Magnificent adsorption capacity of hierarchical mesoporous copper oxide nanoflakes towards mercury and cadmium ions: Determination of analyte concentration by DPASV

Powder Technology 258 (2014) 11–19

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Magnificent adsorption capacity of hierarchical mesoporous copper oxide nanoflakes towards mercury and cadmium ions: Determination of analyte concentration by DPASV Kumarswamy Yogesh Kumar a, Handanahally Basavarajaiah Muralidhara a,⁎, Yenjerappa Arthoba Nayaka b a b

Centre for Emerging Technologies, Jain University, Bangalore 562112, India Department of PG Studies and Research in Chemistry, School of Chemical Sciences, Kuvempu University, Shankaraghatta 577451, India

a r t i c l e

i n f o

Article history: Received 27 September 2013 Received in revised form 14 February 2014 Accepted 1 March 2014 Available online 7 March 2014 Keywords: CuOs Adsorption Mercury Cadmium DPASV

a b s t r a c t Hierarchical mesoporous copper oxide nanoflakes (CuOs) were successfully synthesized by hydrothermal method. The complete characterization of product was done by XRD, SEM, TEM, SAED, EDX, BET and FT-IR studies. The batch experiments were conducted under different adsorbate concentration, contact time, pH and temperature conditions. The residual analyte concentration was determined using differential pulse anodic stripping voltammetry (DPASV) technique with high accuracy and reproducibility. Adsorption equilibrium was studied with Langmuir and Freundlich isotherm models. Equilibrium data were best fitted with the Langmuir and Freundlich isotherm models and the maximum adsorption capacities of Hg(II) and Cd(II) were determined to be 1767.97 and 1577.78 mg/g respectively. All these values are significantly higher than those reported on other hierarchical nanostructures. Thermodynamic parameters and adsorption kinetics were studied in detail to know the nature and mechanism of adsorption. A regeneration study is proposed in order to reuse the adsorbent for better economy of the process. The results demonstrate that CuOs can be used as a possible alternative low-cost adsorbent for the efficient removal of heavy metals from aqueous solutions. © 2014 Elsevier B.V. All rights reserved.

1. Introduction Water contaminated by heavy metal ions has become much more serious with the rapid development of industries and competitive use of fresh water in many parts of the world. The occurrence of heavy metals, especially mercury and cadmium, in industrial effluents beyond permissible level brings serious environmental pollution, threatening human health and ecosystem [1]. Therefore, heavy metal ion removal from water has become an important subject today [2,3]. Generally ion exchange, precipitation, adsorption, flocculation and membrane processes have been used for the removal of toxic metal ions. In particular, adsorption is recognized as an effective and economic method for the removal of pollutants from waste waters [4,5]. Hence, one of the ultimate goals for the pollutant removal is to design and fabricate highly efficient adsorbents with high adsorption capacity and rapid adsorption rate. Numerous adsorbents such as zeolites, clays, metal oxides and activated carbons were developed for the adsorption of heavy metal ions [6–9]. Metal oxide nanoparticles have revealed their vast application in the field of sensing, optoelectronics, catalysis, solar cells and so on, owing to their distinctive physical and chemical properties differing from the ⁎ Corresponding author. E-mail address: [email protected] (H.B. Muralidhara).

http://dx.doi.org/10.1016/j.powtec.2014.03.015 0032-5910/© 2014 Elsevier B.V. All rights reserved.

bulk [10,11]. Moreover, in recent years, considerable efforts have been made to design, synthesize and fabricate hierarchically structured mesoporous metal oxide nanoparticles for the adsorption of heavy metal ions [12–16]. Their novel structures possess a high surface to bulk ratio and a large surface area, contributing to high adsorption capacities and fast adsorption rates, as well as remarkably facilitating the mass transportation of the target. CuO has been extensively studied because it is an important component of copper oxide superconductors [17,18]. With regard to its commercial value and interesting properties, CuO has also been widely exploited in a versatile range of applications such as catalysts [19], magnetic storage media [20], field emission devices [21], gas sensors [22], lithium batteries [23], and solar cells [24]. Chemical and physical properties of this material depend on its composition, structure, phase, shape and size distribution. Thus, diverse morphologies have been obtained for the CuO with and without the use of surfactants or templates [25]. Many synthesis methods have been reported to prepare CuO nanocrystals with an aim to achieve tunable properties. Annealing, arc spray, and precipitation–pyrolysis methods are frequently used to produce CuO [26,27]. As is well known, the samples prepared by hightemperature calcinations and pyrolysis have the drawbacks of nonuniformity [28]. Alternatively, hydrothermal and solvothermal processes are powerful to synthesize nanocrystals with homogeneous particle

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

sizes [29,30]. However CuO nanocrystals by solvothermal conditions are always reduced to Cu2O or Cu by the solvents in forming unwanted impurity phases [31]. Therefore, a simple and green synthetic route for CuOs with single-crystalline is sorely needed. Herein, we demonstrate a hydrothermal method for synthesizing CuOs using low cost materials instead of toxic and dangerous reagents. The purpose of this study is to investigate the feasibility of CuOs as high efficient adsorbent for the removal of Hg(II) and Cd(II) from aqueous solutions. Operational conditions such as contact time, pH and adsorbent dosage were systematically studied. Adsorption kinetics, isotherms and thermodynamics were analyzed to reveal the adsorption mechanisms and to achieve a better understanding of the adsorption process.

(111) (200)

12000

10000

Intensity (a.u)

12

8000

6000 (202) (113)

4000

(110)

(311) (220)

(021) (020)

2. Experimental

(112)

(112)

2000

2.1. Materials All the reagents including copper(II) nitrate trihydrate (Cu(NO)3· 3H2O), sodium hydroxide (NaOH), Triton X-100 (C14H22O(C2H4O)n), cadmium nitrate (Cd(NO)3·4H2O), mercuric chloride (HgCl2) and potassium chloride (KCl) used in the experiments were of analytical grade purchased from Fisher Scientific India Pvt. Ltd., (Mumbai, India) and used without further purification. Double distilled water was used throughout the experiments. 2.2. Preparation of CuOs In a typical procedure, 20 ml of 0.2 mol/L NaOH solution was slowly added into a 20 ml of 0.1 mol/L copper(II) nitrate trihydrate solution containing 0.004 mol/L of Triton X-100 with constant stirring. After vigorous stirring for 3 h, the mixture was autoclaved at 200 °C for 5 h. After the reaction system was naturally cooled to room temperature, the precipitates was separated from solution and thoroughly washed several times first with deionized water and then absolute ethanol and later dried in an oven at 50 °C for 8 h. 2.3. Characterization techniques X-ray diffraction (XRD) patterns were obtained on a Bruker D2 Phaser XRD system. Surface morphology (SEM) was studied using scanning electron microscope (JEOL JSM 840A), coupled with energy dispersive X-ray analyzer (EDX). Transmission electron microscope (TEM) and selected area electron diffraction patterns (SAED) were recorded by using a Philips CM-200 instrument. BET surface area, total pore volume and average pore size of the CuOs were measured using ASAP 2010 Micrometrics instrument by the Brunauer–Emmett–Teller (BET) method. Finally, the Fourier transform infrared (FT-IR) analysis was applied to determine the surface functional groups, using FT-IR spectroscope (Bruker ATR). 2.4. Heavy metal ion adsorption experiments Solutions with different concentrations of Hg(II) and Cd(II) were prepared using HgCl2 and Cd(NO)3·4H2O as the sources of heavy metal ions respectively. To study the effect of parameters such as initial concentration (100–400 mg/L), contact time (0–90 min), solution pH (2.0–8.0) and temperature (30–50 °C) for the removal of Hg(II) and Cd(II) on CuOs, experiments were conducted in batch method, in a 250 ml of stoppered flasks (Erlenmeyer flasks) that contained 100 ml of metal ion solution of varying concentrations, an equal amount of CuOs (250 mg/L) was added into each flask and subjected to agitation using an incubator shaker at 200 rpm, until equilibrium was reached (90 min). The resultant solutions were centrifuged and the supernatant liquids were determined for Hg(II) and Cd(II) ions. The reaction was carried at room temperature and pH of the solution maintained was

30

40

50

60

70

2θ Fig. 1. Typical XRD pattern of CuOs.

unaltered unless otherwise mentioned. The metal ion concentration was analyzed by DPASV using electrochemical workstation (CHI 660D). All experiments were conducted in a three electrode electrochemical cell with glassy carbon (Ø = 3 mm) as working electrode, saturated calomel as reference electrode and platinum wire as auxillary electrode. The amount of adsorbate adsorbed at equilibrium condition, qe (mg/g) was calculated by the following Eq. (1): qe ¼

ðC 0 −C e Þ V W

ð1Þ

where C0 and Ce are the initial and equilibrium concentrations (mg/L) respectively. V is the volume of solution (L) and W is the mass of adsorbent used (g). The effect of pH on the adsorption capacity of metal ion was evaluated by agitating 200 mg/L metal ion solution with 250 mg/L of CuOs for predetermined equilibrium time at pH ranging from 2.0 to 8.0. The pH of metal ion solution was adjusted by using 0.5 M HCl or 0.5 M NaOH. Similarly the rate of Hg(II) and Cd(II) adsorption onto CuOs was investigated by kinetics and isotherm study. Duplicate experiments were carried out for all operating variables studied and only average values were taken into consideration. 3. Results and discussion 3.1. Characterization of adsorbent Fig. 1 shows the XRD pattern of as-prepared CuOs obtained by the hydrothermal process, where all diffractions can be indexed as monoclinic phased CuOs by comparison with JCPDS card file nos. 80–1916 (a = 4.69 Å, b = 3.42 Å and c = 5.13 Å). No characteristic peaks of other impurities were detected. The diffraction peaks were very sharp, indicating that CuOs obtained are well-crystallized. The average particle size and lattice parameters are summarized in the Table 1. The SEM study of the as synthesized CuOs reveals the surface topography. The typical SEM image of the CuOs is shown in Fig. 2 a–b. The low-magnification image (Fig. 2a) indicates that the panoramic Table 1 Parameters derived from XRD of CuOs. Adsorbent

2θ

hkl

Size

Structure

CuOs

32.47 35.49 38.89

(110) (111) (200)

12.5 6.61 5.90

Monoclinic

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

(a)

13

(b)

Fig. 2. SEM micrographs of CuOs: (a) low and, (b) high magnified images.

morphology of the as-prepared sample is mainly composed of uniform flake like structures. The clear view of Fig. 2b displays that the surface of the flake like architectures is not smooth. They are densely packed possessing compact texture. High resolution TEM (HRTEM) (Fig. 3a) observations were conducted for these nanoflakes, and SAED (Fig. 3b) were collected on a single nanoflake; it has a bunch of spindle shape and its length was 41.28 nm. The color contrast of this CuOs is homogeneous, which indicates that it has a uniform dense structure and its surface is smooth [32]. This observation is in accordance with the corresponding SEM image. The specific surface area (SSA) is one of the important parameters for adsorbent. Fig. 4 shows the BET measurement curve of CuOs. Their BET SSA was determined to be 9.72 m2/g and their total pore volume is 0.023 cm3/g. Majority of the pore sizes are in the range of 9.35 nm. HRTEM image of the CuOs shows that the size of different color contrast regions is in the similar size range. Thus, the HRTEM observation agrees well with the pore size distribution measurement. Thus, these CuOs synthesized by the hydrothermal process have a relatively large surface area and pore volume, which are beneficial for its sorption capability.

3.2. Growth mechanism of CuOs Based on the experimental results, the possible formation mechanism of the CuOs was proposed. The morphologies of material are mainly determined by: the crystal nucleation and crystal growth direction. A large quantity of the growth units ([Cu(OH)4]2−) were formed under the condition of heavy alkaline solution like NaOH, which contribute

(a)

to the growth of CuOs nuclei and the crystal nuclei exhibiting fast growth orientation as shown in Fig. 5. 3.3. Effect of initial concentration Fig. 6 summarizes the effect of initial concentration, when the initial Hg(II) and Cd(II) concentrations were 100 to 400 mg/L and the loading concentration of CuOs was 250 mg/L. With the increase of contact time, the remaining metal ion concentration in the water samples decreased rapidly. For example, when the loading concentration of CuOs was 250 mg/L and initial concentration was 200 mg/L, CuOs can remove approximately 75 to 80% of Hg(II) and Cd(II) from water in just 30 min. It was evident that for lower initial concentrations of Hg(II) and Cd(II), adsorption was very fast. The removal efficiency decreased with increase in initial concentration and took longer time to reach equilibrium. With increase in concentration, competition for the active adsorption sites increases and the adsorption process will increasingly slow down. 3.4. Effect of pH Hg(II) and Cd(II) removal by the adsorbent increased with the increase in pH from 2 to 7, and was maximum at a pH range of 6.0–7.0 as shown in Fig. 7. Removal percentage increased from 24.19 to 98.81% for Hg(II) and from 24.63 to 97.08% for Cd(II) with the increase in pH from 2.0 to 7.0. In solutions having low pH value, there would be competition between protons and metal ions in the adsorption on surface hydroxyl groups, resulting in low metal adsorption. Moreover, when the active sites are protonated, the surface becomes positively charged and consecutive adsorption of metal ions in solution on the

(b)

Fig. 3. (a) Typical TEM micrograph of CuOs, (b) SAED patterns of CuOs.

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K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

40

Removal efficiency in %

35

Volume Adsorbed (cm³/g)

100

Adsorption Desorption

30 25 20 15

80

60 100 ppm Hg2+ 200 ppm Hg2+ 300 ppm Hg2+ 400 ppm Hg2+ 100 ppm Cd2+ 200 ppm Cd2+ 300 ppm Cd2+ 400 ppm Cd2+

40

20

Hg2+ - - - - Cd2+

10 0 5 0

0 0.2

0.4

0.6

0.8

20

40

60

80

100

Time in minutes

1.0

Relative Pressure (P/P0)

Fig. 6. Effect of contact time on adsorption of Hg(II) and Cd(II) by CuOs at different initial concentrations.

Fig. 4. Nitrogen adsorption–desorption isotherms of CuOs.

surface is less likely to occur. On the other hand, the active sites are deprotonated when the solution pH is higher, resulting in negatively charged sites and the adsorption of metal cations on CuOs could possibly take place via electrostatic interaction with the negatively charged sites on the surface [33]. As a matter of fact, at higher pH the determination of reliable adsorption capacity is not possible, due to the possible precipitation of cations as hydroxides. 3.5. Adsorption isotherm studies The adsorption capacity of these CuOs on Hg(II) and Cd(II) was investigated by the equilibrium adsorption isotherm study. In this research, adsorption isotherm study was carried out on two well known isotherms, i.e., Langmuir and Freundlich, and the parameters of both models were calculated and listed in Table 2. The equilibrium adsorption data could be best fitted with the Langmuir isotherm as follows:  qe ¼

KLCe 1 þ aL C e

in solution at equilibrium. The constant KL (L/mg) is the Langmuir equilibrium constant and the KL/aL gives the theoretical monolayer saturation capacity, Q0. Therefore, a plot of Ce/qe versus Ce gives a straight line of slope aL/KL and intercepts 1/KL. The square of the correlation coefficient (R2) values indicate that the Langmuir isotherm fitted the experimental data well (Fig. 8). The maximum adsorption capacities are 1767.97 mg/g and 1577.78 mg/g for Hg(II) and Cd(II) respectively. These two values are likely the highest among the adsorption data reported in the literatures and are significantly higher than those of many reported hierarchical nanostructures. Essential characteristics of the Langmuir isotherm parameters can be used to predict the affinity between the adsorbate and adsorbent using separation factor or dimensionless equilibrium parameter, “RL”, expressed as in the following equation:

RL ¼

1 : 1 þ aL C e

ð3Þ

 ð2Þ

where qe (mg/g) and Ce (mg/L) are the amounts of adsorbed adsorbate per unit weight of adsorbent and unadsorbed adsorbate concentration

The values of RL indicate the shapes of isotherms to be either unfavorable (RL N 1), linear (RL = 1), favorable (0 b RL b 1) or irreversible (RL = 0). RL was found to be in the range of 0–1 which indicates the favorable adsorption.

Random aggregation

CuO nuclei

CuO flake arrays Fig. 5. Schematic representation of growth mechanism of CuOs.

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

100

80

Ce/Qe (mg/g)

Removal efficiency in %

90

70 60 50 40

Hg2+ - - - - Cd2+

30

200 ppm Hg2+ 200 ppm Cd2+

20

15

0.032 0.030 0.028 0.026 0.024 0.022 0.020 0.018 0.016 0.014 0.012 0.010 0.008 0.006 0.004 0.002 0.000

Hg2+ - - - - Cd2+ 0

2

3

4

5

6

7

10

20

30

40

50

Ce(mg/L)

8

Variation of pH Fig. 8. Langmuir adsorption isotherms for Hg(II) and Cd(II) on CuOs. Fig. 7. Variation in removal efficiency of Hg(II) and Cd(II) on CuOs as a function of pH.

3.2

The Freundlich isotherm is suitable for a highly heterogeneous surface and expressed by the following: 1 log C e þ log K f n

3.0

ð4Þ

where qe (mg/g) is the equilibrium value for removal of adsorbate per unit weight of adsorbent, Ce (mg/L) is the equilibrium concentration of metal ion in solution, Kf and n are Freundlich isotherm constants which are related to the adsorption capacity (or the bonding energy) and intensity of the sorbent respectively. From the correlation coefficients, it can be seen that the adsorption data for Hg(II) and Cd(II) onto CuOs fit the Freundlich isotherm model (Fig. 9) better than the Langmuir isotherm model (R2 = 0.99). The values of 1/n less than 1 represent a favorable adsorption.

log Qe

log qe ¼

3.1

2.9 2.8 2.7

Hg2+ - - - - Cd2+

2.6 2.5 -0.4 -0.2

0.0

0.2

0.4

3.6. Kinetic studies

1.0

1.2

1.4

1.6

1.8

Fig. 9. Freundlich adsorption isotherms for Hg(II) and Cd(II) on CuOs.

3.0

2.5

log (qe-qt)

2.0 100 ppm Hg2+ 200 ppm Hg2+ 300 ppm Hg2+ 400 ppm Hg2+ 100 ppm Cd2+ 200 ppm Cd2+ 300 ppm Cd2+ 400 ppm Cd2+

1.5

1.0

0.5 Table 2 Adsorption isotherm model parameters of adsorption of Hg(II) and Cd(II) on CuOs.

Hg(II) Cd(II)

0.8

log Ce

The kinetics of adsorption, which describes the solute uptake rate governing the residence time of adsorption reaction, is one of the most important characteristics that define the efficiency of adsorption. It is clear from figures (Figs. 10 and 11) that the adsorption rate of Hg(II) and Cd(II) is rather fast. The adsorption is rapid at first and then slows down considerably. The probable adsorption model may involve metal ions that are initially adsorbed by CuOs on their exterior surface. When the adsorption on the exterior surface reaches saturation level, the metal ions begin to enter the exterior via the pores and move into the interior of CuOs. When metal ions diffuse into the interior of the CuOs, the diffusion resistance is increased; this in turn leads to a decrease in diffusion rate [34]. In such experimental conditions, most of Hg(II) and Cd(II) ions could be removed after 90 min. In addition, the above adsorption kinetic experimental data can be best fitted into a

Metal ion

0.6

Langmuir

Hg2+ - - - - Cd2+

0.0

Freundlich

Q0 (mg/g)

KL (L/mg)

R2

Kf (mg/g)

n

R2

1767.97 1577.78

0.17 0.13

0.95 0.98

470.78 376.85

2.52 2.93

0.99 0.99

0

10

20

30

40

50

60

70

80

90

Time in minutes Fig. 10. Pseudo-first-order kinetic plots for the adsorption of Hg(II) and Cd(II) on CuOs.

16

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

1600

0.25 100 ppm Pb2+ 200 ppm Pb2+ 300 ppm Pb2+ 400 ppm Pb2+ 100 ppm Cd2+ 200 ppm Cd2+ 300 ppm Cd2+ 400 ppm Cd2+

0.15

Hg - - - - Cd2+

1400 1200

1000

Hg2+ - - - Cd2+

Qt

t/qt (min / mg/g)

0.20

100 ppm Hg2+ 200 ppm Hg2+ 300 ppm Hg2+ 400 ppm Hg2+ 100 ppm Cd2+ 200 ppm Cd2+ 2+ 300 ppm Cd 400 ppm Cd2+

2+

800

0.10

600 0.05 400 0.00 0

20

40

60

80

200

100

3.0 3.5 4.0 4.5 5.0 5.5 6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5

Time (min)

t0.5

Fig. 11. Pseudo-second-order kinetic plots for the adsorption of Hg(II) and Cd(II) on CuOs.

Fig. 12. Intraparticle diffusion model fitting of the adsorption kinetics Hg(II) and Cd(II) on CuOs.

pseudo-second-order rate kinetic model. The mathematical expressions of the first order and second order kinetic model are given by: k t logðqe −qt Þ ¼ logqe − 1 2:303

ð5Þ

t 1 1 ¼ þ t qt k2 q2e qe

ð6Þ

where qe and qt are the amounts of adsorbate adsorbed (mg/g) at equilibrium at contact time t (min) respectively, and k1, k2 are the rate constants (1/min) of first order and second order kinetics. The linearized first and second order kinetics of Hg(II) and Cd(II) onto CuOs are presented in Figs. 10 and 11 respectively. The parameters of both models were calculated and listed in Table 3. The pseudo-first-order model showed poor fitting to the experimental data with very low correlation coefficients, i.e., ranging from 0.77 to 0.87 for Hg(II) and from 0.77 to 0.91 for Cd(II). The calculated equilibrium adsorption capacities (qe, cal) deviated too much from the measured values (qe, exp). This suggested that the pseudo-first-order model failed to describe the adsorption process correctly. But experimental data could be well fitted with the linear form of the pseudo-second-order model. The correlation coefficient value for pseudo-second-order model is 0.99, suggesting that the pseudo-second-order model best represents the adsorption kinetics in our adsorbent systems. And, we also found that the value of calculated qe was close to that of experimental qe. The intraparticle diffusion model proposed by Weber and Morris is applied to study the further adsorption process, which is written as: Q t ¼ kid t

1=2

þ C

ð7Þ

where Kid is the intra-particle diffusion rate constant (mg/g min1/2) and C is the intercept. The overall reaction kinetics for the adsorption of Hg(II) and Cd(II) is a pseudo-second-order process. However, this could not highlight the rate-limiting step. The rate-limiting step (slowest step of the reaction) may be either the boundary layer (film) or the intra-particle (pore) diffusion of solute on the solid surface from bulk of the solution in a batch process. It can be seen from Fig. 12 that the plots possess multi-linear portions; it indicates that two or more steps influence the sorption process. It was found that three straight lines relate the points, the first linear portion is due to the film diffusion and the second linear portion is due to the pore diffusion. Non-linearity of the plots had indicated the multi stage adsorption of Hg(II) and Cd(II) by CuOs. The extrapolation of the first linear portion gives the intercept equal to the boundary layer thickness or film diffusion. 3.7. Thermodynamics studies To study the effect of temperature parameter on the uptake of Hg(II) and Cd(II) onto CuOs, temperatures of 30, 40 and 50 °C were selected. Thermodynamic parameters, free energy (ΔG°), enthalpy (ΔH°) and entropy (ΔS°) changes were also calculated using the following equation: kC ¼

Ca Cb

ð8Þ

0

ΔG ¼ −RT ln kC

ð9Þ

Table 3 Kinetic parameters for the adsorption of Hg (II) and Cd(II) on CuOs. Hg(II) concentration (mg/L)

Cd(II) concentration (mg/L)

Kinetic models

Parameters

100

200

300

400

100

200

300

400

First order kinetic model R2 qe, exp (mg/g) Second order kinetic model R2 qe, exp (mg/g)

qe,cal (mg/g) k1 (1/min) 0.87 397 qe, cal (mg/g) k2 (1/min) 0.99 397

439 0.06 0.87 782 438 0.10 0.99 782

682 0.05 0.77 1161 869 0.10 0.99 1161

1214 0.05 0.77 1528 1273 0.10 0.99 1528

1331 0.04 0.84 393 1739 0.07 0.99 393

325 0.05 0.77 769 427 0.11 0.99 769

760 0.05 0.86 1094 833 0.12 0.99 1094

805 0.04 0.91 1429 1183 0.12 0.99 1429

1183 0.05

1579 0.11

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

17

Δs0 ΔH 0 − R RT

ð10Þ

ln kC ¼

5.6

lnKC

5.4

5.2

5.0

Hg2+ - - - - Cd2+

4.8

4.6

3.8. EDX analysis 0.00300

0.00305

0.00310

0.00315

0.00320

1/T(K-1) Fig. 13. Effect of temperature on adsorption of Hg(II) and Cd(II) on CuOs.

Table 4 Thermodynamics parameters for the adsorption of Hg(II) and Cd(II) on CuOs. Temperature (°C) Hg(II) 30 °C 40 °C 50 °C Cd(II) 30 °C 40 °C 50 °C

where kC is the distribution coefficient for the adsorption, ΔH° is the enthalpy change, ΔS° is the entropy change, ΔG° is the Gibb's free energy change, R is the gas constant, T is the absolute temperature, Ca is the Hg(II) and Cd(II) adsorbed per unit mass of the adsorbent, and Cb is the equilibrium adsorbate concentration in the aqueous phase. ΔH° and ΔS° can be calculated from the slope and intercept of the Van't Hoff plot of log Kc versus 1/T (Fig. 13). The standard free energy change, ΔG° reflects the feasibility of the process and standard entropy change ΔS° determines the disorderliness of the adsorption at solid–liquid interface. The results are shown in Table 4. The positive value of ΔS° suggests that the process of adsorption is spontaneous and thermodynamically favorable.

ΔG° (kJ/mol)

ΔS° (J/K mol)

ΔH° (kJ/mol)

−1.31 −1.43 −1.50

94.16

−15.79

−1.16 −1.25 −1.30

69.21

The solid residues of the adsorbent before and after adsorption of Hg(II) and Cd(II) ions were subjected to EDX analyses. Fig. 14 shows the EDX spectra of CuOs adsorbent before and after loading with Hg(II) and Cd(II) ions respectively. Spectrum clearly shows the peak for the presence of copper and oxygen as major constituents before adsorption. Comparing the spectrum of the CuOs loaded with Hg(II) and Cd(II) with that of unloaded one, peak for the mercury and cadmium could be observed. It was suggested that heavy metals had been adsorbed on the surface of CuOs successfully. Moreover, before and after adsorption also there are no characteristic peaks observed for any impurities. 3.9. FT-IR studies

−9.56

Fig. 15 shows the spectrum of CuOs samples before and after adsorption. The peaks near 548 cm−1 agreed with stretching vibrations

Fig. 14. EDX patterns of CuOs: (a) Before adsorption, (b) after adsorption of Hg(II), (c) after adsorption of Cd(II).

18

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

2.0 1046

1.8 1.6

Transmittance(%)

1631

(c)

2919

778

548

3511 1379

1.4 778

1.2

1046

(b)

2919

1621

adsorption with heavy metal ions. The spectrum clearly shows that the intensities of the peaks were low when compared to the intensity of sample before adsorption, which indicates the interaction of heavy metal ions on the adsorbent. And moreover as it will not show the major differences in the peaks it is maybe due to less involvement of functional groups. Therefore, it was concluded that electrostatic attraction might play a major role in the initial bulk diffusion.

1388

548

3511

3.10. Regeneration studies

1.0 548 845

0.8

2919

1046

(a)

1641

0.6

CuO before adsorption 3411 CuO after adsorption with Hg(II) CuO after adsorption with Cd(II)

0.4 0.2

1393

500

1000

1500

2000

2500

3000

3500

4000

Wavenumber(cm-1) Fig. 15. FTIR spectra of CuOs: (a) Before adsorption, (b) after adsorption of Hg(II), (c) after adsorption of Cd(II).

Desorption is an important process in adsorption studies because metal ion loaded adsorbents may create an environmental disposal issue as they are hazardous in nature. If desorption and leaching occur, desorption also enhances the economical value of the adsorption process. Desorption efficiency of the spent adsorbent was checked with 0.1 M HCl solution and the adsorption–desorption cycle was repeated three times with the same adsorbent using 0.1 M HCl for regeneration (Fig. 16). For Hg(II) after three cycles, the removal efficiency of CuOs decreased from 97.8% to 58.3%. For Cd(II), the removal efficiency after three cycles decreased from 96.2% to 56.2%. 3.11. Comparison with other adsorbents

60 40 20

st

Fir

No

of

n

tio

era

n ge

I re

re

on

i rat

ne

ge

ne

ra

II

tio

n

e reg

A Ad dso so rpt rp ion tio o n fH of C g(II d( ) o II) n on C C uOs uO s

0 ge

a us

Removal ef

80

ficiency in %

100

The comparison between the maximum adsorption capacities for Hg(II) and Cd(II) with CuOs and other adsorbents reported in the literature are given in Table 5. To the best of our knowledge the maximum adsorption capacity obtained in this study is comparable to the result from the reported adsorbents, so it is a real advantage that renders it to become a suitable alternative for the cleanup of industrial effluents from the heavy metals. 3.12. Real sample analysis To investigate the feasibility of the CuOs in actual ground water, Cd(II) contaminated ground water was sampled from the Peenya industrial area and used to test the Cd(II) removal performance. The Cd(II) concentration was found to be 0.004, which is slightly higher than the maximum contamination level of drinking water regulated by WHO (0.003 ppm). The Cd(II) removal efficiency was tested by adding 250 mg CuOs into 1000 mL contaminated groundwater. The DPASV results demonstrated that the Cd(II) removal efficiency could reach more than 98.4%. These results confirmed that CuOs generally exhibited a high efficiency of Cd(II) and other heavy metal ion removal in groundwater. 4. Conclusions

Fig. 16. Effect of CuOs regeneration on removal efficiency of Hg(II) and Cd(II) adsorption.

of Cu\O bonds [32]. The broad peak at the range of 3431 and 1635 cm−1 corresponded to the stretching vibration of \OH group, dienes respectively. Fig. 15 b–c, shows the spectrum of CuOs after

Novel hierarchical mesoporous copper oxide nanoflakes were successfully synthesized via a low-cost hydrothermal method. The results of various characterizations showed the formation of hierarchical flakes. Adsorption of Hg(II) and Cd(II) was dependent on metal ion concentration, adsorbent dose, contact time, pH and temperature. The Langmuir

Table 5 Comparison of monolayer maximum capacities of some adsorbents for Hg(II) and Cd(II) from aqueous solutions. Adsorbents

Adsorbate

Adsorption capacity (mg/g)

References

Almond shell Lentinus edodes pellets Lewatit Monoplus TP-214 Tryethylenetetramine modified polystyrin resin CuOs ZnO nanoparticles Carbon aerogel NiO nanoparticles Flower like MgO nanostructures CuOs

Hg(II) Hg(II) Hg(II) Hg(II) Hg(II) Cd(II) Cd(II) Cd(II) Cd(II) Cd(II)

135.1 403.3 456.0 344.8 1767.9 384.0 400.8 625.0 1500 1577.7

[35] [36] [37] [38] Present Study [39] [40] [41] [42] Present Study

K.Y. Kumar et al. / Powder Technology 258 (2014) 11–19

and Freundlich adsorption isotherm models were used for the mathematical explanation of the adsorption equilibrium of Hg(II) and Cd(II) ions. The maximum adsorption capacities of Hg(II) and Cd(II) were determined to be 1767.97 and 1577.78 mg/g respectively. All these values are significantly higher than those reported on other hierarchical nanostructures. The pseudo second order kinetic model was found to be best correlated to the experimental data for Hg(II) and Cd(II) adsorption. Thermodynamic studies conclude that the process of sorption of Hg(II) and Cd(II) ions were feasible, spontaneous and thermodynamically favorable. The spent adsorbents were regenerated with HCl solutions and regenerated adsorbents showed very good adsorption efficiency. All the above results demonstrated that CuOs could be used as a possible alternative low-cost adsorbent for the efficient removal of heavy metals from aqueous solutions. Acknowledgments The authors are grateful to Visvesvaraya Technological University for providing financial support under research grant scheme (Project No VTU/Aca/2010-11/a-9/11353) and for K.S. Institute of Technology, Bangalore for their great support. References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11]

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