selective removal of hg(ii) from aqueous solution by functionalized magneticmacromolecular hybrid...
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Selective removal of Hg(II) from aqueous solution by functionalized magnetic-macromolecular hybrid material
Khalid Z. Elwakeel, Eric Guibal
PII: S1385-8947(15)00798-6DOI: http://dx.doi.org/10.1016/j.cej.2015.05.110Reference: CEJ 13751
To appear in: Chemical Engineering Journal
Received Date: 17 April 2015Revised Date: 28 May 2015Accepted Date: 29 May 2015
Please cite this article as: K.Z. Elwakeel, E. Guibal, Selective removal of Hg(II) from aqueous solution byfunctionalized magnetic-macromolecular hybrid material, Chemical Engineering Journal (2015), doi: http://dx.doi.org/10.1016/j.cej.2015.05.110
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Selective removal of Hg(II) from aqueous solution by functionalized
magnetic-macromolecular hybrid material
Khalid Z. Elwakeela, Eric Guibal
b
a Environmental Science Department, Faculty of Science, Port-Said University, Port-Said, Egypt
b Ecole des mines d’Alès, Centre des Matériaux des Mines d’Alès (C2MA), 6, avenue de Clavières,
F-30319 Alès cedex, France
Abstract:
A new hybrid material was prepared using chitosan, glycidyl methacrylate and magnetite
microparticles. The concentration of amine groups on the sorbent was increased by grafting
diethylenetriamine. These materials were tested for the sorption of Hg(II) metal ions. These
materials showed high affinity and selectivity for Hg(II) uptake from aqueous solutions:
maximum sorption capacity reached 2.6 mmol g−1
at pH 4.0. The influence of pH was tested
on the selective separation of Hg(II) from a mixture of Hg(II), Co(II), Cu(II), Fe(II), Ni(II),
Zn(II) and Mg(II). The effect of counter anions was also studied at different pH values.
Sorption may occur on amino groups by chelation of cationic mercury species on amino
groups or by ion exchange of chloroanionic mercury species on protonated amino groups.
Uptake kinetics and sorption isotherms were modeled using the pseudo-second order rate
equation and the Langmuir equation, respectively. The distribution coefficient was obtained at
different temperatures and the thermodynamic parameters have been calculated: the sorption
is endothermic, spontaneous and contributes to increase the entropy of the system. Potassium
iodide was used for Hg(II) desorption from loaded sorbents: desorption yields 99%, and the
sorbent could be efficiently recycled for a minimum of three sorption/desorption cycles.
Keywords: Hg(II) sorption; magnetic hybrid sorbent; diethylenetriamine-grafted chitosan;
uptake kinetics; sorption isotherms; thermodynamics.
1. Introduction
Corresponding author: Fax: +33(0) 466782701 – Phone: +33 (0)466782734
E-mail address: [email protected] (E. Guibal).
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The regulations at international level are becoming more and more drastic concerning the
discharge of heavy metal ions into the environment. Therefore controls on industrial
wastewater are stressed because of their potential threat on human health and on the
biosphere. In addition, due to the rarefaction of some metal resources (critical, strategic and
precious metals, for example) strongly incentive politics of material reuse and recycling have
been recently elaborated.
The Agency for Toxic Substances and Disease Registry ranked mercury as the third priority
hazardous substance, after arsenic and lead [1]. Water pollution by Hg(II) is thus a serious
environmental issue, though many national or international directives are recommending as
far as possible substituting other metal to mercury or changing the production processes to
avoid using mercury . This metal is non-biodegradable and can threat the ecosystem by being
accumulated in the food chain; the dramatic Minamata intoxication of local population clearly
illustrates this hazardous effect. Hg(II) may be discharged into the environment through
wastewaters issued from different industries such as: electroplating, leather tanning, metal
finishing and petroleum refining.
Hg(II) removal from contaminated water is imperative to make water fit for the human uses
(drinking, agriculture, etc.). The allowed concentrations for discharge into surface water and
for drinking water are 10 µg/L and 1 µg/L, respectively [2]. The aforementioned threats of
mercury require developing treatment processes capable to face, alone or in combination with
other processes, the target levels for discharge into the environment. Generally, metal ions can
be recovered from solutions through conventional processes such as precipitation, solvent
extraction, membrane techniques, ion-exchange and chelating resins. Frequently these
techniques fail to fit with target regulations (technical limitations), or economic constraints
(solvent extraction is not appropriate for dilute effluents; membrane processes may be
expensive for large-scale applications) or produce huge amounts of contaminated sub-
3
products (with potential complementary hazards for the environment, as it may occur with
precipitation processes). Sorption is one of the preferred methods for the removal of metal
ions from the wastewater as it is effective as well as cost-effective [3]. In the field of sorption
processes, several alternative materials have been investigated for the last decades making
profit of materials of biological origin for metal binding: biosorption consists in using
functional groups held on materials of biological origin for metal recovery with mechanisms
similar to those found in ion-exchange and chelating resins. These materials can be issued
from agriculture, fishing or as by-product of other industries : fungal, algal, bacterial biomass,
agriculture or food industry residues, for example. Chitosan ((1,4)-2-amino-2-deoxy-β-D-
glucan) is an emblematic biopolymer that was abundantly studied for metal binding. Being
produced at the commercial scale from the shells of crustacean, it is obtained by alkaline
deacetylation from chitin (the primary resource) and bears many hydroxyl groups (which
bring hydrophilic behavior to the biopolymer, and also possible sites for chemical
modification and grafting) and amino groups. This biopolymer has unique properties among
polysaccharide having cationic behavior in acid solutions. Amino groups are involved in
metal binding through different mechanisms including (a) metal cation chelation in near
neutral solutions (through the free electron doublet of nitrogen), and (b) binding of metal
complex anion by ion-exchange/electrostatic attraction on protonated amino groups (in acidic
solutions). With a pKa that depends on the degree of acetylation (incomplete deacetylation of
chitin) and the degree of neutralization but ranges between 6.3-6.7 for most common
commercial samples [4] the protonation of chitosan leads, in most cases (with the remarkable
exception of sulfuric acid) to the dissolving of chitosan in acid media. This is the base of the
mechanisms used for changing the conditioning of the biopolymer but this is also a serious
drawback for potential applications in the field of metal sorption. A cross-linking treatment or
a chemical modification is thus required for extending the pH field of application (for both
4
sorption and desorption steps). On the other hand, Glycidyl methacrylate (GMA) is an
attractive vinyl monomer because of its low toxicity, lower cost compared with other acrylic
monomers, versatile properties and especially due to the presence of epoxy group in its
molecule [5], which makes it very reactive for chemical modification or for reacting with
other materials. These polymers are potent sorbents for various transition metals [6, 7].
Various studies on metal sorption on chitosan-based materials have pointed out the limitations
in uptake kinetics due to the resistance to intraparticle diffusion (which may be explained by
the residual crystallinity of the biopolymer and by poor porous and surface properties). This
drawback was partially solved changing the conditioning of the biopolymer (manufacturing
gel beads, membranes and so on). Another possibility consists in reducing the size of sorbent
particles but at the expense of serious drawbacks concerning particle recovery at the end of
the process (in batch systems) or head loss (in fixed-bed applications). Recently, a number of
studies have developed an alternative for solving the problem of filtration or recovery of fine
(micro- or nano) particles: the incorporation of a magnetic core (or its in situ production, with
simultaneous coating with biopolymer) in the particle makes easier the handling and recovery
of these fine sorbent particles. These methods are cheap and often highly scalable. Moreover,
these techniques employing external magnetic fields are more amenable to automation [8].
The reactivity of the sorbent can be improved by grafting new functional groups such as
polyamine compounds [9-13]. The objective may consist in increasing the density of sorption
sites, in improving the sorption selectivity, in changing the sorption mechanism (ion-
exchange vs. chelation or the reciprocal), or in extending the field of application (especially
regarding pH range).
The objective of this work consists in synthesizing a magnetic composite made of glycidyl
methacrylate (GMA) and chitosan (coating the magnetic Fe3O4 core) that is functionalized by
grafting diethylenetriamine (DETA) through GMA (to increase the density of sorption sites
5
(and their potential selectivity). These composite materials are being tested for Hg(II)
recovery through the study of pH effect, the investigation of competition effects (composition
of the matrix: competitor metal ions, counter anions), the determination of sorption isotherms
and thermodynamic characteristics and the identification of controlling steps in uptake
kinetics. Finally the desorption of Hg(II) is studied with the objective of verifying the
possibility to recycle the sorbent.
2. Materials and methods
2.1. Materials
All chemicals used were of analytical grade and demineralized water was used for the
preparation of all aqueous solutions. Glycidyl methacrylate (GMA) was provided by Riedel-
de Haën (Germany), while N,N’ methylenebisacrylamide (MBA) and benzoyl peroxide
(Bz2O2) were supplied by Fluka AG (Switzerland). Chitosan, epichlorohydrin
(chloromethyloxirane) and diethylenetriamine (DETA) were obtained Sigma-Aldrich
(Switzerland) while isopropyl alcohol was provided by Carlo Erba (France). Chitosan was
provided by Aber Technologies (France) ; the deacetylation degree of the biopolymer was 87
ù and the molecular weight (MWn) was 125000 g mol-1
. All other chemicals were Prolabo
products (France) and were used as received. HgCl2 salt was used for the preparation of the
stock solution, except in the case of the study of counter anions (for which the relevant salt of
mercury was used).
Fe3O4 particles were synthesized by co-precipitation of ferric and ferrous salts following a
procedure derived from the so-called Massart method [14]. The amounts of 6.480 g of
FeCl3·6H2O and 3.334 g of FeSO4·7H2O were dissolved into 150 mL of demineralized water.
Then 2 mL of HCl was added under continuous stirring for 30 min, until complete dissolving
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of precursors salts. The chemical precipitation was achieved at 30 °C under vigorous stirring
by drop wise addition of 10 mL of NaOH solution (50%, v/v). During the reaction process,
pH was maintained around pH 11. The reaction system was maintained under agitation at 30
°C for 15 min. The product was then washed with demineralized water several times and
stored for use.
2.2. Preparation of the magnetic-macromolecular hybrid material
Step 1: Magnetic chitosan particles were prepared by the co-precipitation method, 4 g
chitosan powder was dissolved in a volume of 150 mL of acetic acid solution (10 %, w/v).
The magnetite precursors were mixed with 150 mL of demineralized water (i.e., 6.479 g of
FeCl3·6H2O and 3.334 g of FeSO4·7H2O) using a volume of 2 mL of HCl (concentrated) for
improving salt dissolution under continuous stirring for 30 min. After complete dissolution of
iron salts, the solution was mixed with the chitosan solution. Then the chemical precipitation
of magnetic chitosan material was performed at 30 °C under vigorous stirring by dropwise
addition of 20 mL of NaOH solution (50 %, v/v) for 30 min. The pH was roughly controlled
around pH 11. The solid product that was recovered by centrifugation was freeze-dried using
a freeze-dryer (Bioblock scientific, Christ) at 223 K and 0.01 mbar; this intermediary product
was called MC (magnetic chitosan).
Step 2: The magnetic-chitosan-glycidyl-methacrylate macromolecular hybrid material was
prepared through the polymerization of GMA in the presence of MC particles (Step 1). GMA
(5.0 g) corresponding to a mass ratio of 50 % (w/w, referred to MC amount) was used for
GMA grafting; 0.2 g of MBA was used as the cross-linking agent and 0.1 g of Bz2O (acting as
the initiator of the polymerization reaction) was added under agitation before adding 1.0 g of
magnetite. Three mL of isopropyl alcohol and 25 mL of cyclohexane were mixed and added
to the former solution. The solution was then poured into a flask containing 100 mL (1 %)
7
polyvinyl alcohol and heated on a water bath at 75–80 C under continuous stirring for 3 h.
The product was filtered off and washed repeatedly with demineralized water and acetone
before being air-dried. The product was called MCGMA. The preparation and modification of
magnetic-macromolecular hybrid material is described in Scheme 1, and the suggested
chemical structure of MCGMA is shown in Scheme 2.
2.3. Chemical modification of MCGMA
2.3.1 Diethylenetriamine grafting – First method
MCGMA was suspended in dioxane (100 mL) and then treated with 5 mL of
diethylenetriamine (DETA). The suspension was stirred at 70 °C for 12 h. The product was
successively washed with water and acetone. The product was then air-dried and called
MCGMA-I.
2.3.2 Diethylenetriamine grafting – Second method (via epichlorhydrin)
MCGMA was suspended in 70 mL of isopropyl alcohol. Then 7 mL of epichlorohydrin (62.5
mmol) dissolved in 100 mL acetone/water mixture (1:1 v/v) was added to the suspension. The
reaction was performed under continuous stirring for 24 h at 60 °C. The solid product was
filtered off and washed several times with water. The obtained product was treated with
DETA according to the method reported in section 2.3.1. The product was finally air-dried; it
was sieved and the fraction below 1 mm was retained for experiments. This sorbent was
called MCGMA-II.
2.3.3. Estimation of the amine content
The amine content in the obtained resin was estimated using a volumetric method previously
described by Atia et al. [15]. Fifty mL of HCl (0.05 M) was added to 0.5 g of MCGMA-II or
MCGMA-II and conditioned for 15 h on a shaker at 20 °C. The residual concentration of HCl
8
was measured by the titration against 0.05 M standardized NaOH (using phenolphtalein as the
titration indicator). The number of moles of HCl that reacted with the amino groups was
determined and the concentration of amino group was calculated by the following equation
resin) (mmol/g 50
50212
.
)M(Mgroups)(NHofionConcentrat
(1)
where M1 and M2 are the initial and final concentrations of HCl.
2.3. Analytical methods
Mercury was analyzed using inductively coupled plasma atomic emission spectrometer ICP-
AES (Jobin-Yvon Activa M, Horiba-Jobin Yvon, FRANCE). The morphology and the
elemental distribution of Hg in the magnetic-macromolecular material were analyzed with a
Scanning Electron Microscope coupled with an Energy Dispersive X-ray analysis system
(SEM-EDX; Environmental Scanning Electron Microscope (ESEM) Quanta FEG 200,
equipped with an OXFORD Inca 350 Energy Dis-persive X-ray microanalysis (EDX)
system). The magnetic property of the sorbent was measured on a vibrating-sample
magnetometer (VSM) (Lake Shore 730T, Westerville OH, USA) at room temperature.
2.4. Sorption experiments
Demineralized water was used for the preparation of metal ion solutions. A stock solution (20
mM) of HgCl2 was prepared in demineralized water. The other solutions were obtained by
dilution of the stock solution with demineralized water just prior experiments. HCl (0.5 M)
and NaOH (0.5 M) were used to control the pH of the solutions. 0.1 M of KI was used for
elution of Hg(II) loaded on the sorbent.
Sorption experiments were performed in batch systems using polyethylene flasks and the
temperature was set to 20 ± 1 °C (unless specified). For the study of pH effect 20 mL of 10
9
mM Hg(II) solutions at different pH values (in the range 1-6) were mixed with 50 mg of
sorbent (dried weight) for 5 h, and the stirring speed was maintained at 100 rpm using a
reciprocal agitator Rotabit, J.P. Selecta (Spain). Samples were collected and filtrated through
1.0 µm pore size filtration membrane and/or magnetic separation and the filtrate was analyzed
for residual Hg(II) concentration using ICP-AES. The pH was not controlled during the
sorption but the final pH was systematically recorded.
For sorption isotherms 50 mg of sorbent (m) were mixed with 20 mL (V) of Hg(II) solutions
at different initial concentrations (C0, ranging between 3 and 10 mmol Hg L-1
) for 5 h. The pH
of the solutions was initially set at 4. After solid/liquid separation, the residual concentration
(Ceq, mmol Hg L-1
) was determined by ICP-AES and the sorption capacity (qeq, mmol g−1
)
was determined by the mass balance equation: qeq = (C0-Ceq)V/m.
For uptake kinetics 300 mg of sorbent were mixed with 120 mL of Hg(II) solutions (C0: 10
mmol Hg L-1
) at pH 4. Samples (a volume of 4 mL) were collected (the sorbent was
magnetically separated) at fixed times and the residual concentrations were determined by
ICP-AES. The agitation speed was set at 100 rpm while the temperature was maintained at 20
± 1 °C. The sorbed amount of Hg(II) per unit weight of the sorbent at time t (q(t), mmol Hg g-
1), was calculated from the mass balance equation (taking into account the decrement in the
volume of the solution) as:
n
im
itVitCitCtq
1
)1()())()()1()(()( (2)
where C(t)(i) (mmol Hg L-1
) is the Hg(II) concentration of the withdrawn sample number i at
time t and C(t)(0) = C0, V(t)(i) (mL) is the volume of the solution in the flask at sample number
i and time t, and m is the amount of resin added into the flask. Here V(t)(i)-V(t)(i-1) equals 4
mL.
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Regeneration experiments were performed by contact of 250 mg of the sorbent with 100 mL
of 10 mmol Hg L-1
solution at pH 4 Hg(II) for 5 h. The amount of metal sorbed (and the
sorption capacity) was determined by the mass balance equation. The solution was
magnetically decanted and the sorbent was washed by distilled water. The loaded sorbent was
mixed with 100 mL of 100 mM KI for 50 min. Samples were collected at different time
intervals and the residual concentration of Hg(II) was determined by ICP-AES after
solid/liquid separation. The regenerated sorbent was carefully washed by distilled water for
reuse in the second run. The regeneration efficiency (RE, %) was calculated according to the
following equation:
100 X % RE1)-(nrun at (mmol) Hg(II) desorbed ofAmount
(n)run at (mmol) Hg(II) desorbed ofAmount (3)
3. Results and discussion
3.1. SEM and SEM-EDX analysis
Figure 1 shows the SEM and SEM-EDX analyses of the sorbents MCGMA-I and MCGMA-II
before and after Hg(II) sorption. The surface of MCGMA-I and MCGMA-II is generally
smooth with several small pores inside and interconnected macropores can be identified
within the sorbents (Figure 1-a,b). However, after Hg(II) sorption, differences in the surface
structure of the sorbents clearly appear: agglomerated spherical particles are formed on the
surface and the porous structure disappears (Figure 1-c,d). These conclusions are consistent
with observations made on SEM-EDX analysis (Figure 2): the typical signals of Hg element
are appearing at 1.7, 2.2, 2.6 and 10 keV on both MCGMA-I and MCGMA-II sorbents.
Though a quantitative analysis of the Hg load on the sorbent by SEM-EDX would not be
accurate, the levels of Hg are high enough to demonstrate qualitatively the homogeneous
11
distribution of the metal on the surface of the sorbent. The mass percentages of Hg element
are found close to 27.5 % and 32.9 % for MCGMA-I and MCGMA-II, respectively. This is a
first evidence of the good affinity of the sorbents for mercury. The presence of Cl element is
due to the binding of mercury under the form of chloroanionic species: the sample used for
analysis was collected after metal sorption at pH 3 (controlled with HCl).
Figure AM1 (See Additional Material Section) shows the SEM images of the two sorbents at
different magnitudes. There is a large dispersion in the size of sorbent particles (in the range
10-200 µm) and the particles are characterized as irregularly-shaped flakes.
3.2. Amine content
The amine content in MCGMA, MCGMA-I and MCGMA-II (which was determined by
titration) are 3.12, 5.92 and 6.5 mmol g−1
, respectively. The chitosan used in this study was
characterized by a deacetylation degree of 87 % (this means 6 mmol N g-1
). The magnetite
fraction in the sorbents was characterized by weight loss at 700 °C: the materials contain 36
% and 34 % of magnetite for MCGMA-I and MCGMA-II, respectively. On the basis of
similar magnetite fraction and chitosan (without chemical modification) the theoretical
nitrogen content would be close to 3.9 mmol N g-1
. This confirms the efficient grafting of
additional amino groups on the reference material (MCMA). The semi-quantitative analysis
of nitrogen content obtained by SEM-EDX analysis confirms this trend. Indeed, the nitrogen
mass percentages were close to 8.6 % and 9.1 % for MCGMA-I and MCGMA-II, respectively
(this means 6.15 mmol N g-1
and 6.5 mmol N g-1
, respectively). These observations and
analyses confirm the successful experimental procedure described in Scheme 1 for the
grafting of additional amino group on chitosan backbone.
3.3. Magnetic properties
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The magnetic properties of the materials were determined using VSM (vibrating sample
magnetometry). Figure 3 shows their typical magnetization loop. There was no remanence
and coercivity, contrary to certain supported-magnetite materials [16]. This means that
MCGMA sorbents are paramagnetic. The saturation magnetization of MCGMA-I and
MCGMA-II were found to be about 19.1 and 17.9 emu g-1
, respectively. These values are
much smaller than the levels reported for bulk phase magnetite (i.e., 92 emu g-1
) and also
smaller than the values obtained for magnetite nanoparticles [17]. The slight decrease in the
magnetization in case of MCGMA-II may be due to the grafting of DETA via epichlorhydrin,
which increase the mass of organic material and decrease the relative percentage of magnetite
in the sorbent. The reported decrease in saturation magnetization can be explained by several
factors including size effect and particle crystallization [17], and obviously by the fact that
only a fraction of the sorbent (about 50 %) is constituted by the magnetic core. Similar
decrease in magnetization was observed for other chitosan-magnetite composites [18].
Therefore, MCGMA sorbents can be readily separated with the help of an external magnetic
field. This may be very helpful for solid/phase separation or for handling the material in
hazardous environment.
3.4. Infra-red spectrometry
FT-IR spectrometry was used to characterize the structure of the sorbents (Figure 4). The two
sorbents have very similar FT-IR profiles: the same bands are appearing on the two spectra,
the only small differences observed are small shifts in some of these bands. The large band
around 3410 cm-1
is generally attributed to the combination of stretching vibrations of OH and
NH2 groups, while a series of –CH vibrations are reported around 2940 cm-1
and 2850-2860
cm-1
. These bands are present in a wide number of compounds and they are bringing poor
information on the specific properties of the sorbents. More interesting and representative are
13
the different bands identified in the range 1800-400 cm-1
. The bands at 1264, 1162 and
986/990 cm-1
are representative of the saccharide ring. The band at 1653-1657 cm-1
is
associated to the CONH2 group [19], while the band at 1561-1573 cm-1
is probably assigned
to the in-plane bending vibration of free amine groups and/or the asymmetric carboxylate
stretching vibration [19, 20]. The grafting of diethylenetriamine, via glycidylmethacrylate,
can be identified by the strong peak observed around 1724 cm-1
(assigned to C=O stretching
vibration) [20, 21]. Additional methylene (-CH2-) groups can probably be associated to the
band observed at 1466-1470 cm-1
[22]. The band at 586-596 cm-1
was attributed to γ-Fe2O3
[23].
3.5. Hg(II) sorption properties
3.5.1. Effect of pH
The pH of the aqueous solution is a key parameter that controls the sorption of the Hg(II).
This effect can be associated to metal speciation, to chemical properties of the polymer
(protonation/deprotonation of amino groups). Figure 5 shows the effect of pH on Hg(II)
sorption (varying the initial pH from 0.9 to 6), and the pH variation observed during sorption
operation. Figure 5a confirms that the pH influences sorption capacity: the sorption increases
with the pH up to pH 4-4.2 before beginning to stable and even reducing when pH exceeds 5.
It is remarkable that the pH variation during metal sorption is not varying significantly
(Figure 5b): at low pH (below pH 4) a slight increase of the equilibrium (below 0.4 pH unit)
is observed, while above pH 4-4.5 the equilibrium pH tended to slightly decrease (variation
being less than 0.5 pH unit). This is remarkably stable compared to other chitosan-based
systems where a kind of “buffering effect” is obtained around 5-6 due to the acid-base
properties of the biopolymer (pKa in the range 6.3-6.7) that binds protons. It is suspected that
14
the grafting of new functional groups (including other amino groups with different acid-base
properties, associated to primary and secondary amino groups) contribute to smooth this pH
variation effect.
As the pH increases the protonation of amine groups decreases, which, in turn, contributes to
make free the amino groups for interacting with metal cations by chelation. For MCGMA-II
sorbent the sorption capacity increases from pH 1 to pH 4 and then tends to stabilize. At
higher pH (above pH 5) the formation of hydrolyzed species (Hg(OH)+ or soluble Hg(OH)2,
before precipitation) that have less affinity for amino groups, may explain the slight decrease
in sorption capacity (which represents less than 0.2 mmol Hg g-1
). This effect is completed by
the impact of deprotonation of amine groups. On the other hand, at low pH (below 2-3), the
sorption capacity remains quite high (up to 1.5 mmol Hg g-1
), especially compared to the
levels reached with pure chitosan [24]. This can be explained by either (a) differences in the
acid-base properties of amino groups grafted on chitosan (lower pKa compared to chitosan
amino groups), and/or (b) the binding of chloro-anionic mercury species on protonated amino
groups. Indeed, in the presence of HCl, at low pH, metal speciation can be displaced to the
formation of HgCl42-
or HgCl3- species which may be bound to protonated amino groups
through an electrostatic attraction mechanism of an ion-exchange mechanism: chloride ions
bound to protonated amino groups may be exchanged with anionic chloroanions. The sorbent
MCGMA-I has a slightly different behavior: the sorption capacity tends to increase with pH
increasing from 0.9 to 2.5 and progressively decreases when pH increases: the acid-base
properties of the different amino groups present on the sorbent are different to those on the
MCGMA-II sorbent due to differences in their chemical environment. The decrease in amino
group protonation at pH above 2.5 (which contributed to bind anionic mercury chloroanions)
is not compensated by the chelation properties of amino groups and by the speciation of
mercury in solution (metal speciation is shifted from chloroanions to positively charged
15
mercury species with low affinity for protonated amino groups). The two sorbents have
comparable sorption capacities at 2.3: below this value MCGMA-I is more efficient for Hg(II)
binding while above this value the best results were obtained by MCGMA-II sorbent.
Above pH 5 the formation of colloidal mercury species (Hg(OH)2) may occur and their
precipitation may overestimate sorption performance [25]. This phenomenon should be taken
into account. For further experiments the pH was set to 4-4.2.
3.5.2. Kinetics
The uptake kinetics of Hg(II) using MCGMA-I and MCGCMA-II is shown in Figure 6.
Figure 6a shows the relative decay of concentration with time, which allows comparing the
two sorbents both in terms of equilibrium performance and relative decay slopes. Figure 6b
shows the evolution of the concentration of Hg(II) in the sorbent (i.e., q(t) vs. time): the
approach of equilibrium is another method for comparing the kinetic profiles. Experimental
conditions have been selected to be able to illustrate the effect of resistance to diffusion.
Indeed, an excess of sorbent (compared to metal) would lead to fast and complete removal of
the metal bound on the sorption sites at the surface of the sorbent and minimizing the impact
of mass transfer mechanisms. Under these conditions, the solutions is not fully
decontaminated (Figure 6a); this issue (complete metal recovery) is addressed below (see
section 3.5.7)
The kinetic profiles (both in terms of relative concentration decay; i.e., Figure 6a, and metal
concentration in the sorbent; i.e., Figure 6b) are characterized by two main phases in the
uptake: (a) a first initial step that lasts for about 60 min and counts for more than 60 % of total
sorption, (a) a second step that takes about 6 hours and corresponds to a much slower metal
accumulation. The initial section of the curve corresponds to a great availability of reactive
groups (surface coverage is progressively increasing) and a large concentration gradient
16
between the solution and both the surface and the internal sorption sites. These two conditions
may explain the fast initial accumulation of Hg(II). The sorption mainly occurs on the
reactive groups covering the surface of the sorbent. The second step, much slower, is
controlled by the decrease of the concentration gradient and by the resistance to intraparticle
diffusion and requires much longer time for reaching the equilibrium (i.e., about 6 h).
Actually the binding kinetics is controlled by a series of mechanisms including: (a) the bulk
diffusion, (b) the resistance to film diffusion (or external diffusion), (c) the resistance to
intraparticle diffusion, and (d) the proper reaction rate (chemical reaction rate) [26]. Usually a
sufficient agitation allows neglecting the resistance to bulk diffusion and minimizes the
resistance to film diffusion whose contribution in the control of uptake kinetics is mainly
significant within the first minutes of contact. The modeling of such a complex system, which
is beginning even more complex when the sorbent is heterogeneous, when the solution is
subject to changes in the speciation of the metal ions, etc., requires complex numerical tools.
Experimental data have been modeled using simplified conventional equations to fit kinetic
profiles and make possible the comparison of kinetic parameters for the two sorbents. Hence,
the kinetics of Hg(II) sorption on modified MCGMA sorbents were analyzed using the
pseudo-first order rate equation (PFORE) [27], the pseudo-second order rate equation
(PSORE) [28], the simplified resistance to intraparticle diffusion equation [29] and the
Elovich equation [30]. These models and their linear forms are reported in Table 1, where k1
is the pseudo-first order rate constant (min-1
) of sorption and qeq and q(t) (mmol Hg g−1
) are
the amounts of Hg(II) sorbed at equilibrium and time t, respectively, k2 is the pseudo-second
order rate constant (g mmol−1
min−1
), Ki is the intraparticle diffusion rate (mmol g−1
min-0.5
), α
the initial sorption rate (mmol g−1
min−1
) and β the desorption constant (g mmol−1
). The
validity of each model is checked by the correlation coefficient associated to the linear fits.
Table 2 reports the parameters of the different models for both MCGMA-I and MCGMA-II
17
sorbents. Systematically, the best correlation coefficients were found for the PSORE model;
this is confirmed by the plot of experimental data according the linearized forms of these
models: Figure AM2a and Figure AM2b (see Additional Material Section) for PFORE and
PSORE, respectively, show a best fit of kinetic profiles by PSORE. In addition, the
comparison of equilibrium sorption capacities for the calculated values and the experimental
values are only consistent for the PSORE model: the equilibrium sorption capacities are found
close to 1.4 and 2.0 mmol Hg(II) g-1
for MCGMA-I and MCGMA-II, respectively. PSORE
modeling gave values of 1.49 and 2.09 mmol Hg(II) g-1
closer from experimental values than
PFORE (0.91 and 1.27 mmol Hg g-1
, respectively). It was more likely to reflect that the rate-
determining step might be chemical sorption and that the sorption behavior might involve a
chelation mechanism through coordination between Hg(II) ions and the reactive groups of the
sorbents [25]. The Hg(II) ions may form complexes with amino groups of the modified
MCGMA sorbents (with possible contribution of hydroxyl groups in the stabilization of
metal binding) [31].
However, the PSORE describes kinetics data through a global approach, and does not take
into account the contribution of diffusion mechanisms in the control of the kinetics. Under
these conditions, the kinetic parameters should be considered as apparent rate coefficients.
The influence of resistance to intraparticle diffusion has been approached using a simplified
model: the so-called Weber and Morris plot (Table 1). The intraparticle diffusion model
provides a more comprehensive approach for defining of sorption mechanism, and the plot
generally allows identifying different successive steps in the global process [32]. The Weber
and Morris shows multi-linear sections (Figure AM3a, see Additional Material Section), i.e.,
three linear sections (on the plot q(t) vs. t0.5
) with fast kinetics in first step followed by the
gradual attainment of equilibrium for both sorbents, and a pseudo saturation plateau. The
multi-linear plot does not pass through the origin suggesting that the resistance to intraparticle
18
diffusion is not the sole rate-limiting step: other steps, e.g. resistance to film diffusion and/or
reaction rate, are probably involved in the control of uptake kinetics (Table 2, Figure AM3a,
see Additional Material Section). We can assume that the first linear section corresponds to a
regime controlled by the resistance to film diffusion and that the binding is limited in this
stage at the sorption of external sorption sites or on the macropores in the first external layers
of the material. The second section is characterized by a much lower kinetic rate and leads to
a slow approach to equilibrium with the control by the resistance to intraparticle diffusion
(into internal macroporous and mesoporous network). The last step is very slow and
represents only a few percentage of the total sorption: this phase can be associated to the
resistance to diffusion in the microporous network of the sorbent. In addition the progressive
saturation of available and accessible sorption sites influences the local equilibrium on the
surface between surface sorption and desorption. The low values of intraparticle rate constants
(ki) appearing in Table 2 with values available in the literature indicates that the sorbents
(both MCGMA-I and MCGAM-II) are significantly affected by the resistance to intraparticle
diffusion, at least compared to the fast reaction rate of chemical sorption of Hg(II) ions at the
surface of the sorbent [33]. The analysis of the rate constants for the Weber and Morris model
shows that except for the first step, the parameters were systematically higher for MCGMA-II
compared with MCGMA-I. In the initial stage, MCGMA-I seems to be faster in sorption than
MCGMA-II: this first stage corresponds to a kinetic control by external film diffusion and/or
diffusion in the macropores present at the external surface of sorbent particles. While in the
next steps the uptake kinetics is controlled by the diffusion in meso- and micro-pores, which
appears to be a little more accessible for MCGMA-II material.
The Elovich equation was developed for modeling chemisorption processes [30]. Table 1
displays the model equation and its linearization as well the plots to be used for determining
the parameters (which are reported in Table 2). The values of α and β were determined from
19
the intercept and slope, respectively, of the linear plot of qt vs ln t (Figure AM3b, see
Additional Material Section). The values of α for the sorption of Hg(II) ions on the modified
sorbents are 0.46 and 0.73 (mmol g−1
min−1
) for MCGMA-I and MCGMA-II, respectively.
These values are higher than the values cited in literature [34, 35], which may be attributed to
the high concentration of active sites on the sorbent surface allowed for reacting with Hg(II)
ions. The values of β (desorption constant) are found to be 4.54 and 3.54 g mmol−1
for
MCGMA-I and MCGMA-II, respectively. These values are smaller compared with the levels
cited in the literature [34, 35]. This is another confirmation of the high affinity of the sorbents
for Hg(II) ions.
3.5.3. Equilibrium sorption isotherm
The sorption isotherms represent the distribution of the sorbate between the solid (sorption
capacity or sorbate concentration in the solid) and liquid phase (residual concentration of the
sorbate in the solution), at equilibrium. This distribution is only controlled by the temperature
and is independent on the size of sorbent particles, on the experimental procedure (batch,
fixed-bed column) providing a sufficient time has been given to the process of reaching
equilibrium. The distribution of the sorbate between the two phases can be modeled using
different equations; however, the most frequents models used for describing solid/liquid
equilibrium systems have been established by Langmuir [36], Freundlich [37], and Dubinin-
Radushkevich [38]. These models and their linear forms are reported in Table 3, where qeq the
sorbed value of Hg(II) at equilibrium concentration (mmol Hg g−1
), qm,L is the maximum
sorption capacity (corresponding to the saturation of the monolayer, mmol Hg g−1
) and KL is
the Langmuir binding constant which is related to the energy of sorption (L mmol−1
), Ceq is
the equilibrium concentration of Hg(II) in solution (mmol Hg L−1
). KF (mmol g−1
) (L
mmol−1
)1/n
and n are the Freundlich constants related to the sorption capacity and intensity,
20
respectively. KDR (J2 mol
−2) is a constant related to the sorption energy, qDR (mmol g
−1) is the
theoretical saturation capacity, ε (J2 mol
−2) is the Polanyi potential.
At low initial concentrations (1.0–3.0 mM) the sorption of Hg(II) was almost quantitative.
Brunauer et al. [39] divided the sorption isotherms into five types. Type I isotherm represents
unimolecular sorption and applies to non-porous, microporous and sorbents with small pore
sizes (not significantly greater than the molecular diameter of the sorbate). The isotherm
curves (Figure 7) are following the typical shape of I-type systems, according to the BET
classification [39]: they are characterized by a high degree of sorption at low concentrations.
At higher concentrations, the sorption sites becoming progressively occupied the sorption
tends to stabilize and forms the saturation plateau.
The Langmuir model is the simplest theoretical model that is used for describing monolayer
sorption onto a surface with a finite number of identical sites. It was originally developed to
represent gaz/solid sorption before being extrapolated to liquid/solid sorption. The values of
qm,L and KL were determined from the slope and intercept, respectively, of the linear plot of
Ceq/qeq vs. Ceq (Figure AM4a, see Additional Material Section); parameters of the Langmuir
equation are reported in Table 3. The maximum sorption capacities (qm,L) are in good
agreement with the experimental values. The correlation coefficients reported in Table 4
confirm the better fit of experimental data by the Langmuir model compared with the other
two models. Though the fact that the mathematical fit of the sorption equation fits well
experimental data does not necessarily mean that the hypotheses associated to the model are
verified this is indicative of the relative homogeneity of the surface; or at least that the active
sites are energetically equivalent. Furthermore, the Langmuir parameters can be used to
predict the affinity between the sorbate and sorbent using the dimensionless separation factor
RL
o
LC1
1R
LK (4)
21
where KL is the Langmuir equilibrium constant and Co is the initial concentration of Hg(II)
ions. Values of 0 < RL < 1 indicates the “suitability” of the process. In this study, the values of
RL for the resin for the sorption of Hg(II) ions lie between 0.002 and 0.088 for each sorbents
and all concentrations at 20 oC. These values confirmed the affinity of modified MCGMA
sorbents for Hg(II) ions.
The Freundlich isotherm model is applicable to highly heterogeneous surfaces, and a sorption
isotherm lacking to form a saturation plateau indicates a multi-layer sorption. The values of n
and KF were determined from the slope and intercept, respectively, of the linear plot of ln qeq
vs. ln Ce (Figure AM4b, see Additional Material Section) and the calculated parameters of the
Freundlich equation are summarized in Table 4. Both the saturation plateau formed on Figure
7 at high residual mercury concentration and the relatively low correlation coefficients
confirm that the Freundlich poorly fit experimental data.
Another popular equation for the analysis of isotherm was proposed by Dubinin and
Radushkevich (Table 3). This isotherm was developed taking into account the effect of the
porous structure of the sorbent, and the energy involved in the sorption process. The Polanyi
potential ( ) given as Eq. (5) [38]:
) C
1 (1ln RT
eq
(5)
R is the universal gas constant (8.314 J mol−1
K−1
) and T is the absolute temperature (K). The
slope of the plot of ln qeq vs. 2 (Figure AM4c, see Additional Material Section) gives KDR
and the intercept yields QDR (Table 4). The D–R constant (KDR) can give valuable information
regarding the mean energy of sorption (Ea, J mol-1
) by Eq. (6):
5.0)2(
1
aE
DRK (6)
The results of D–R isotherm are reported in Table 4. Figure AM4c (See Additional Material
Section) shows that the experimental data are poorly fitted to the model (consistently with the
22
low value of correlation coefficients): the values of the parameters of the model should be
taken as indicative values (order of magnitude). The values of the mean energy of sorption
range between 8.7 and 10.3 kJ mol-1
: this is consistent with the proposed mechanism of
chemisorption. Indeed, it is generally admitted that 8 kJ mol-1
is the limit energy for
distinguishing physical (below 8 kJ mol-1
) and chemical sorption. A comparison of the
correlation coefficient values obtained from the Langmuir, Freundlich and D–R isotherm
models in Table 4 reveals that the correlation coefficients for the Langmuir isotherm are
somewhat higher than those for the Freundlich and D–R isotherm. This result suggests that
the binding of Hg(II) may occur as a monolayer on the surface of the sorbent and that the
uptake occurs on a homogenous surface by monolayer sorption. This should be confirmed by
experimental observation for confirmation (as reported above the mathematical fit of isotherm
curve by a given model equation does not mean that the relevant hypotheses of the model are
verified). The uptake can be described in terms of chemisorption as the formation of ionic or
covalent bonds between the sorbent (free mercury species or chloroanionic species, depending
on the pH and the sorption mechanism) and the sorbate (free amine groups or protonated
amino groups depending on the pH).
The coexistence of different metal species that could be sorbed in function of the pH as well
as the different mechanisms, and the simultaneous presence of different types of amino
groups are not comforting the hypothesis of homogeneous surface (or homogeneous energies
of sorption).
3.5.4 Influence of temperature
The sorption performance may be controlled by the temperature both in terms of kinetics
(depending on the activation energy) and equilibrium: the thermodynamics of the process
influences the instantaneous kinetics of sorption and desorption, and consequently the
23
equilibrium. The sorption capacities were compared at different stabilized temperatures
(under identical experimental conditions) for both MCGMA-I and MCGMA-II (Table 5): the
sorption capacity increases with temperature and the sorption is an endothermic process. The
sorption equilibrium constant, Kc was determined (Eq. 7) and used with the van’t Hoff
equation (Eq. 8) and conventional thermodynamic equation (Eq. 9) for evaluating the
thermodynamic constants of the sorbents (i.e., the standard enthalpy change, ∆Ho, the
standard free Gibbs energy, ∆Go, and the standard entropy change, ∆S
o).
eqC
eqqcK (7)
where qeq and Ceq are equilibrium concentrations of Hg(II) on the sorbent and in the solution,
respectively.
∆Go = -RT lnKc (8)
∆Go = ∆H
o − T∆S
o (9)
Therefore the van’t Hoff equation becomes:
R
S
RT
H- ln C
K (10)
The values of standard enthalpy change (∆Ho) and standard entropy change (∆S
o) for the
sorption process are thus determined from the slope and intercept of the plot of ln Kc versus
1/T: the values of thermodynamic parameters are reported in Table 6. The positive values of
∆Ho confirm the endothermic nature of sorption process. The negative values of ∆G
o indicate
that the sorption reaction is spontaneous. The increase in the negativity of ∆Go with
increasing temperature confirms that the “favorability” increases with temperature. In
addition, the values of standard free energy change for MCGMA-II are more negative than
these of MCGMA-I: Hg(II) sorption is more favorable on MCGMA-II than on MCGMA-I,
especially at low temperature. The density of sorption sites and the spatial arrangement of
24
functional groups are more favorable for Hg(II) binding on MCGMA-II sorbent. It is
noteworthy that in the case of MCGMA-I, the standard free Gibbs energy becomes positive
only at higher temperature (i.e., above 40 °C): the spontaneous nature of sorption process is
only possible when the system is “activated” by temperature. With MCGMA-II more amine
groups are available; in addition their distribution along a free long chain makes them more
accessible. This could explain the better “spontaneity” of sorption process for MCGMA-II
compared to MCGMA-I (which possibly requires an increase in temperature for the
enhancement of metal ion or polymer chains “mobility” that, in turn, improves sorption
performance).
In industry and water purification plants the optimum temperature at which the sorption is
highly feasible and spontaneous is essential. The sorption of metals onto sorbent surfaces may
be either spontaneous or non-spontaneous in function of temperature. The limit temperature
value corresponding to a null value of standard free energy can thus be deduced from Eq. 11.
The range of temperature can be predicted from the value of temperature at which the
standard free energy is zero (T0), and then the minimal temperature for the process to being
spontaneous.
S
HT0
(11)
Here, the calculated values of zero standard free energy temperature (T0) are 312.8 and 283.8
for MCGMA-I and MCGMA-II, respectively. The low T0 for MCGMA-II indicates the
feasibility of Hg(II) removal at ambient temperature, while MCGMA-I sorbent requires a
thermal activation: the spontaneous reaction occurs at around 40 °C.
3.5.5. Effect of counter anions and metal speciation
The effect of the counter anion on the uptake of Hg(II) by modified MCGMA sorbents was
investigated at different pH values using different mercury salts. The pH of the solution was
25
adjusted using HCl for HgCl2 solution, HNO3 for Hg(NO3)2 solution and H2SO4 for HgSO4
solution. Table 7 shows that at low pH (i.e., 0.95) an appreciable sorption was obtained in the
case of HgCl2 for both sorbents, while with Hg(NO3)2 and HgSO4 salts the sorption was
negligible. This result can be explained by the possibility for the sorbents to bind mercury by
an ion-exchange between the counter anions (chloride) bound to amino groups and the Hg(II)
chloroanions (HgCl3- or HgCl4
2-) [25]. In the case of other counter-anions, mercury does not
form stable anionic complexes that could be bound to protonated amine groups by ion-
exchange and the sorption strongly decreases. The limited uptake in the case of Hg(NO3)2 or
HgSO4 salts may be also attributed to the higher stability of (MCGMA-NH+)NO3
− or the
(MCGMA-NH+)2SO4
2− that retards the ion exchange mechanism. At pH 2.12 for both
MCGMA-I and MCGMA-II Hg(II) was significantly sorbed when using HgCl2 or Hg(NO3)2
salts, while the sorption was negligible with HgSO4 salt. Sulfate salts have more affinity for
protonated amine groups than other mono-anions making more difficult metal anion
exchange. The composition of the solution may thus have a significant impact on the sorption
performance, especially at pH where the ion-exchange is the predominant sorption
mechanism.
3.5.6. Selectivity of modified MCGMA for Hg(II) sorption
In order to evaluate the possible selectivity of the sorbents Hg(II) sorption was carried out in
multicomponent solutions (containing Hg(II), Co(II), Cu(II), Fe(II), Ni(II), Zn(II) and Mg(II))
at different pH values (0.4 -1.6). Table 8 shows that the recovery of Hg(II) is almost
quantitative in the presence of equimolar concentrations of other base metal cations,
regardless of the pH (at least in the selected pH range). Whatever the pH, the maximum
sorption capacities of competitor metal ions by the modified MCGMA sorbents is less than
5.0 % of sorbed amounts of Hg(II) ions. The sorbents appears to have a selective recognition
26
pattern for Hg(II) ions against other selected heavy metal ions. The reported negative values
for the uptake of Fe(II) can be attributed to the release of iron from magnetite at low pH:
when the pH increases this release decreases. The stability of the composite material is an
important issue to take into account and it sounds preferable managing the hybrid material in
solution whose pH is higher than pH 1.5-2. The selectivity of the sorbents for Hg(II) against
other base metals may be explained by the sorption mechanism which is attributed to ion-
exchange of Hg(II) chloro-anions on protonated amine groups in acid conditions: the other
selected metal ions are not supposed to form stable chloro-anions under selected experimental
conditions.
3.5.7. Effect of sorbent dose
The sorption of Hg(II) on modified MCGMA sorbents was studied by changing the sorbent
dosage (in the range 1.5-10 g L-1
), metal concentration being fixed to 9.8 mM (pH 4 and T: 20
°C). Figure AM5a (See Additional Material Section) shows the variation in sorption capacity
with sorbent dosage: the sorption capacity decreases from 1.52 to 0.97 mmol Hg g-1
and from
2.07 to 0.97 mmol Hg g-1
for MCGMA-I and MCGMA-II, respectively when the SD
increases from 1.5 to 10 g L-1
. Figure AM5b (See Additional Material Section) shows the
effect of sorbent dosage on the relative equilibrium concentration (C/C0). As expected when
the sorbent dosage increases, the equilibrium concentration of Hg(II) decreased. The results
shown indicate that the residual concentration C/C0 may decrease up to 0.0014 for SD = 10 g
L-1
. At low sorbent dosage, all sorption sites are exposed and the can be easily saturated. With
an increase of sorbent dosage the amount of sorption sites becomes in excess compared to
present metal ions and the saturation of the sorption sites is not achieved. According to these
results, the sorbent dosage of 1.5 g L-1
allows achieving the saturation of the sorbent and
makes possible the concentration effect and the enrichment of mercury after desorption step.
27
On the opposite hand with a sorbent dosage of 10 g L-1
it is possible removing about 99.6 %
of mercury from the solution (achieving environmental target). The sorbent dosage to be used
depends on the target of the sorption process: concentration effect or maximum
decontamination.
3.5.8. Comparison of sorption capacity for Hg(II) ions with various sorbents
Table 9 shows the comparison of maximum sorption capacities of MCGMA-I and MCGMA-
II with a series of values found in the literature (together with the best operating conditions
reported by respective authors). A direct comparison of sorption performance is difficult due
to different experimental conditions; however, this is a useful criterion for roughly evaluating
the potential of these materials. The modified MCGMA sorbents have a sorption capacity of
the same order of magnitude as other sorbents; although chitosan foam [40] and thiol-
modified activated coke [41] showed better sorption capacity. It is noteworthy that the
modified MCGMA sorbents have an important advantage related to their fast kinetics. The
high sorption capacity of the modified MCGMA sorbents towards Hg(II) ions reveals that
sorbents could be promising for practical application in Hg(II) ions removal from
wastewater.
3.5.9. Regeneration
Three successive sorption/desorption cycle runs were performed using 100 mL of 0.1 M KI
solutions for 50 min. Samples were collected at different time intervals and the residual
concentration of Hg(II) was determined. Figure 8 shows the time course of released Hg(II)
ions. About 97% of the sorbed Hg(II) ions is released within 15 min. The desorption is even
faster than the sorption. Though a slight and progressive decrease in the amount of desorbed
Hg(II) ions is observed along the successive cycles the desorption efficiencies remained
28
higher than 90 % after three sorption/desorption cycles. The sorbent MCGMA-I seems to be
slightly more efficient than MCGMA-II sorbent in terms of desorption with desorption yield
of 97.2, 98.2 for MCGMA-I and 99.6%, 89.0% for MCGMA-II for the first three
sorption/desorption cycles, respectively.
4. Conclusion
Novel magnetic macromolecular hybrid materials were prepared by grafting glycidyl
methacrylate (synthetic polymer) to chitosan (biopolymer) and functionalizing the
intermediary compound with diethylenetriamine (directly, or via epichlorhydrin spacer,
MCGMA-I and MCGMA-II, respectively). The two sorbents are characterized by efficient
and selective sorption towards Hg(II) ions from aqueous medium at approximately pH 4. The
uptake kinetics are well fitted by the PSORE, while the distribution of the metal at
equilibrium between the solid and the liquid is modeled by the Langmuir equation. The
maximum capacity reached 1.43 mmol g−1
for MCGMA-I and 2.01 mmol g−1
for MCGMA-
II at 20 oC. The tests performed in the presence of other heavy metals (multicomponent
solutions) demonstrated the selectivity of the sorbents for Hg(II). Depending on the pH two
sorption mechanisms can be involved: ion-exchange on protonated amine groups in acidic
solutions, or chelation on amine groups at mild acidic pH. However, this interpretation of
sorption mechanisms is affected by the composition of the solution and more specifically the
nature of counter-anions present in the solution. The presence of counter anions that form
anionic complexes in acidic solutions makes possible Hg(II) sorption while those counter
anions that do not form stable complexes with mercury reduce the affinity of the sorbent for
metal ions. The investigation of temperature effect on Hg(II) sorption isotherm confirms that
the sorption mechanism is endothermic. Mercury ions bound to the sorbents can be
29
successfully desorbed using 0.1 M potassium iodide solutions and the sorbent can be
effectively re-used.
Acknowledgements
This study was supported by the French Government through a fellowship granted from the
French Embassy in Egypt (Institut Français d'Égypte). The authors would like to thank
Thierry Vincent, André Brun and Jean-Marie Taulemesse for their technical and scientific
contributions to this work.
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aqueous solution by chitosan-coated magnetic nanoparticles modified with alpha-ketoglutaric
acid, J. Colloid Interface Sci., 330 (2009) 29-37.
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derivatives: A review, J. Hazard. Mater., 167 (2009) 10-23.
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(2005) 267-275.
[26] X. Li, Z. Liu, J.-Y. Lee, Adsorption kinetic and equilibrium study for removal of
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253 (2013) 419-427.
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Amer. Chem. Soc., 40 (1918) 1361-1402.
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carbons. I. Adsorption of organic vapors, Zh. Fiz. Khim., 21 (1947) 1351-1362.
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Am. Chem. Soc., 60 (1938) 309-319.
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35
List of Tables
Table 1: Kinetics models and their linear forms
Table 2: Kinetic parameters for Hg(II) sorption
Table 3: Sorption isotherms and their linear forms
Table 4: Parameters of the models for sorption isotherms
Table 5: Effect of temperature on Hg(II) sorption capacity
Table 6: Standard enthalpy, entropy and free energy changes for Hg(II) sorption
Table 7: Effect of counter ion on Hg(II) sorption
Table 8: Multicomponent sorption – Effect of pH on the sorption capacity for Hg(II) Co(II),
Cu(II), Fe(II), Ni(II), Zn(II) and Mg(II) (initial concentration of 1.0 mmol metal L-1
for each
metal ion).
Table 9: Comparison of sorption capacity for Hg(II) ions with various sorbents
36
Table 1: Kinetics models and their linear forms
Kinetic model Non-Linear form Linear form Plot References
Pseudo-First
order ]1[ 1 tk
eqq et
t)
2.303
k(q log )q(q log 1
ete log (qe- qt) vs. t [27]
Pseudo-Second
order tqk
tkq
e
t
2
2
1 t)
q
1(
qk
1
q
t
e2
e2t
(t/qt) vs. t [28]
Intraparticle
diffusion
- qt = Kit
0.5 + X qt vs. t
0.5 [29]
Elovich
equation q
edt
dqt
lnt
1ln
1
tq
qt vs. lnt [30]
37
Table 2: Kinetic parameters for Hg(II) sorption
PFORE PSORE Weber & Morris model Elovich equation
Sorbent qe,exp k1 qe, calc R2 k2 qe, calc R
2 Ki X R
2 α β R
2
MCGMA-I 1.43 0.0130 0.91 0.942 0.035 1.49 0.9976
Ki1 0.124 0.247 0.953
0.470 4.55 0.993 Ki2 0.042 0.751 0.942
Ki3 0.013 1.178 0.889
MCGMA-II 2.01 0.0095 1.27 0.979 0.017 2.10 0.9899
Ki1 0.091 0.664 0.991
0.736 3.54 0.951 Ki2 0.079 0.689 0.989
Ki3 0.035 1.342 0.996
Units: qe, exp/calc: mmol g-1
., k1: min-1
), k2: g mmol-1
min-1
; Ki: mmol g-1
min0.5
; α: mmol g-1
min-1
; β: g mmol-1
.
38
Table 3: Sorption isotherms: equation in their linear and non-linear forms
Isotherm Non-Linear form Linear form Plot References
Langmuir eqL
eqLLmeq
CK
CKqq
1
,
Lm,Lm,
eq
eq
eq
q
1
q
C
q
C
LK
eqqeqC
vs. eqC [36]
Freundlich n
eqFeq CKq /1 eqCln
n
1Kln q ln feq ln qeq vs. ln Ceq
[37]
Dubinin–Radushkevich 2
expDR
DR
KQ
eqq
2
e lnqln DRDR KQ ln qeq vs. 2 [38]
39
Table 4: Parameters of the sorption isotherm models
Langmuir model Freundlich model Dubinin-Radushkevich (D-R) model
Sorbent qmax,exp
(mmol g−1
)
qm,L
(mmol g−1
)
KL
(L mmol−1
)
R2
n KF (mmol g
−1)
(L mmol−1
)1/n
R2 QDR
(mmol g−1
)
KDR
(J2 mol
−2)
Ea
(kJ mol−1
)
R2
MCGMA-I 1.43 1.55 27.78 0.999 12.87 1.37 0.965 1.50 4.72 × 10-9
10.3 0.892
MCGMA-II 2.02 2.28 16.48 0.999 7.98 1.89 0.945 2.16 6.48 × 10-9
8.78 0.894
40
Table 5: Effect of temperature on Hg(II) sorption capacity
Temperature oC
Sorption capacity (mmol Hg g-1
)
MCGMA-I MCGMA-II
20 1.48 2.12
30 1.74 2.40
40 1.91 2.58
50 2.21 2.65
Table 6: Standard enthalpy, entropy and free energy changes for Hg(II) sorption
Adsorbent
∆Ho
(kJ mol−1
)
∆So
(J mol−1
K−1
)
T0
(K)
∆Go (kJ mol
−1)
293 K 303 K 313 K 323 K
MCGMA-I 19.2 61.4 312.8 1.22 0.60 -0.01 -0.62
MCGMA-
II 15.3 54.0 283.8 -0.50 -1.04 -1.58 -2.12
41
Table 7: Effect of counter ion on Hg(II) sorption
Table 8: Multicomponent sorption – Effect of pH on the sorption capacity for Hg(II) Co(II),
Cu(II), Fe(II), Ni(II), Zn(II) and Mg(II) (initial concentration of 1.0 mmol metal L-1
for each
metal ion).
*: negative values for Fe(II) are associated to metal release due to magnetite partial
dissolving.
Initial
pH
Sorption capacity (mmol Hg g-1
)
HgCl2 HgNO3 HgSO4
MCGMA-I MCGMA-II MCGMA-I MCGMA-II MCGMA-I MCGMA-II
0.95 1.51 2.07 0.15 0.16 0.048 0.054
2.12 1.45 2.07 1.00 1.41 0.153 0.164
3.2 1.36 1.95 1.33 2.14 - -
Adsorbent Initial
pH
Uptake (mmol g-1
)
Final
pH Metal ion
Hg(II) Co(II) Cu(II) Fe(II)* Ni(II) Zn(II) Mg(II)
MCGMA-I
0.4 0.82 0.280 0.016 0.001 -0.209 0.002 0.034 0.000
0.8 1.12 0.342 0.006 0.000 -0.118 0.011 0.021 0.001
1.2 1.26 0.350 0.016 0.005 -0.008 0.015 0.008 0.004
1.6 1.80 0.301 0.014 0.004 -0.005 0.016 0.000 0.000
MCGMA-II
0.4 0.90 0.310 0.014 0.000 -0.192 0.004 0.027 0.000
0.8 1.22 0.372 0.000 0.003 -0.109 0.005 0.008 0.000
1.2 1.29 0.380 0.001 0.006 -0.007 0.000 0.000 0.005
1.6 1.86 0.331 0.000 0.013 -0.004 0.000 0.000 0.002
42
Table 9: Comparison of sorption capacity for Hg(II) ions with various sorbents
Adsorbent material
Initial pH Contact
time (min)
Temperature
(oC)
Initial
concentration
(mM)
Sorbent
dosage (g L-1
)
Sorption capacity
(mmol Hg g−1
) References
Thiourea modified Hg(II)
ion-imprinted cellulosic cotton fibers 5.00 180 30 1.99 0.30 0.54 [42]
Schiff-base modified guar gum 5.00 120 25 0.24 1.00 0.20 [43]
Hydrolyzed acrylamide-grafted PET
films 4.5 40 25 0.49 4.00 0.07 [44]
ZnCl2-MCM-41 6.00 30 20 0.24 0.33 0.43 [45]
Chitosan foam 4.00 2880 20 0.49 0.45 1.74 [40]
Thiol modified activated coke 5.00 480 25 0.49 0.10 2.29 [41]
Glycidylmethacrylate grafted on
cellulose 5.00 50 - 0.24 20.00 0.18 [6]
Sheep bone charcoal 3.0 360 25 0.39 4.00 0.06 [46]
MCGMA-I 4.0 300 20 5.00 2.50 1.43 This work
MCGMA-II 4.0 300 20 5.00 2.50 2.02 This work
43
OH
NH2
H2NO
O
O
O
O
n
HO
OH
O CH2 CH2CH
OH
NH2
H2N
HO
OH
polymerization
NH2
NHNH2
CHHO
CH2
NH2
NHNH2
NH2
NHNH2
CH
CH2
HO
(MCGMA) O CH2 CH2CH
OH
NH2
H2N
HO
OH
Cl CH2 HC CH2
OHN
CH2H2C CH
O
OHO
O
O
(MCGMA)
NH
CH2 HC CH2
OHO
CH2CH
OH
O CH2
HO
HN
CH2H2C CHNH2
NHNH
OH
OHO
NH2
NHNH2(MCGMA-ECH)
NH2
NHNH2
CHHO
NH
CH2H2C CH
OH
NH2
NHNH
CHHO
CH2
NH2
NHNH2
CH2CH
OH
O CH2
(MC) + (GMA) + (magnetite) (MCGMA)
(MCGMA-I)
(MCGMA-ECH)
(MCGMA-II)
Scheme 1: Scheme for the synthesis and modification of magnetic-macromolecular hybrid
material
44
OO
O
OHO NH2
CH2
HC OH
H2C
O
C O
CCH3H2C
H2C CH
CO
NH
CH2
NH
CO
CHH2C
OO
O
OHO NH2
CH2
HC OH
H2C
O
C O
C
CH3
CH2
CH
CO
NH
CH2
NH
CO
CH
CH2
Chitosan-GMA (MBA) Macromolecular hybrid material
(MCGMA)
Scheme 2: proposed chemical structure of chitosan–glycidyl methacrylate - N,N’methylene-
bis-acrylamide (macromolecular hybrid material)
45
Figures captions
Figure 1: Scanning electron micrographs of; (a) unloaded MCGMA-I, (b) unloaded
MCGMA-II, (c) Hg(II)-loaded MCGMA-I and (d) Hg(II)-loaded MCGMA-II.
Figure 2: Energy dispersive X-ray analysis (EDX) of MCGMA-I and MCGMA-II sorbents:
(a) unloaded MCGMA-I, (b) unloaded MCGMA-II, (c) Hg(II)-loaded MCGMA-I and (d)
Hg(II)-loaded MCGMA-II.
Figure 3: Magnetization curves for MCGMA-I and MCGMA-II sorbents.
Figure 4: FT-IR spectra of MCGMA-I and MCGMA-II sorbents.
Figure 5: pH effect on Hg(II) sorption using MCGMA-I and MCGMA-II: (a) sorption
capacity, (b) pH change (T: 20 oC; C0: 10 mmol Hg L
-1).
Figure 6: Hg(II) uptake kinetics using MCGMA-I and MCGMA-II (a: relative concentration
decay of Hg(II) in the solution, b: Hg(II) concentration in the sorbent (i.e., q(t)) vs. time) (C0:
10 mmol Hg L-1
; pH 4; T: 20 oC; sorbent dosage, SD: 2.5 g L
-1).
Figure 7: Hg(II) sorption isotherms using MCGMA-I and MCGMA-II (pH 4; T: 20 °C).
Figure 8: Desorption kinetics for successive sorption/desorption cycles for MCGMA-I (a)
and MCGMA-II (b).
46
(a) (b)
(c) (d)
FIG. 1
47
-2500 -2000 -1500 -1000 -500 0 500 1000 1500 2000 2500
Field (G)
-25
-20
-15
-10
-5
0
5
10
15
20
25
Mo
me
nt/M
ass (
em
u g
-1)
MCGMA-I
MCGMA-II
(a) (b)
(c) (d)
FIG. 2
FIG. 3
48
FIG. 4
40080012001600200024002800320036004000
Tra
nsm
itta
nce
Wavenumber (cm-1)
MCGMA-I
MCGMA-II
3406.6
3414.3
2940.0
2940.9
2850.3
2857.0
1724.0
1652.7
1657.5
1561.1
1572.7
1469.5
1465.6
1392.3
1390.4
1264.1
1162.9
986.4
750.2
586.3
990.3
752.1
3406.6
3414.3
2940.0
2940.9
2850.3
2857.0
1724.0
1652.7
1657.5
1561.1
1572.7
1469.5
1465.6
1392.3
1390.4
1264.1
1162.9
986.4
750.2
586.3
990.3
752.1
595.91265.1
1161.9
49
0 1 2 3 4 5 6 7
Initial pH
0
1
2
3
4
5
6
7
Fin
al p
H
MCGMA-I
MCGMA-II
0 1 2 3 4 5 6 7 8 9 10
Equilibrium pH
0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
q e
(m
mol g
-1)
MCGMA-I
MCGMA-II
(a)
(b)
FIG. 5
50
0 50 100 150 200 250 300 350 400 450 500
t (min)
0.0
0.3
0.6
0.9
1.2
1.5
1.8
2.1
q (
t)
(mm
ol g-1
)
MCGMA-I
MCGMA-II
Simulated from PSORE (MCGMA-I)
Simulated from PSORE (MCGMA-II)
0 50 100 150 200 250 300 350 400
t (min)
0.4
0.5
0.6
0.7
0.8
0.9
C(t
)/C
0
MCGMA-I
MCGMA-II
(a)
(b)
FIG. 6
51
FIG. 7
0 1 2 3 4 5 6 7 8 9 10
Ce (mM)
0.0
0.3
0.6
0.9
1.2
1.5
1.8
2.1
2.4q
e (m
mol g
-1)
MCGMA-I
MCGMA-II
Simulated Langmuir isotherm (MCGMA-I)
Simulated Langmuir isotherm (MCGMA-II)
52
0 10 20 30 40 50 60
t (min)
2.6
2.8
3.0
3.2
3.4
3.6
3.8
4.0
Re
lea
se
d H
g(I
I)
(mM
)
Cycle I
Cycle II
Cycle III
0 10 20 30 40 50 60
t (min)
4.0
4.2
4.4
4.6
4.8
5.0
5.2
5.4
5.6
5.8
6.0
Rele
ase
d H
g(I
I) (
mM
)
Cycle I
Cycle II
Cycle III
(a)
(b)
FIG. 8
53
Additional Material Section
Figure AM1: SEM images of MCGMA-I and MCGMA-II sorbents at different
magnitudes).
Figure AM2: Modeling of uptake kinetics with: (a) PFORE, (b) PSORE.
Figure AM3: Modeling of uptake kinetics with (a) simplified model of resistance to
intraparticle diffusion (Morris and Weber equation), (b) Elovich equation.
Figure AM4: Linearized plots for Hg(II) sorption isotherms: (a) Langmuir equation, (b)
Freundlich equation, (c) Dubinin–Radushkevich equation.
Figure AM5: Effect of sorbent mass (SM) on Hg(II) sorption using MCGMA-I and
MCGMA-II: (a) sorption capacity vs. SM, (b) relative residual concentration (C/C0) vs. SM
(V: 20 mL; C0: 10 mmol Hg L-1
; T: 20 °C; pH 4).
54
FIG. AM1
MCGMA-I MCGMA-I MCGMA-I
500 µm 100 µm 20 µm
20 µm100 µm500 µm
MCGMA-II MCGMA-II MCGMA-II
55
0 40 80 120 160 200 240 280 320 360 400
t (min)
-3.0
-2.5
-2.0
-1.5
-1.0
-0.5
0.0
0.5
1.0
1.5
log (
q e
- q
t )
MCGMA-I
MCGMA-II
0 50 100 150 200 250 300 350 400 450 500
t (min)
0
50
100
150
200
250
300
t/q
t (m
in g
mm
ol-
1)
MCGMA-I
MCGMA-II
(a)
(b)
FIG. AM2
56
0 5 10 15 20 25
t 0.5 (min) 0.5
0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0q
e (m
mol g
-1)
MCGMA-I
MCGMA-II
1 2 3 4 5 6 7
ln t
0.0
0.5
1.0
1.5
2.0
2.5
3.0
3.5
4.0
q e
(m
mol g
-1)
MCGMA-I
MCGMA-II
(a)
(b)
FIG. AM3
57
0 1 2 3 4 5 6 7 8
Ce (mM)
0
1
2
3
4
5
6
Ce/q
e (g
L-1
)
MCGMA-I
MCGMA-II
-4.5 -3.0 -1.5 0.0 1.5 3.0
ln Ce
-0.2
0.0
0.2
0.4
0.6
0.8
1.0
ln q
e
MCGMA-I
MCGMA-II
(a)
(b)
(c)
FIG. AM4
0.0E+0 3.0E+7 6.0E+7 9.0E+7 1.2E+8 1.5E+8-0.2
0.0
0.2
0.4
0.6
0.8
1.0
ln q
e
MCGMA-I
MCGMA-II
ε2 (J2 mol−2)
58
0.00 0.04 0.08 0.12 0.16 0.20 0.24
Adsorbent mass (g)
0.0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1.0
C/C
0
MCGMA-I
MCGMA-II
0.00 0.04 0.08 0.12 0.16 0.20 0.24
Adsorbent mass (g)
0.5
1.0
1.5
2.0
2.5
3.0
q e
(m
mol g
-1)
MCGMA-I
MCGMA-II
(a)
(b)
FIG. AM5
Highlights
• Magnetic hybrid material for mercury recovery from slightly acidic solutions
• Diethylenetriamine grafting on chitosan for improved sorption performance
• Pseudo-second order rate equation for modeling uptake kinetics
• Selectivity for mercury recovery against other base metals
• Mercury can be easily desorbed for sorbent for efficient recycling