Sorption of lead and silver from aqueous solution on phosphoric acid dehydrated carbon

Journal of Environmental Chemical Engineering 1 (2013) 934–944

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Sorption of lead and silver from aqueous solution on phosphoric acid dehydrated carbon E.I. El-Shafey *, A.H.R. Al-Hashmi Department of Chemistry, College of Science, Sultan Qaboos University, P.O. Box 36, Al-Khodh, 123 Muscat, Oman

A R T I C L E I N F O

A B S T R A C T

Article history: Received 9 April 2013 Received in revised form 11 July 2013 Accepted 3 August 2013

Dehydrated carbon was prepared from date palm leaflets via the chemical treatment with phosphoric acid. Carbon preparation was investigated by varying the acid concentration at 150 8C and selection criterion was based on the maximum removal of Pb2+ and Ag+ from aqueous solution. 13 M phosphoric acid was selected for the preparation of dehydrated carbon referred as DC13 that shows maximum removal for both metals from aqueous solution. DC13 was tested for the sorption of Pb2+ and Ag+ at different pH, contact time, metal concentration, temperature and sorbent status (moistened or dry). Maximum sorption was obtained at initial pH 5.0 for both metals. Equilibrium was reached in 1 h for Pb2+ sorption and 80 h for Ag+ sorption with higher activation energy for Ag+ sorption (31.9 kJ mol1) than that for Pb2+ sorption (11.0 kJ mol1). Moistened carbon showed better performance than the dry carbon and sorption capacity using the Langmuir isotherm was 41.5 and 37.9 mg g1 for Pb2+ sorption and 312.5 and 285.7 mg g1 for Ag+ sorption on moistened and dry carbons, respectively. Sorption of Pb2+ and Ag+ from a binary mixture showed a decrease in their uptake. The peculiar behavior of metal sorption is related to the chemical reduction of Ag+ to Ag0 on the carbon surface, however, no chemical reduction was involved in Pb2+ sorption. ß 2013 Elsevier Ltd. All rights reserved.

Keywords: Sorption Pb2+ Ag+ Dehydrated carbon Phosphoric

Introduction Lead and silver, among heavy metals, are involved in many industries. Lead is utilized in electronic devices, battery recycling, traditional medicine, paint and cosmetics. Silver is used in silver batteries, electronic devices, mirroring, photographic, electroplating, catalyst and ink-formulation industries. Lead is a potent neurotoxicant with adverse health effects ranging from death [1] to impaired cognitive and behavioral development that can have lifelong consequences for children. Lead can adversely affect the liver, kidney, lungs, brain, spleen, muscles and heart. Silver also causes many adverse effects to human health such as stomach pain, breathing problems and irritation of throat and lungs [2]. Silver compounds can lead to the deposition of silver–protein complexes in body tissues causing argyria, a permanent bluishgray discoloration of skin [2]. Based on the Omani legislations [3], maximum allowed concentrations of lead in the aquatic environment is 0.1 mg L1 for irrigating public parks and agricultural areas and 0.2 mg L1 for irrigating areas with no public access. However, the maximum allowed concentration of silver in both of the above areas is 0.01 mg L1.

* Corresponding author. Tel.: +968 99822317; fax: +968 24141469. E-mail address: [email protected] (E.I. El-Shafey). 2213-3437/$ – see front matter ß 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jece.2013.08.004

Existing physical and chemical methods for the removal of heavy metals from wastewater include chemical precipitation, electro-coagulation, membrane filtration and reverse osmosis. Chemical precipitation processes produce voluminous sludge that requires further treatment such as chemical precipitation. In addition, membrane fouling and membrane limited life are shortcoming for membrane filtration and reverse osmosis [4]. Above all, these methods are either expensive or not appropriate for the treatment of dilute metal solutions [5]. Adsorption technologies utilizing activated carbon appear to be expensive due to the high cost of carbon preparation. Different adsorbents were tested for the removal of Pb2+ from aqueous solutions including dehydrated carbon from flax shive via sulfuric acid treatment [6], lignin [7], tea waste [8], brewery waste biomass [9], hazel nut [10], almond shell [10], waste maize bran [11] and activated carbon [12]. Ag+ removal was also investigated using activated carbon [13] perilite [14], peat [15], coal [16] and sulfuric acid dehydrated carbon from [17]. Date palm (Phoenix dactylifera L.) is the main crop in Oman and the Gulf States [18]. 180,000 and  3 million tons of palm leaflets, as agricultural waste, are produced annually in Oman and the Gulf States, respectively. Burning in the field is a common practice that poses environmental air pollution. Leaflets of date palm (P. dactylifera) as plant material possess cellulose (33.5%), hemicelluloses (26%) and lignin (27%), as main components [19]. Concentrated phosphoric acid acts as a dehydrating agent, however, with less power than concentrated

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

sulfuric acid [20]. It is often used in place of sulfuric acid to dehydrate alcohols to alkenes because phosphoric acid is a weaker acid, less powerful oxidizing agent and less destructive than sulfuric acid [21]. Accordingly, dehydrated carbon via phosphoric acid is expected to possess good reduction properties. In addition, residual phosphoric acid from the preparation process can be easily converted to phosphate fertilizer by neutralization. The objective of this research was to prepare dehydrated carbon from date palm leaflets using phosphoric acid and investigate its adsorption properties in the removal Pb2+ and Ag+ from single and binary solutions. Materials and methods Carbon preparation Clean air-dried palm leaflets (40 g) were cut into small pieces and were mixed with 160 mL of 9, 11, 13 and 14.7 mol L1 phosphoric acid to produce dehydrated carbons DC9, DC11, DC13 and DC15, respectively. The mixture was heated in 20 min to 150  2 8C and was maintained in that temperature range for 25 min with occasional stirring. The black mixture was allowed to cool at room temperature, and then filtered using porcelain Buchner funnel under vacuum. Spent phosphoric acid with fine carbon particles were filtered off. The carbonized material was then washed several times with distilled water and stored under dilute phosphoric acid condition to avoid bacterial effects. Before use for metal sorption, a sample of carbonized product was washed in Gooch crucible (G1) by distilled water several times to remove the acidity until the wash water did not show a change in methyl orange color. The moistened carbon was washed with a stream of distilled water between two sieves of 16 and 60 mesh to remove fine particles and to select a size range for sorption experiments. The sample was transferred to a Gooch crucible and left under suction for 30 min. Appropriate amounts of the moistened carbon were then used in sorption experiments and a sample of 1 g of the moistened carbon was separated to measure the moisture content using oven drying at 120 8C to constant weight. The moisture content was determined in each experiment using the moistened carbon. For work under dry conditions a sample of the moistened carbon (acid free) was dried in an oven at 120 8C to constant weight, transferred to the desiccators to cool. The carbon was ground and a size range between two sieves of 16 and 60 mesh were selected for sorption experiments. The carbon was then stored in a dry, clean and well closed polyethylene jar. Dry dehydrated carbons (DC9, DC11, DC13 and DC15) were investigated for their capability to remove Pb2+ and Ag+ from aqueous solution. The yield of the carbon product was also determined. Physico-chemical properties Surface functionality was investigated via Boehm titrations method [22]. Scanning electron microscope (SEM) analysis was carried out using a Goel JSM 840-A scanning electron microscope (Jeol, Japan). X-ray powder diffraction (XRD) was carried out using a Philips PW 1830 generator with a Philips PW 1050 powder goniometer (Philips, USA). Copper Ka was used as the incident radiation. Both SEM and XRD were carried out for metal loaded and unloaded carbon samples. The surface area of the dehydrated carbon was measured using Autosorb-1 (Quantachrome Instruments, USA) via nitrogen adsorption at 77 K. FTIR was conducted for date palm leaflets and dehydrated carbon. Pellets made of a mixture of 1.0 mg of carbon or date palm leaflets and 10 mg of KBr were pressed at high pressure and the pellet was scanned in transition mode using FT-IR spectrometer (spectrum BX, Berkin Elmer, Germany) through a wavelength range from 4000 to

935

800 cm1. Apparent density and ash contents were determined using standard methods [23,24]. Metal sorption All the chemicals used were of analytical grade. Stock solutions, 1000 mg L1 of both Pb2+ and Ag+, were prepared in deionized water using lead and silver nitrate (PbNO3, AgNO3), respectively. All the working solutions were prepared by diluting the stock solution in deionized water. For carbon selection, samples of 0.1 g of the dehydrated carbons were mixed with 50 mL (200 mg L1) of Pb2+ or Ag+ solution at initial pH 5. The samples were shaken mechanically in a shaking water bath at a shaking rate of 100 rpm at 25 8C. Maximum sorption of both metals was obtained at initial pH 5.0 (as shown later). Once equilibrium was reached, aliquots of supernatant were separated and analyzed using AAS (Varian/SpectraAA/220FS, Australia). Since the sorption capacity of Pb2+ was lower than that of Ag+ on the selected carbon (DC13), from carbon selection experiments as shown later, further Pb2+ sorption studies (kinetics and equilibrium) were carried out using lower concentration range than that tested for Ag+ sorption. In the kinetic experiments 0.2 g of the dry carbon, (or its equivalent mass of the moistened carbon) was mixed with 100 mL of metal solution (50 mg L1 for Pb2+ and 100 mg L1 for Ag+) at initial pH 5. At different periods of time, aliquot of supernatant was withdrawn for metal analysis. The kinetic studies were conducted at 25, 35 and 45 8C with mechanical agitation. The effect of initial pH was studied as follows. Samples of 0.1 g of DC13, were mixed with 50 mL (50 mg L1 for Pb2+ and 100 mg L1 for Ag+) at different initial pH values. The pH was adjusted using drops of dilute NaOH or HNO3 before the addition of the pre-weighed sorbent. Samples were shaken mechanically at 25 8C until equilibrium was reached and final pH was recorded. For the isotherm studies, 0.1 g of the dry carbon (or its equivalent weight in the moistened status) was mixed with 50 mL of metal solutions (25–200 mg L1 for Pb2+ and 50–1000 mg L1 for Ag+) at different temperature (25–45 8C) at initial pH 5. For sorption from binary metal solution, samples of 0.1 g of the dry DC13 were added to mixtures of different metal concentrations at initial pH 5 and at different temperature (25–45 8C). The samples were shaken mechanically until equilibrium was reached. After equilibrium, aliquot of supernatant was withdrawn and metal concentration was analyzed. Metal loaded carbon samples for SEM and X-ray diffraction analysis were obtained as follows. 0.1 g of carbon was mixed with 50 mL of Pb2+ (200 mg L1) and Ag+ (1000 mg L1) at initial pH 5.0 (optimum pH as shown later). After the equilibrium was reached under continuous shaking, loaded carbons were separated, washed with distilled water and left to dry at 120 8C and then tested for SEM and X-ray diffraction. Results and discussion Carbon selection Under the sorbent preparation conditions, carbonization of cellulose and hemicelluloses via dehydration (removal of water) took place. Partial oxidation is expected to take place to cellulose, hemicellulose and lignin. Criterion of carbon selection was mainly related to the sorption capacity of Pb2+ and Ag+ from aqueous solution on produced dehydrated carbons. Preliminary studies showed that concentrations of phosphoric acid lower than 9 M could not carbonize the date palm leaflets efficiently. As presented in Fig. 1, Pb2+ sorption did not show much variation for the carbons prepared. On the other hand, Ag+ sorption has increased from

[(Fig._1)TD$IG]

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

Metal Sorbed (mg/g)

936 120

acid respectively.

100

Yield ð%Þ ¼

80

60

(1)

The slight loss in the yield with the increase of acid concentration is related to the harsher preparation conditions using concentrated phosphoric acid that led to more fragmentation of carbon particles produced. Selected DC13 showed a yield of 54%. In a previous study [20], sulfuric acid was used to prepare dehydrated carbon from flax shive at 160 8C. The carbon yield was 58.3% using 12 M acid while using the concentrated acid (18.2 M) the yield has decreased to 51.3% due to harsher conditions.

Ag(I) Pb(II)

40

20

Physico-chemical characterization

0 8

9

10

11 12 13 14 Phosphoric Acid Concentration (mol/L)

15

16

Fig. 1. Sorption of Pb2+ and Ag+ on dehydrated carbons prepared at different concentrations of phosphoric acid. (1 g leaflets: 4 mL acid, reaction time 25 min, temperature 150 8C). Error bars indicate standard deviation (n = 3).

32 mg g1 on DC9 to 111 mg g1 on DC13 and DC15. Based on these results, DC13 was selected for the comprehensive sorption studies of both metals from aqueous solutions. DC15 did not show an advantage over DC13 for metal sorption, thus, DC15 was avoided to minimize acid consumption. Spent acid was collected after the preparation process and was found to be 14.5 M. The acid can be reused or utilized in phosphate fertilizer preparation. Carbon yield, Eq. (1), was in the range of 56–53.5% using 9–14.7 M

[(Fig._2)TD$IG]

mass of carbon produced ðgÞ  100 mass of dry date palm leaflets ðgÞ

Dry DC13 was selected for physico-chemical characterization and metal sorption. Apparent density and ash content were 0.79 g cm3 and 13.9%. FT-IR spectrum of DC13 compared to date palm leaflets (Fig. 2) shows the presence of hydroxyl (3399– 3422 cm1), carboxyl (1738–1707 cm1), –COO or C5 5C (1654– 1624 cm1) and other carbon-oxygen species (1252–1080 cm1) [25,26]. In addition, the band at 1109 cm1 for DC can also be attributed to P–O asymmetric stretching vibration [27]. It can be concluded that the bands at 3422, 1707 and 1109 cm1 for the dehydrated carbon are more intense than comparable peaks for the leaflets as a result of surface carbonization and partial oxidation accompanied the interaction of date palm leaflets with phosphoric acid. Surface acidic functional groups were determined via Boehm titrations [22] where NaHCO3 neutralize carboxylic, Na2CO3 neutralize carboxylic and lactonic, and NaOH neutralize carboxylic, lactonic and phenolic functional groups. Number of milliequivalents (meq) calculated based on those titrations for the surface

Fig. 2. FTIR of (A) date palm leaflets and (B) dehydrated carbon (DC13).

[(Fig._3)TD$IG]

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937

80

16

14 70 12

+

50

Ag sorbed (mg/g)

8

2+

Pb sorbed (mg/g )

60 10

6 40 4 30 2

Pb(II) Ag(I)

0

20 1

2

3

4

5

6

7

Initial pH Fig. 3. Sorption of Pb2+ and Ag+ at different initial pH values and 25 8C. Error bars indicate standard deviation (n = 3).

carboxyl, lactone and phenol groups were 0.82, 1.1 and 0.92 meq g1, respectively. The surface area (SBET), pore volume and pore diameter of dry DC13 was 7.8 m2 g1, 1.34  103 cm3 g1 and 6.8 A˚ respectively. Similar low surface areas were found for dehydrated carbons prepared via sulfuric acid treatment from date palm leaflets (24 m2 g1) [18]. The low surface area could be related to the abundance of polar functional groups, such as –COOH, –OH, which occupy large fraction of the surface [28]. Hot phosphoric acid dehydrates and partially oxidizes date palm leaflets under the preparation conditions creating ion exchange sites such as –COOH, –OH. X-ray diffraction showed an amorphous structure (Fig. 1, Appendix).

Kinetics of sorption Pb2+ sorption shows a much faster kinetics, however, with less metal uptake than Ag+. Equilibrium was reached within 1 h for Pb2+ sorption, but for Ag+ sorption, approximate equilibrium was reached within 80 h (Fig. 5A and B) with better performance for the moistened carbon than for the dry carbon for both metals. The dry carbon showed visible limited swelling when immersed in the aqueous solution giving narrower pores for metal diffusion. The uptake of Pb2+ and Ag+ was found to vary almost linearly with the half power of time, t0.5, in the initial stages of sorption (Eq. (2)) [32]. q ¼ kd t 0:5

(2) 1

pH effect As presented in Fig. 3, at low pH values, sorption of both metal ions was low, and with the rise in the initial pH, metal uptake increased. No significant changes took place between pH 4 and 6 for Pb2+ sorption, and between pH 5 and 6 for Ag+ sorption. Previous results showed a similar increase in Pb2+ sorption [14,29] and Ag+ sorption [17,30] on different sorbents with initial pH increase. The increase in removal capacity at higher pH may also be attributed to the decrease of H+ ions which compete with Pb2+ or Ag+ ions for sorption sites. The speciation of Pb2+ and Ag+ in aqueous solution is pH dependent. At pH of 2–5 lead exists as Pb2+. Above pH of 6, it is hydrolyzed to PbOH+ and Pb(OH)2 [29]. The predominant sorbing forms of lead are Pb2+ which occurred in the pH range of 4–6. On the other hand, Ag+ species are dominant in aqueous solution of silver nitrate at pH 5 [31]. Higher initial pH values were avoided to prevent precipitation of Pb2+ or Ag+ as hydroxides [29]. Metal sorption was accompanied by a decrease in the final pH indicating protons release in solution and suggesting an ion exchange mechanism. Plotting [H+] released in solution against Pb2+ and Ag+ uptake would give straight lines of slope 2 and 1 respectively, if the ion exchange process is the dominating mechanism of metal uptake. The slopes of [H+] released/[metal] sorbed were 1.91 and 0.99 for Pb2+ and Ag+ sorption, respectively (Fig. 4A and B) suggesting ion exchange as a dominating mechanism for both metals.

where q (mg g ) indicates the metal sorbed and kd is the intraparticle diffusion rate constant. kd values are presented in Table 1 in units of (mg g1 h0.5). Kinetic sorption data were tested for pseudo first and pseudo second order kinetic models [14] (Eqs. (3) and (4)). The pseudo second order model shows better fitting for kinetic sorption data than pseudo first order model, as presented in Table 1. This points out that the sorption of Pb2+ and Ag+ fulfills very well the pseudo second order kinetic reaction. This reflects that the rate of sorption depends on both of the adsorbent and adsorbates [33]. logðqe  qt Þ ¼ logqe  t 1 t ¼ þ qt k2 q2e qe

K1t 2:30

(3)

(4)

where k1 and k2 are the rate constants for the pseudo first order and pseudo second order models, respectively. qt (mg g1) is the amount of metal sorbed per unit weight of sorbent at any time t. h = kqe2, where h is the initial adsorption rate. The values of k2, h and kd (Table 1) show an increase for the moistened carbon over the dry carbon. This could be related to the shrinkage and compaction that took place on drying. Drying has probably resulted in narrower pores for the diffusion of metal ions and this is reflected on the values of these constants as similarly found for Zn2+ and Hg2+ sorption on dehydrated carbon via sulfuric acid treatment [34]. Dehydrated carbon, once dried, shows less performance in metal sorption than when kept moistened [6]. In

[(Fig._4)TD$IG]

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938

exchange mechanism. However, for Ag+ sorption, Ea was found to be 31.4 and 31.9 kJ mol1 for moistened and dry carbons respectively. Such higher Ea values could be related to the involvement of another chemical process in Ag+ sorption in addition to ion exchange, as shown later. In a previous study, Pb2+ sorption on modified spent grain showed activation energy 12.33 kJ mol1 and sorption mechanism was related to ion exchange [36]. Ea for Ag+ sorption on amino methylene phosphonic acid resin was 22.8 kJ mol1 and mechanism was dominantly related to coordination bonds between the silver cation and the resin [37].

0.25

(A)

0.15

y = 1.9164x 2 R = 0.9922

0.1

+

H released (mmol/g)

0.2

0.05

Equilibrium sorption

0 0

0.02

0.04

0.06

0.08

0.1

0.12

2+

Pb sorbed (mmol/g) 0.8

(B)

0.6

y = 0.9918x 2 R = 0.9875

0.5

+

H released (mmol/g)

0.7

0.4

0.3 0.3

0.35

0.4

0.45

0.5

0.55

0.6

0.65

0.7

0.75

0.8

Ce 1 Ce þ ¼ qe bq q

+

Ag sorbed (mmol/g) Fig. 4. Plot of protons released in solution versus sorbed (A) Pb2+ and (B) Ag+.

addition, rising the temperature has led to a rise in rate constant values. Temperature has probably led to more swelling and thus, wider pores for the diffusion of metal ions. Similar behavior was found for the increased rates of sorption for both moistened and dry dehydrated carbons on rising the temperature for the removal of different metal ions [6,34]. In comparison, the values of k2 and h were much higher for Pb2+ sorption than for Ag+ sorption. In addition, qe (calculated from the pseudo second order model) for Ag+ sorption is higher than that for Pb2+ sorption (Table 1). This reflects a different sorption mechanism for both metals. Because of the low correlation values of the pseudo first order rate constant, k1 and related qe were not discussed. As pseudo second order model fit well the kinetic sorption data, k2 values at different temperature were used to calculate the activation energy, Ea, from the Arrhenius equation (Eq. (5)). k ¼ AeEa =RT

Sorption of both Pb2+ and Ag+ on DC13 follows an L- type adsorption isotherm, with better performance for the moistened carbon over the dry (Fig. 6A and B). The increase of metal sorption for the moistened over the dry carbon could be related to the shrinkage and compaction that took place on drying as mentioned earlier. Better removal capacity of metals was found for moistened carbon over the dry for dehydrated carbons prepared via sulfuric acid treatment [6,34] and was related to the shrinkage of the latter. By rising the temperature, the uptake of Pb2+ and Ag+ has increased, (Fig. 6, Table 2), and this could be due to an increased swelling as a result of temperature leading to more access for metal ions to sorption sites. Similar explanation was reported earlier utilizing sulfuric acid dehydrated carbons in metal removal [6,34]. The equilibrium sorption data were found to fit the Langmuir equation (Eq. (6)), for both metals (Table 2) more than the Freundlich equation (Eq. (7)) indicting that the sorption process generates monolayer formation.

(5)

Ea is the activation energy of metal sorption (KJ mol1), R is the gas constant (8.314 J mol1 K1), T is the solution temperature (K) and A is the pre-exponential factor (frequency factor). Ea for physical adsorption is usually less than 4.2 kJ mol1 as the equilibrium is rapidly obtained and easily reversible representing weak adsorption forces [35]. Chemical adsorption is specific and involves stronger forces than physical adsorption. Ea for chemical adsorption is of the same extent as the heat of chemical reactions and this means that the rate varies with temperature according to finite activation energy (between 8.4 and 83.7 kJ mol1) in the Arrhenius equation [35]. Ea for Pb2+ sorption was found to be 11.2 and 10.2 kJ mol1 for moistened and dry carbons, respectively, representing chemical ion

(6)

where q (mg g1) and b (L mg1) are the Langmuir constants related to maximum adsorption capacity and the relative energy of adsorption, respectively. 1 log qe ¼  log C e þ log K n

(7)

where K (l1/n mg11/n/g) and 1/n are the constants, which are considered to be the relative indicators of adsorption capacity and adsorption intensity, respectively. Sorption of Pb2+ and Ag+ in their nitrate from a binary mixture of both metals was examined at different temperature (25–45 8C) using dry DC13. The equilibrium data were examined for the Langmuir and Freundlich isotherms (Table 2). Compared to sorption from single metal system, the Langmuir sorption capacity for Pb2+ in a binary mixture decreased to 40%, at 25 8C, and 65%, at 45 8C. On the other hand, Ag+ sorption capacity decreased in the binary mixture to 37%, at 25 8C, and 40%, at 45 8C, compared to Ag+ sorption from a single Ag+ solution. It is evident that the presence of both metal ions in the same sorption solution has mutually decreased the sorption capacity of each other. These results are in good agreement with previous reports showing that the sorption of metals in a multiple metal system is lower than that in a single metal system [38,39]. Sorption capacities of Pb2+ and Ag+ from single metal solution (Table 3), in this study, were compared with other sorbents in other studies. Thermodynamic calculation Thermodynamic parameters were calculated from the variation of the equilibrium constant, Kc, at different temperature (Eq. (8)). Kc ¼

C Ae Ce

(8)

[(Fig._5)TD$IG]

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939

20

(A)

18 16

12 10

2+

Pb sorbed (mg/g)

14

8 6 4

25 C, moistened 35 C, moistened 45 C, moistened 25 C, dry 35 C, dry 45 C, dry

2 0 0

20

40

60

80

100

120

140

160

180

200

Time (min)

90

(B)

80

70

+

Ag sorbed (mg/g)

60

50

40

30 25 C, moistened 35 C, moistened 45 C, moistened 25 C, dry 35 C, dry 45 C, dry

20

10

0 0

20

40

60

80

100

120

Time (hr) Fig. 5. Sorption of (A) Pb2+ and (B) Ag+ on DC13 against time at different temperature. (Initial pH 5.0, volume of metal solution 100 mL, shaking speed 100 rpm, initial concentration 50 mg L1 for Pb2+ and 100 mg L1 for Ag+). Error bars indicate standard deviation (n = 3).

where CAe the equilibrium amount of metal (mg) adsorbed per liter of the solution, and Ce is the equilibrium concentration (mg L1) of metal in the solution. The initial part of the adsorption isotherm in which qe versus Ce is linear is used to estimate Kc [40]. As presented in Table 4, the value of Kc increases with temperature indicating an endothermic nature of the sorption of both metals from their respective single solutions. The Gibbs free energy change of the

adsorption process is related to Kc as in Eq. (9) [40].

DG ¼ RT ln K c

(9)

The changes in enthalpy (4H8) and entropy (4S8) for metal adsorption were calculated from the slope and the intercept of the plot of ln Kc against 1/T according to the van’t Hoff equation [40],

0.9984 0.9986 0.9996 0.9978 0.9958 0.9994 (1.26%) (1.20%) (0.67%) (1.49%) (2.06%) (0.752%) 78.1 81.3 84.7 71.4 74.6 79.4 8.98 (4.30%) 13.9 (5.33%) 23.5 (5.2%) 6.14 (4.8%) 10.30 (4.71%) 16.81 (5.24%) (4.64%) (5.04%) (3.45%) (3.04%) (2.22%) (4.34%) 0.00147 0.00210 0.00327 0.00120 0.00185 0.00267 0.9862 0.9898 0.9440 0.9864 0.9823 0.9788 (0.19%) (6.44%) (14.4%) (6.80%) (6.25%) (10.9%) 53.1 49.9 40.7 52.6 53.7 44.9 (11.4%) (4.29%) (9.20%) (5.33%) (4.15%) (6.59%) 0.0350 0.0382 0.0408 0.0329 0.0424 0.0424

(0.68%) (0.68%) (0.54%) (1.1%) (0.87%) (0.707%) 14.03 15.38 17.15 12.33 13.87 14.53 (4.34%) (4.34%) (4.61%) (3.42%) (4.75%) (4.70%) 3.47 4.98 6.85 1.60 2.39 2.89 1.05 (2.28%) 1.26 (3.63%) 1.392 (3.19%) 0.63 (1.36%) 0.744 (2.55%) 0.816 (3.04%) 0.9916 0.9513 0.9512 0.9102 0.9935 0.9809 (6.835%) (18.6%) (24.0%) (15.2%) (5.27%) (11.7%) 5.58 5.25 5.61 5.93 6.34 7.11

R Monolayer, qe, mg g1 Rate const, k1, h1

2.172 (3.76%) 2.322 (8.55%) 2.988 (8.56%) 1.41 (12.8%) 1.878 (3.06%) 2.376 (5.27%)

Monolayer, qe, mg g1 Initial adsorption rate, h, mg g1 h1 Rate const, k2, g mg1 h1

Pseudo second order model

2

Pseudo first order model

Eq. (10). ln K c ¼

DS R



DH  RT

(10)

In single solution, the negative values of 4G8 (Table 4) indicate a favorable and spontaneous process for the temperature range evaluated which is usually the case for many adsorption systems in solution [38] 4G8 values are more negative for Ag+ sorption than for Pb2+ sorption. This probably indicates more spontaneity for Ag+ sorption. However, for sorption from binary mixture, 4G8 has increased and stayed negative for Pb2+ sorption, however, 4G8 showed an increase to positive values for Ag+ sorption. This may be related to the difference in the interaction between the metals and the carbon surface reflecting non-equal competition of the two cations to sorption sites on a heterogeneous surface. In a previous study, Atar et al. [41], found that 4G8 values were negative for the adsorption of Acid Red 183 dyes and Reactive Blue 4 from single solution on boron industrial waste. However, 4G8 values changed to be positive for the sorption of the mixture of both dyes with temperature rise. In the current study, the results reflect the preference of Pb2+ sorption, as a divalent cation over Ag+, a monovalent cation for sorption sites. The system of Ag+ sorption under such binary conditions has probably gained energy from an external source at high temperature [41]. The positive values of 4H8 show that metal adsorption is endothermic for both metals from single or binary solutions. 4H8 values between 0 and 20.9 kJ mol1 refer to physisorption processes [42], however, values between 20.9 and 418.4 kJ mol1 1 are considered for chemical adsorption process [42]. In this study, 4H8 values for the sorption Pb2+ (Table 4) represent physisorption that is dominated by ion exchange processes as shown earlier. As it is a reversible process, ion exchange process is considered as a physical process [43]. Other research studies showed low 4H8 values for the adsoption of Pb2+ on pineapple sludge (5.2 kJ/mol) [44], paper sludge (2.9 kJ/mol) [44] and magnetic eggshell-Fe3O4 (7.83 kJ mol1) [45]. On the other hand, Ag+ sorption showed higher 4H8 values than 20 kJ mol1. This indicates the involvement of chemical process (chemical reduction as shown later) together with ion exchange. The positive value of 4S8 indicates an increase in randomness at the solid–solution interface during the adsorption of metals onto the carbon adsorbents.

Dry

15.0 (1.11%) 22.7 (1.45%) 31.35 (2.23%) 10.8(1.23%) 16.0 (1.17%) 23.3 (1.32%) 25 35 45 25 35 45 Wet Ag+

Dry

72.0 (2.16%) 102.0 (3.06%) 112.2 (2.33%) 64.80 (2.26%) 73.80 (2.13%) 75.60 (2.22%) 25 35 45 25 35 45 Wet Pb2+

Pore diffusion constant, kd mg g1 h0.5 Sorbent

Temp (8C)

SEM and X-ray diffraction

Metal

Table 1 Pore diffusion and rate constants of pseudo first order and pseudo second order models for the kinetics of Pb2+ and Ag+ sorption.

0.9997 0.9996 0.9998 0.9992 0.9994 0.9996

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

R2

940

SEM photograph of DC13 showed that the dehydrated carbon possesses the plant fibrous structure (Fig. 7a). Ag particles were clearly identified on the carbon surface (Fig. 7c) however no changes were identified on lead-loaded carbon indicating the absence of chemical reduction of Pb2+ on the carbon surface (Fig. 7b). X-ray powder diffraction pattern (Fig. 8) has confirmed the identification of silver metal deposits with the interplanar spacing (d) and relative line intensities (I) agreeing with those recorded [46]. This indicates the reduction of Ag+ to Ag0 on the carbon surface Metal recovery and carbon recycle For a carbon sample (0.5 g) loaded with Pb2+ (38 mg/g), desorption using nitric acid (0.5 M) reached 95% at such acidic conditions. The protons replaced Pb2+ cations on ion exchange sites. The carbon was washed to become acid-free, and was the reused in Pb2+ sorption and sorption capacity has reached to 37.7 mg/g. Desorption for the second time has led to a recovery of 96%. For silver loaded samples, elemental silver was recovered by immersion in 10 mL 7 M nitric acid at 90 8C for about 20 min.

[(Fig._6)TD$IG]

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

941

40

(A) 35

25

20

2+

Pb sorbed (mg/g)

30

15

10

25 C, moistened 35 C, moistened 45 C, moistened 25 C, dry 35 C, dry 45 C, dry

5

0 0

20

40

60

80

100

120

140

160

Equilibrium Concentration (mg/L) 350

(B) 300

+

Ag sorbed (mg/g)

250

200

150

100

25 C, moistened 35 C, moistened 45 C, moistened 25 C, dry 35 C, dry 45 C, dry

50

0 0

100

200

300

400

500

600

700

800

900

Equilibrium Concentration (mg/L) Fig. 6. Sorption of (A) Pb2+ and (B) Ag+ at initial pH 5 at different temperatures.

Nitric acid at such temperature was found to dissolve silver particles as silver nitrate and oxidize the dehydrated carbon completely to carbon dioxide.

Eqs. (11) and (12). 2COOH þ Pb2þ ¼ ðCOOÞ2 Pb þ 2Hþ

(11)

Interaction analysis

2COH þ Pb2þ ¼ ðCOÞ2 Pb þ 2Hþ

(12)

As mentioned earlier, at pH 2–5, Pb2+ species become dominant in aqueous solutions [29]. In addition, as presented earlier, the molar ratio of [H+] released in solution, as a result of Pb2+ sorption, and Pb2+ sorbed was 2. Consequently, ion exchange processes seem to dominate sorption mechanism for Pb2+ as presented in

On the other hand, Ag+ species are dominant in aqueous solution of silver nitrate at pH 5 [31]. Presence of silver crystals on the sorbent surface, as shown in SEM and X-ray diffraction pattern, reflects Ag+ reduction. Ion exchange processes for Ag+ sorption can be represented as in Eqs. (13) and (14). Possible redox

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

942

Table 2 Langmuir and Freundlich parameters for the sorption of Pb2+ and Ag+ at pH 5 at different temperature. Sorption system

Metal

Single

Pb2+

Sorbent

Sorption temp. (8C)

Langmuir constants 1

Dry

Ag+

Wet

Dry

Binary

Pb2+

Dry

Ag+

Dry

Freundlich constants

Correlation value, R2

)

b (L mg

1/n

K

(3.25%) (2.46%) (1.78%) (3.25%) (2.63%) (1.78%) (0.99%) (2.50%) (2.40%) (1.84%) (2.18%) (2.82%)

0.0409 0.0426 0.0445 0.0410 0.0423 0.0440 0.0113 0.0122 0.0135 0.0103 0.0114 0.0124

(4.84%) (3.54%) (3.13%) (4.32%) (4.05%) (3.42%) (4.10%) (1.09%) (1.12%) (5.34%) (4.5%) (4.67%)

0.9955 0.9959 0.9977 0.9958 0.9976 0.9987 0.9994 0.9963 0.9966 0.9980 0.9972 0.9953

0.401 (8.55%) 0.412 (9.83%) 0.432 (11.8%) 0.391 (8.85%) 0.414 (9.56%) 0.420 (12.8%) 0.443 (9.66%) 0.415 (9.04%) 0.417(8.49%) 0.439 (8.86%) 0.424 (7.63%) 0.412 (6.51%)

4.15 (13.2%) 4.55 (15.8%) 4.86 (19.0%) 3.90 (13.5%) 4.02 (15.4%) 4.60 (20.5%) 14.4 (21.9%) 18.7 (19.0%) 21.8 (17.3%) 14.0 (20.0%) 16.9 (16.5%) 20.2 (13.4%)

0.0481 0.0364 0.0359 0.0017 0.0019 0.0282

(4.21%) (3.11%) (4.21%) (4.28%) (3.77%) (3.62%)

0.9955 0.9978 0.9983 0.9953 0.9967 0.9930

0.338 0.399 0.462 0.727 0.713 0.663

q (mg g Wet

Correlation value, R2

1

25 35 45 25 35 45 25 35 45 25 35 45

32.4 37.2 41.5 29.4 32.7 37.9 250.0 270.3 312.5 243.9 263.2 285.7

25 35 45 25 35 45

11.73 (2.14%) 18.12 (2.21%) 24.69 (3.14%) 91.74 (1.12%) 103.1 (1.24%) 116.3 (2.33%)

)

(7.23%) (12.36%) (8.03%) (6.17%) (6. 3%) (7.12%)

1.99 (17.7%) 2.25 (21.2%) 2.512 (16.9%) 0.450 (11.7%) 0.581 (11.1%) 1.129 (12.5%)

0.9716 0.9628 0.9471 0.9696 0.9647 0.9381 0.9470 0.9530 0.9580 0.9550 0.9663 0.9752 0.9529 0.9332 0.9686 0.9877 0.9867 0.9830

Table 3 Maximum sorption capacities of Pb2+ and Ag+ on different sorbents. Metal ion

Sorbent material

Maximum sorption (mg/g)

References

Pb2+

Sulfuric acid dehydrated carbon C200 Waste maize bran Lignin Sulfuric acid dehydrated carbon C160 Brewery waste biomass Tea leaves Phosphoric acid dehydrated carbon at (moistened) Phosphoric acid dehydrated carbon at (dry) Hazelnut shell Activated carbon Almond shell

147.1 142.86 102.40 87.0 85.6 65.0 41.5 37.9 28.18 21.9 8.08

[6] [18] [7] [6] [9] [8] This study This study [10] [11] [10]

Ag+

Phosphoric acid dehydrated carbon (moistened) Phosphoric acid dehydrated carbon (dry) Coal Brewery waste biomass Multisorb Peat Expanded perlite (EP) Coke

312 286 107 42.7 36.1 10.8 8.60 4.90

This study This study [15] [9] [14] [14] [13] [14]

Table 4 Thermodynamic parameters of Pb2+ and Ag+ sorption from single and binary systems at different temperature. System

Metal

Carbon

Temp. (K)

Kc

4G8 (kJ/mol)

Single

Pb2+

Wet

298 308 318 298 308 318 298 308 318 298 308 318

2.36 2.58 2.86 2.10 2.28 2.49 2.52 3.56 5.04 2.29 3.25 4.43

2.13 2.43 2.77 1.83 2.11 2.41 2.29 3.25 4.28 2.05 2.91 3.93

298 308 318 298 308 318

1.19 1.28 1.36 0.302 0.415 0.667

0.429 0.626 0.808 2.96 2.25 1.07

Dry

Ag(I)

Wet

Dry

Binary

Pb2+

Dry

Ag+

Dry

4H8 (kJ/mol)

4S8 (J/mol)

7.53

32.4

6.74

28.8

27.3

99.4

26.0

93.9

5.23

31.1

19.0

94.1

[(Fig._7)TD$IG]

[(Fig._8)TD$IG]

E.I. El-Shafey, A.H.R. Al-Hashmi / Journal of Environmental Chemical Engineering 1 (2013) 934–944

943

Fig. 8. X-ray diffraction pattern of silver loaded carbon.

carbon produced, in this study, to reduce Ag+ to Ag0 but not Pb2+ is considered to be an obvious cause of the different sorption behavior of both metals in terms of kinetics and equilibrium. Other carbonaceous sorbents were found to reduce Ag+ to elemental silver such as activated carbon [13], activated carbon fiber [47] and dehydrated carbon via sulfuric acid treatment [17]. Conclusion

 COOH þ Agþ ¼ COOAg þ Hþ

(13)

 COH þ Agþ ¼ COAg þ Hþ

(14)

Dehydrated carbon prepared from date palm leaflets via phosphoric acid treatment possesses ion exchange and reduction properties. Dehydrated carbon DC13 that was prepared using 13 M phosphoric acid was selected for the sorption study of Pb2+ and Ag+. Sorption was much faster for Pb2+ than for Ag+ following well pseudo second order model. Ea was found to be 11 kJ mol1 for Pb2+ sorption representing ion exchange mechanism, however, for Ag+ sorption, Ea was 31 kJ mol1 due to the involvement of chemical reduction that was confirmed through SEM and X-ray diffraction. Sorption of Pb2+ and Ag+ followed well the Langmuir equation with an increase in metal uptake as temperature rises, probably, as a result of an expected increase in the swelling of the sorbent allowing more active sites to become available for metal ions. Moistened carbon in general shows better performance than the dried carbon possibly as a result of shrinkage and compaction on drying. The different behavior of both metals toward the phosphoric acid dehydrated carbon is related to the reduction capability of the sorbent for Ag+ from aqueous solution unlike Pb2+. The produced dehydrated carbon can be used for the removal of metals in industrial effluents such as mining effluents and other wastewater. The carbon preparation is considered to be economic as it is a fast process and the residual phosphoric acid (14.5 M) can be reused in further carbon preparation or it can be used to produce phosphate fertilizer. In addition, the dehydrated carbon produced can be reused for the removal of Pb2+ ions and it can be used to recover the precious silver in its elemental state.

 CH þ Agþ þ H2 O ¼  COH þ Ag0 þ Hþ

(15)

References

 CH=  COH þ Agþ þ H2 O ¼  C¼O þ Ag0 þ Hþ

(16)

 CH=  COH þ Agþ þ H2 O ¼  COOH þ Ag0 þ Hþ

(17)

Fig. 7. SEM photograph of (A) metal unloaded, (B) lead loaded and (C) silver loaded carbon. Error bars indicate standard deviation (n = 3).

processes involving the reduction of Ag+ to elemental silver are shown in Eqs. (15)–(17). In the processes of ion exchange or reduction, the molar ratio of [H+] released in solution to [Ag+] sorbed was 1.

Due to the strong effect of Pb2+ on Ag+ sorption in the binary mixture showing positive 4G8 values for the latter, it seems that Ag+ probably approaches the carbon surface via ion exchange followed by chemical reduction in reduction sites. Pb2+ ions competes for the ion exchange sites preventing to a great extent the approach of Ag+ ions to the carbon surface and, thus, the extent of sorption-reduction decreased. The capability of the dehydrated

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