₄ was investigated utilizing KCl salt at different weight from 0.05 to 0.55 g. The results shown in Figure 8 reveal that, the solution ionic strength has a negative effect on T 99 cO 4 − ions removal percentage, as increasing the ionic strength decreases the T 99 cO 4 − ions adsorption percentage. The significant decrease of removal percentage may be due to the competition between the T 99 cO 4 − ions and chloride ions in the solution. In other words, as a result of the electrostatic attraction between the negatively charged T 99 cO 4 − ions and the positively charged MgFe₂O₄

adsorbent surface, the increase in the solution ionic strength will result in a reduction of the adsorption capacity and this agree with the investigators Ali et al., in 2016 [33] and Reddad et al., in 2002 [34].

Effect of initial T 99 cO 4 − activity concentration

T 99 cO 4 − activity concentration ranged from 128 CPM/vial to 951 CPM/vial (0.43 Bq/ml to 3.17 Bq/ml) and utilizing 0.05 g MgFe₂O₄ as adsorbent. The adsorption capacity increased with increasing the initial activity concentration for T 99 cO 4 − and it reached 186.4 CPM/g (3.11 Bq/g) corresponding to activity 951 CPM/vial (3.17 Bq/ml) (Figure 9). This behavior could be assigned to the higher driving force at more T 99 cO 4 − ions that ease the diffusion of the ions from the solution to the positively charged MgFe₂O₄ nanoparticles surfaces, and hence more interactions between T 99 cO 4 − ions and the active sites of MgFe₂O₄ sorbent [35].

Adsorption isotherm is the relationship at equilibrium between the quantity of adsorbate per unit of adsorbent (q_e) and its equilibrium solution concentration (C_e) at a constant temperature. To determine the adsorption capacity of the MgFe₂O₄ nanoparticles, the equilibrium data for the adsorption of T 99 cO 4 − on MgFe₂O₄ nanoparticles are analyzed in the light of Langmuir and Freundlich adsorption isotherm models [36] [37].

Langmuir isotherm model

Formation of a monolayer adsorbate on the outer surface of the adsorbent, and no further adsorption takes place again, is the assumption of Langmuir isotherm model. It is applicable for monolayer adsorption onto adsorbent surface that has a finite number of identical sites and no transmigration of adsorbate in the plane of the surface. According to these assumptions, the linear equation of Langmuir isotherm model can be written as follows [36].

C e q e = 1 Q 0 b + C e Q 0 (5)

where: C_e is the concentration of T 99 cO 4 − anions solution (CPM/l) at equilibrium, the constant Q₀ refers to the adsorption capacity (CPM/g) and b, Langmuir constant that is related to the energy of adsorption. Linear plot of C_e/q_e versus C_e shows Langmuir isotherm (Figure 10). Values of Q₀ and b were calculated from the slope and intercept of the linear plots and are presented in Table 1. Moreover, the equilibrium parameter R_L which is characteristic of Langmuir isotherm can be, defined as:

R L = 1 1 + b C 0 (6)

where: b is the Langmuir constant and C₀ is the initial T 99 cO 4 − concentration CPM/vial or (Bq/ml), R_L value indicates the adsorption nature to be either favorable if 0 < R_L < 1, linear if R_L = 1, unfavorable if R_L > 1 and irreversible if R_L = 0 [38]. Table 1 shows R_L values between zero and one, which indicate favorable adsorption process.

Freundlich isotherm model:

Freundlich isotherm model assumes that the adsorption takes place on a heterogeneous surface of the adsorbent and the linearized form of this model can be given by Equation (7) [37].

ln q e = ln K f + ( 1 / n ) ln C e (7)

where: K_f is the Freundlich constant (CPM/g) which represents the relative adsorption capacity of the adsorbent. (1/n) is the heterogeneity factor and it is a function of the strength of adsorption in the adsorption process and n has various values depending on the heterogeneity of the sorbent. If n lies between one and ten, this indicate a favorable sorption process [39]. 1/n and lnK_f values were calculated from the slope and intercept of the linear plots of lnq_e versus lnC_e which shows Freundlich isotherm (Figure 11) and are presented in Table 1.

From Table 1, it is observed that, the heterogeneity parameter, 1/n, for T 99 cO 4 − (1/n = 0.793, n = 1.261), which indicates a favorable sorption process as (n) lies between one and ten [39]. Moreover, the correlation coefficient (R²) values of Freundlich isotherm for T 99 cO 4 − is closer to unity and greater than that in Langmuir isotherm, So, we conclude that the adsorption of T 99 cO 4 − by MgFe₂O₄ as adsorbent fits Freundlich isotherm model and this in turn suggests that the adsorption process occurs as a multilayer T 99 cO 4 − molecules adsorb onto the heterogeneous adsorbent surface.

Pseudo first and second order reaction rate models were used to describe the adsorption behavior of the reaction at which the T 99 cO 4 − anions is removed from aqueous solutions. Pseudo first order rate equation of Lagergen is given by Equation (8) [40].

log ( q e − q t ) = log q e − ( K 1 , a d s / 2.303 ) t (8)

where: q_e and q_t are the amounts of T 99 cO 4 − anions adsorbed (CPM/g) at equilibrium and at time t (min), respectively and K_1,ads is the rate constant of pseudo first order adsorption (min⁻¹). Values of K_1,ads were calculated from the plots of log ( q e − q t ) versus t (Figure 12).The rate equation for pseudo-second order model is given by equation (9) [41].

t / q t = 1 / ( k 2 , a d s q e 2 ) + ( 1 / q e ) t (9)

where: k_2,ads (g/CPM min) is the pseudo second order rate constant and its value was obtained from the plots of t/q_t versus t (Figure 13). Table 2 lists the rate constants, calculated equilibrium adsorption capacity q_e(cal) and experimental equilibrium adsorption capacity q_e(exp) for T 99 cO 4 − obtained using the pseudo-first and second order models. It is obvious that q_e value calculated from pseudo first order kinetic model was too small compared to the experimental value for T 99 cO 4 − anions. However, the calculated q_e from pseudo second order

kinetic model was relatively close to the experimental value. Additionally, the correlation coefficient (R²) values of pseudo second order model is closer to unity. These results show that the pseudo second order adsorption mechanism is predominant and adsorption process of T 99 cO 4 − anions appeared to controlled by the chemisorption process. T 99 cO 4 − anions can be ionized in solution to form negative ions and the nano-MgFe₂O₄ carry a positive charge, which may lead to a chemical reaction through electrostatic interaction between a positive charge at the surface of MgFe₂O₄ nanoparticles and the negative charge on pertechnetate anions [42].

Van’t Hoff Equation (10), was used to calculate the thermodynamic parameters such as change in enthalpy (∆H^o), change in entropy (∆S^o) and change in free energy (∆G^o). Adsorption of T 99 cO 4 − on MgFe₂O₄, at different temperatures were determined by using Equations (11) & (12) and illustrated in Figure 14. The values of ∆H^o and ∆S^o were calculated from Figure 14 and reported in Table 3.

ln K d = Δ S ° R T − Δ H ° R T (10)

K d = C A e C e (11)

Δ G ° = Δ H ° − T Δ S ° (12)

where: K_d is the equilibrium constant, C_Ae is the solid phase concentration at equilibrium (mg/L), T is the temperature in kelvin and R is the gas constant. Plotting lnK_d against 1/T gives a straight line with slope and intercept equal to―∆H^o/R and ∆S^o/R, respectively. The values of ∆H^o and ∆S^o were calculated from Figure 14 and reported in Table 3. The negative values of ∆H^o indicate the exothermic nature of adsorption process. The negative value of entropy change (∆S^o) corresponds to a decrease in degree of freedom of the adsorbed species. This suggests the decrease in concentration of adsorbate in solid solution interface indicating an increase in adsorbate concentration on the solid phase. This is the

normal effect of chemical adsorption phenomenon, which takes place through electrostatic attraction. Gibbs free energy of adsorption (∆G^o) was calculated from Equation (12) and the values are given in Table 3. The negative value of ∆G^o indicates that the adsorption of T 99 cO 4 − on Nano-magnesium ferrite is a spontaneous process, whereby no energy input from outside of the system is required. However, the values of ∆G^o decreased with increasing of temperature suggesting that adsorption became less favorable at higher temperatures. As the temperature increases, the mobility of T 99 cO 4 − anions increases, causing the anions to escape from the solid phase to the liquid phase. Therefore, the amount of T 99 cO 4 − anions that can be adsorbed will decrease. The increased mobility of T 99 cO 4 − anions at elevated temperature may be reflected in the values of K_d (Table 3). As the temperature increased the values of K_d decreased indicating a lower affinity of the MgFe₂O₄ nanoparticles towards T 99 cO 4 − anions at higher temperatures [43] [44].