Adsorption of Mn 2+ from Aqueous Solution Using Manganese Oxide-Coated Hollow Polymethylmethacrylate Microspheres (MHPM)

Abstract

Results of investigation on adsorption of Mn²⁺ from aqueous solution by manganese oxide-coated hollow polymethylmethacrylate microspheres (MHPM) are reported here. This is the first report on Mn-coated hollow polymer as a substitute for widely used materials like green sand or MN-coated sand. Hollow polymethylmethacrylate (HPM) was prepared by using a literature procedure. Manganese oxide (MnO) was coated on the surface of HPM (MHPM) by using the electroless plating technique. The HPM and MHPM were characterized by using optical microscopy (OM), scanning electron microscopy (SEM), Fourier-transform infrared spectroscopy (FTIR), X-ray diffraction (XRD), and thermogravimetric analysis (TGA). Optical and scanning micrographs were used to monitor the surface properties of the coated layer which revealed the presence of MnO on the surface of HPM. TGA showed the presence of 4-5% of MnO in MHPM. Adsorption isotherm studies were carried out as a function of pH, initial ion concentration, and contact time, to determine the adsorption efficiency for removal of Mn²⁺ from contaminated water by the synthesized MHPM. The isotherm results showed that the maximum adsorption capacity of MnO-coated HPM to remove manganese contaminants from water is 8.373 mg/g. The obtained $R^{2}$ values of Langmuir isotherm and Freundlich isotherm models were 1 and 0.87, respectively. Therefore, $R^{2}$ magnitude confirmed that the Langmuir model is best suited for Mn²⁺ adsorption by a monolayer of MHPM adsorbent. The material developed shows higher adsorption capacity even at a higher concentration of solute ions, which is not usually observed with similar materials of this kind. Overall findings indicate that MHPM is a very potential lightweight adsorbent for removal of Mn²⁺ from the aqueous solution because of its low density and high surface area.

1. Introduction

Heavy metal contamination of water bodies from improperly disposed industrial wastes is a serious environmental problem across the world. Manganese is one among those heavy metal contaminants and is used principally in the manufacture of iron and steel alloys [1]. Manganese dioxide and other manganese compounds are used in products such as batteries, glass, and fireworks. Potassium permanganate is used as an oxidant for cleaning, bleaching, and disinfection purposes (HSDB, 2001) [2]. Manganese green sands are used in water treatment plants. An organic manganese compound, methylcyclopentadienyl manganese tricarbonyl (MMT), is used as an octane-enhancing agent in unleaded petrol in Canada, the United States of America (USA), Europe, Asia, and South America [3]. Some manganese compounds are used in fertilizers, varnish, and fungicides and as livestock feeding supplements. All these activities lead to anthropogenic causes for Mn contamination of water sources. Manganese is found in abundance on earth’s crust and usually present in the form of compounds, rather than the elemental form. In the compounds, “Mn” is present in different oxidation states such as Mn²⁺, Mn⁴⁺, or Mn⁷⁺ (USEPA, 1994) [4]. It is generally present in groundwater as Mn²⁺ and is considered a pollutant beyond a certain limit, because of its organoleptic properties. According to the World Health Organization, the maximum acceptable “Mn” concentration in drinking water is 0.05 mg/l [WHO, 2011] [5, 6]. According to Indian Standard (IS), 10500 :2012, the acceptable limit for Mn in drinking water is 0.1 mg/l.

In surface waters, manganese occurs in both dissolved and suspended forms, while in groundwater, it remains in a dissolved state [7]. Mn is mainly formed by hydrogeochemical reactions and is found in the form of hydroxides, sulfates, or carbonates of the divalent Mn²⁺ ion. Thus, anaerobic groundwater often contains elevated levels of dissolved manganese. The divalent form (Mn²⁺) predominates in most water at pH 4–7, but at higher pH values, more highly oxidized forms may occur [8]. Upon exposure to air, Mn²⁺ forms MnO₂ particulates that can cause black or brown stains on household utensils, plumbing fixtures, and clothing. Mn otherwise is an essential nutrient, and research findings indicated that [9] a minimum quantity of 3.5–7 mg/day for adults is necessary for regular physiological functions [9]. But, a higher concentration of Mn²⁺ in drinking water is hazardous for health which may cause damage to the central nervous system and neurological systems in humans [10]. This may even result in Parkinson’s disease [11, 12]. Excess manganese ion concentration in drinking water imparts an unpleasant metallic taste, stains plumbing fixtures, and also leads to discoloration of water and laundry [13]. Similarly, when manganese(II) in solution undergoes oxidation, it gets precipitated, resulting in encrustation problems and may damage the water storage tanks and pipelines by forming coating [14, 15].

Thus, several physical, chemical, and biological Mn remediation strategies have been developed over the last few decades to mitigate the toxic effects of excess Mn and to ensure a safe supply of drinking water.. Currently, many treatment methods, such as oxidation [16–19], membrane filtration [20], and adsorption [21–24], are employed in removing high manganese concentrations from groundwater [25]. Although the oxidation method can remove manganese rapidly, the generation of some toxins demonstrates the disadvantages of this method.

Membrane filtration [26–29] is an effective method for heavy metal removal, but it has the limitations like a requirement of specific pH, which accelerates membrane fouling [15, 16]. An adsorption technique however has appeared to be more economical and has received extensive attention for the removal of manganese from groundwater [30, 31]. The application of the adsorption method is highly dependent on the adsorptive materials [32]. Some natural materials like zeolites, chitosan, and clay are known as low-cost materials that are used for the purpose [33]. Especially, zeolites, the low-cost materials with exceptional ion exchange and sorption properties, can abundantly be used [34, 35]. However, pH has a significant impact on the ion exchange properties of these highly porous zeolite-based materials, which may even sometimes negate its cost economic nature.

Therefore, the development of new highly porous and highly efficient cost economic materials for removal of cations from drinking water sources is important to research objectives and many scientific interests are focused on that. Manganese oxide has been successfully used as a catalytic adsorbent for the removal of manganese [36]. But this too has certain strict limitations that restrict its use as an adsorbent for manganese removal. Immobilization of manganese oxide on materials like sand [25], clay [37], zeolites [38–40], active carbon [41, 42], and glass beads [43] has been carried out, and these MnO-coated materials are used as a catalyst for effective removal of Mn from water. The MnO-coated sand showed the best results in the pH range of 5.5 to 8.0. Mn²⁺ uptake also increased with increasing temperature from 25 to 45°C. The study also showed that the mechanism was chemisorption. The equilibrium adsorption data were best represented by the Langmuir isotherm [25]. Similarly, MnO-coated zeolite (MOCZ) materials are also known which have a good binding affinity and hence strong adsorption capacity for Mn²⁺ ions [40]. Other known MnO-coated zeolite showed maximum removal of Mn at pH 7, with an Mn concentration of 5 mg/l. The equilibrium adsorption data were best represented by the Langmuir isotherm, indicating monolayer adsorption on a homogenous surface, and the adsorption capacity was found to be 0.21 mg/g at 20°C. It was assumed that the adsorption process was controlled by physical mechanism rather than chemical [39]. Zohreh and Abedi investigated the efficiency of impregnated active carbon in manganese removal from aqueous solutions which showed highest adsorption of Mn²⁺ ions at the pH of 9 (Mn²⁺ removal: 67.19%) with the adsorption kinetics fitting the pseudo-second-order and intraparticle diffusion models. They confirmed the suitability of the Langmuir model for Mn²⁺ adsorption using impregnated active carbon [42]. The manganese is also being removed by using MnO-coated glass beads, which are developed using dry and wet coating techniques [43]. In most of the cases, pH has a decisive role in the Mn²⁺ removal process.

In the present study, we have attempted to utilize manganese oxide-coated hollow polymethylmethacrylate microspheres, to develop a low-cost and highly efficient novel adsorbent (MHPM) for the removal of Mn contaminant from water. The adsorption capacity of MHPM was evaluated by adsorption isotherm studies as a function of parameters such as pH, initial ion concentration, and contact time. Langmuir and Freundlich isotherm parameters were determined and compared. Therefore, the main objectives of this work were to investigate the ability of manganese removal onto MHPM and to reveal a probable removal mechanism of manganese from the contaminated water.

2. Materials and Methods

2.1. Materials

Polymethylmethacrylate (PMMA) [Sigma-Aldrich, MW (avr.): 120,000, 98%, viscosity 0.20 g/dl (lit.)], dichloromethane [Merck, mfg date: Jan 2013, 99.5%, $M = 84.93 g / mol$ ], polyvinyl alcohol [Central Drug House, Delhi, MW (avr.): 125,000, 99.25% viscosity 35-50 cP at 4% cold aqueous solution], acetic acid [Merck, $M = 60.052 g / mol$ ], manganese(II) sulfate (MnSO₄) [Aldrich, $M = 151.001 g / mol$ ], and potassium permanganate (KMnO₄) solution [Aldrich, $M = 158.034 g / mol$ ] were used. All other reagents used are of analytical grade and were obtained from Aldrich, India. The limestone was acquired from the mining site, Meghalaya, India.

2.2. Preparation of HPM Microspheres

The solvent evaporation technique [44] is used to synthesize hollow PMMA microspheres. In the first step, a solution was prepared by dissolving PMMA (5-6% $w / v$ ) in dichloromethane with the help of a magnetic stirrer. The solution was then added dropwise to a stirring aqueous medium. The aqueous medium comprises (0.5%, $w / v$ ) of poly(vinyl alcohol) which acts as a stabilizer. The stirring is maintained at 550 rpm with a propeller-type mechanical stirrer. PMMA microspheres (HPM) are formed by slow evaporation of dichloromethane at room temperature. At the end of the reaction, hollow PMMA microspheres are obtained after washing with water and drying at 70°C. The bulk density of the HPM is calculated as 0.69 gm/cc.

2.3. Preparation of MHPM

The manganese oxide-coated PMMA was synthesized, in a 1 l beaker. In the primary step, a solution is prepared by adding 10 g of PMMA to 250 ml of 2 N acetic acid containing 0.5 M MnSO₄. The solution was stirred as 200 ml of 0.43 M potassium permanganate solution was added. The coated material was settled before air drying at 30°C, was rinsed with Milli-Q water, and was redried. The obtained yield is 72%. Then, the product was stored in an amber glass bottle.

2.4. Characterization of Materials

The synthesized materials were characterized by a scanning electron microscope with energy-dispersive X-ray analysis attachment (Carl Zeiss, EVO50), a FTIR (Bruker model Alpha-T), an optical microscope with Leica DMLM/P (Leica Microsystems AG Switzerland at 50x magnification), an transmission electron microscope (TEM) (JEOL 200 kV model no. JEE2100), and thermogravimetric analysis (TGA) (TA Instruments, USA; model 2950 and 2910). FTIR of the samples was taken with the help FTIR spectrophotometer (Bruker Alpha model with KBr). X-ray diffraction patterns were measured with XRD (model: Rigaku Ultima IV IR Technology Services Pvt. Ltd., in powder mode with Cu and Mo with Kβ filter). The UV-visible spectrophotometer is a double beam spectrophotometer from Analytik Jena Model SPECORD 205 (via standard IS 3025 (Part 59): 2006).

3. Results and Discussion

3.1. Characterization of HPM and MHPM

3.1.1. Optical Microscope Analysis

The optical micrographs in Figure 1 show the micrographs for uncoated PMMA (HPM) and MnO-coated microspheres (MHPM). The micrograph of uncoated hollow polymethylmethacrylate microspheres is of white-colored spheres with different sizes. On the contrary, with clear distinction, the MnO-coated microspheres are appeared in pinkish-brown/dark brown color due to the deposition of MnO on the surface of HPM microspheres. Thus, the color difference between the coated (MHPM) and uncoated (HPM) microspheres confirms the presence of MnO coating on the surface of microspheres.

Figure 1

Optical micrographs of (a) uncoated and (b) MnO-coated hollow PMMA MS.

(b)

3.1.2. SEM Analysis

The SEM images in Figure 2 are captured by a scanning electron microscope with X-ray diffraction analysis attachment. These SEM micrographs are used to analyze the surface morphology of MHPM. The SEM micrographs illustrate that the diameter of the HPM microspheres is in the range of 20–80 μm. The micrographs of uncoated HPM microspheres show the smooth surface of microspheres, whereas the surface of MnO-coated HPM microspheres shows the occurrence of a precipitate of MnO, i.e., manganese oxide formed as clusters on to the polymer surface, leading to surface roughness of microsphere. Hence, SEM micrographs indicate the presence of MnO coating on HPM surface.

Figure 2

SEM micrographs of (a) uncoated and (b) MnO-coated hollow PMMA MS.

(b)

3.1.3. TEM Analysis

Figure 3 shows the TEM images of HPM and MHPM. The image textures for the HPM and MHPM are completely different. The image for HPM shows a smooth kind of feature, whereas the image for MHPM shows well-dispersed MnO particles throughout the polymeric matrix. Hence, in general, a good dispersion of MnO particles was obtained except for particle aggregation at some places which is difficult to avoid for small particles. Aggregation of the MnO particles also influences the aspect ratio (average diameter may change) which in turn may influence the surface texture of the polymeric microspheres. There are possibilities of overlapping of MnO particles along the depth direction of the specimen (i.e., $Z$ -direction) resulting in observation of poor resolution along this direction for the TEM images. The study confirms a random coating of the MnO on the surface of PMMA. The uneven surface confirms physical interaction rather than chemicals between the particles and polymeric surfaces. This has further been confirmed by the isotherm studies.

Figure 3

TEM micrographs of (a) uncoated and (b) MnO-coated hollow PMMA MS.

(b)

3.1.4. XRD Analysis

The XRD diffraction patterns of HPM and MHPM are displayed in Figure 4. A broad and shallow peak at around $2 θ$ of 16° is a typical feature of amorphous PMMA. The diffraction curve of MHPM shows peaks at around $2 θ$ of 12°, 24°, 37°, and 66°. Besides, the weak diffraction intensity of the samples reveals a poor crystallization or amorphous form of the MnO₂-based materials prepared by a facile chemical technique.

Figure 4

X-ray diffraction (XRD) patterns of HPM (red) and MHPM (blue).

The patterns of native and impregnated active carbon implied the formation of manganese oxide. The Mn-coated PMMA is confirmed by the peaks at $2 θ$ of 23° and 36° (Figure 4). According to the ICDD 00-012-0720 data, the peaks appeared at $2 θ$ of 23° and 36°, indicating the presence of manganese oxide [45].

3.1.5. TGA Analysis

TGA curves of HPM and MHPM are displayed in Figure 5. The curve shows that there is a slight weight loss during initial heating up to 85°C for both the materials which can be assigned to traces of moisture present in the sample. The decomposition temperature of PMMA is observed to be in the range of 230°C to 420°C, and complete decomposition was observed at 420°C. A similar trend of weight loss was displayed by MnO-coated HPM except for the residual weight of 8.9%. This residual weight corresponds to MnO coating. Thermal analysis of the materials thus confirms good hydrothermal stability of the materials.

Figure 5

TGA analysis of HPM (blue) and MHPM (red).

3.1.6. FTIR Analysis

The presence of MnO on the HPM microsphere can also be analyzed using FTIR spectroscopy. The FTIR spectra of HPM and MHPM are shown in Figure 6. The sharp bands observed at wavelengths lower than 750 cm^-1 are attributed to Mn–O stretching of MnO [16]. The two bands that appeared at 515 and 480 cm^-1 of the FTIR spectra can be associated with the Mn-O bond and thus confirming the deposition of Mn over the polymeric surface of PMMA. After adsorption, MHPM exhibited an extra band at 515 cm^-1, which corresponds to the Mn-O bond.

Figure 6

FTIR analysis of HPM and MHPM.

3.2. Optimization of Adsorption Parameters

3.2.1. Effect of Contact Time

A time-dependent study of Mn removal was carried out as shown in Figure 7. The removal shows a gradual increase with increasing time, and complete effectiveness was obtained within 70 min. After 70 min, it gets almost saturated. The experiment is carried out for the maximum initial concentration of Mn of 150 mg/l as compared to other adsorption tests. This is in accordance with previously reported literature [46].

Figure 7

Effect of contact time of Mn(II) on % removal.

3.2.2. Effect of Mn Concentration on % Removal and Adsorption Capacity

An experiment was carried out on MnO-coated hollow polymethylmethacrylate microspheres (MHPM) to investigate the effects of Mn concentration on the percentage removal of Mn²⁺ from the aqueous solution. The initial concentration of Mn was chosen as 20 mg/l, 30 mg/l, 50 mg/l, 100 mg/l, and 150 mg/l, respectively, at five different contact times. The percentage removal of Mn²⁺ from aqueous solution at equilibrium time was noted as 98.7%, 86.04%, 82.05%, 77.31%, and 55.14%, respectively, for the above solutions. This shows with an increase in initial concentration, the adsorption capacity of MHPM decreases.

This fact is represented in Figure 8(a), which shows that with an increase in initial concentration of Mn(II) ions in solution, percentage removal of Mn²⁺ from aqueous solution decreases. The decrease in removal efficiency is due to the unavailability of active sites on MHPM for the adsorption of the metal ions onto it. The variations of these concentrations are small at the low concentrations and increase more significantly with the increase in concentration. This is occurring because the ratio of available metal ions to the available active sites is small at low initial concentrations.

Figure 8

Effect of the initial concentration of Mn(II) on (a) % removal and (b) adsorption capacity.

(b)

The adsorption capacity for 20 mg/l initial concentration of Mn(II) in aqueous solution was evaluated as 1.974 mg/g from Figure 8(b). The same was found to increase to 2.7012 mg/g for 30 mg/l of Mn(II) ion concentration in aqueous solution and was continuously increasing until an equilibrium adsorption capacity for Mn(II) ion was reached at around 8.3 mg/g. From this data, it is inferred that adsorption capacity varies with the initial concentration of the Mn(II) in an aqueous solution. This is due to the available binding site on the surface of the PMMA microsphere in the initial stage where the binding site is in a higher amount than the Mn(II) in solution. But once the binding sites are saturated at a higher concentration of the Mn(II) in solution, the absorption capacity also reached its saturation. This is an indication of the monolayer adsorption mechanism.

3.2.3. Effect of pH on Adsorption Capacity

As it can be observed in Figure 9, Mn²⁺ adsorption by the MHPM is less significant in the lower pH range. This happens possibly due to the fact that at lower pH, there is a competition between the H⁺ ion and Mn²⁺ ions for the available exchangeable sites on MHPM and due to the preponderance of H⁺ ions, it wins the race. With an increase in pH, there is a concomitant increase in negative charge density on the MHPM surface causing deprotonation of the metal-binding sites, and eventually, Mn²⁺ ion adsorption increases. Beyond pH 8.0, we can observe a dramatic increase in the uptake capacity of Mn²⁺ ions by MHPM and this is continued up to pH 9.0. Beyond this value, Mn²⁺ ions predominantly remain in the precipitate form, and therefore, uptake of Mn²⁺ions by MHPM is impossible.

Figure 9

Effect of pH on % removal of Mn(II).

The dynamic increment in Mn expulsion as the suspension turns out to be progressively basic can be clarified by an expansion in the number of adsorption destinations. In the fluid arrangement, MnO present a hydroxylated surface whose charge is firmly pH reliant. As pH increases, surface negative charge increases due to the deprotonation of functional surface reunions (OH^-) and electrostatic forces bring cations into the surface. Actually, at low pH, prevailing H⁺ particle arrangements possess the restricting destinations of adsorbent, bringing about a net positive surface charge that forestalls cation adsorption.

3.3. Adsorption Isotherms

An adsorption isotherm represents the plot between the amount of solute adsorbed onto the solid and the equilibrium concentration of the solute in solution at a given temperature [47]. The adsorption isotherms indicate the distribution of adsorbate and adsorbent at equilibrium conditions for the adsorption process. These isotherms are represented in the form of graphical plots which set up a relationship between the equilibrium concentration of a solute on the surface of an adsorbent ( $q_{e}$ ) to the concentration of the solute in the liquid ( $C_{e}$ ), with which it is in contact. The nature of the adsorption process can be obtained from these plots. Langmuir and Freundlich adsorption isotherm models are used here to describe the metal ion distribution between the solid and liquid phases. The efficiency of synthesized MHPM to remove Mn(II) from water is determined by using these isotherm adsorption models.

3.3.1. Langmuir Isotherm

The Langmuir isotherm is suggested when all the adsorption active sites are occupied and further adsorption process is terminated. Therefore, this isotherm is used to establish equilibrium between adsorbate and adsorbent in the given system, where the partial pressure of the adsorbate approaches saturation [48]. At this saturation point, all the active sites are equally energetic and there is no interaction among the adsorbed molecules and the neighboring sites [49].

From Figure 6(b), it has been observed that for MHPM material, the equilibrium point was found at 150 mg/l and corresponds to 8.27 mg/g. From this result, it is assumed that the Langmuir model is best fitted at a higher concentration of adsorbate. At equilibrium conditions, the linear form of the Langmuir equation is given as follows: $\begin{matrix} (1) & \frac{q_{e}}{C_{e}} = K_{L} q_{m} - K_{L} q_{e}, \end{matrix}$ where $q_{e}$ (mg/g) is the amount of Mn(II) adsorbed, $C_{e}$ (mg/l) is the concentration of the Mn(II) solution, $K_{L}$ is the Langmuir constant related to adsorption enthalpy, and $q_{m}$ is the maximum adsorption capacity.

From Figure 10, it is shown that the Langmuir model is an empirical model having a linear plot, which indicates that the active sites on the surface are inadequate and the monolayer of MHPM is fitted into the heterogeneous surface. The capability of removing a certain amount of manganese ions by using synthesized MnO-coated HPM is described as the adsorption capacity ( $q_{e}$ ) of the particular adsorbent. The adsorption capacity is mathematically represented as follows: $\begin{matrix} (2) & q_{e} = X / M \end{matrix}$ where $X$ is the amount of adsorbate (Mn in water) and $M$ is the weight of adsorbent (manganese-coated PMMA).

Figure 10

Langmuir adsorption isotherm.

The capacity to retain the maximum amount of manganese (adsorbate) per unit mass of MHPM (adsorbent) is described as maximum adsorption capacity ( $q_{m}$ ). Usually, the $q_{m}$ is achieved at low concentrations and rarely achieved at high concentrations. The enthalpy of the adsorption is correlated with the Langmuir isotherm constant ( $K_{L}$ in g/l) which describes the affinity of MHPM towards Mn(II) ions. From Equation (1), we obtain intercept as $q_{m}$ and slope as $K_{L}$ . The favorable conditions of adsorption are indicated by correlation coefficient $R^{2}$ , whose value should be equal to 1. From the experimental results on MHPM, the obtained maximum adsorption capacity ( $q_{m}$ ) is 8.374 mg/g, the value of Langmuir constant ( $K_{L}$ ) is 1.201 mg/l, and $R^{2}$ value is obtained as 0.976. The Langmuir isotherm parameters obtained from adsorption data are listed in Table 1.

Table 1

Langmuir isotherm parameters.

Langmuir isotherm	$q_{m}$ (mg/g)	$K_{L}$ (mg/l)	$R^{2}$
Langmuir isotherm	8.374	1.201	0.976

The characteristics of Langmuir isotherm are evaluated by a value of a separation factor $R_{L}$ . This separation factor $R_{L}$ is inversely proportional to adsorption capacity ( $q_{e}$ ). As the adsorption capacity increases, the corresponding separation factor decreases thereby confirms that the isotherm model is well fitted and adsorption is favorable. $\begin{matrix} (3) & q_{e} = 1 / [1 + (R_{L} + C_{o}) \end{matrix}$

In Figure 11, a graph is plotted between separation factor $R_{L}$ and adsorption capacity ( $q_{e}$ ) which shows that separation factor depends upon the initial concentration of adsorbents in the solution. The figure represents that the separation factor decreases with the increase of initial concentration from 0.039 at 20 mg/l to 0.026 at 30 mg/l. In this case, the adsorption is said to be favorable as the $R_{L}$ value is less than 1 [47, 50, 51]. Therefore, this model is well fitted to the obtained adsorption data.

Figure 11

Langmuir isotherm for adsorption of Mn(II) ion on MnO-coated HPM surface.

3.3.2. Freundlich Isotherm

The Freundlich isotherm model suggests the multilayer adsorption and describes the adsorption on a heterogeneous surface [35, 52]. During the adsorption process, all the active sites on the adsorbent are filled with adsorbate, and therefore, the energy of adsorption declines exponentially. The obtained plot of the Freundlich isotherm is shown in Figure 9. Mathematically, the isotherm is represented as follows: $\begin{matrix} (4) & \log q_{e} = \log K_{F} + (1 / n) \log C_{e} \end{matrix}$ where the adsorption capacity and the adsorption intensity are given as “ $K_{F}$ ” and “ $n$ ”, respectively, which are also known as Freundlich constants.

$1 / n$ is given as heterogeneity factor and “ $n$ ” is a measure of the deviation from the linearity of adsorption.

The slope and the intercept obtained from the graph (Figure 12) plotted between log $q_{e}$ and log $C_{e}$ give adsorption intensity ( $1 / n$ ) and adsorption capacity ( $K_{F}$ ), respectively. The values of $1 / n$ also indicate the favorability and capacity of the adsorbate/adsorbent system.

Figure 12

Freundlich isotherm as a plot of log $q_{e}$ versus log $C_{e}$ .

From the experiments conducted on MHPM, the obtained value of “ $n$ ” is 3.897. Generally, the value of $n$ should be greater than 1 in case of higher surface density as the surface sites of the adsorbent are distributed inadequately to the adsorbate (manganese ions). As a result, the adsorption process is favorable when $n > 1$ . The affinity of adsorbent (MHPM) towards the adsorbate (manganese) is denoted by the value of $K_{F}$ . At higher values of $K_{F}$ , there is a strong binding between adsorbate molecules and the surface adsorbent. In this case, the $R^{2}$ value obtained for MHPM was 0.856 which is low compared to the Freundlich correlation factor, where $R^{2}$ is 1.

As shown in Table 2, Freundlich isotherm is an empirical model that fits the logarithm data well in the equation, but the obtained $R^{2}$ value as 0.856 is very low and thus proved that adsorption data is best fitted in the case of Langmuir isotherm. This is due to the separation factor of $R_{L}$ which establishes the high efficiency of Langmuir isotherm. Previous studies reported in the literature [34, 35, 40] on other adsorbents (e.g., zeolites) exhibited a good fit of the experimental data to the Langmuir isotherm, suggesting the existence of a plateau region where no further increase in Mn adsorption exists.

Table 2

Freundlich isotherm parameters.

Freundlich isotherm	$1 / n$	$n$	$K_{F}$ (mg/l)	$R^{2}$
Freundlich isotherm	0.257	3.897	2.801	0.856

4. Conclusion

We have successfully immobilized manganese oxide on HPM, to exploit its adsorption efficiency for the removal of Mn²⁺ ions from an aqueous solution. Although manganese oxide directly can be used as an adsorbent for removal of Mn²⁺ from an aqueous solution, it causes lots of inconveniences because it exists in the powder form in water. We have succeeded in overcoming this limitation. The material developed will have significant industrial utility since MHPM apart from being very economical is a lightweight material and offers high surface area compared to other materials used for the purpose. The manganese removal is influenced by various operational parameters such as pH, contact time, and initial concentrations. The most notable feature of the material is that for manganese concentrations from 20 to 50 mg/l, the removal efficiency was >82%. Further at higher pH, the material exhibits remarkable separation efficiency for Mn²⁺ ions from an aqueous solution. Even at a higher concentration of Mn²⁺ ions, the material shows maximum adsorption capacity, a quality rarely observed with other materials of this kind. Adsorption isotherm studies reveal that the adsorption efficiency of Mn(II) metal ions on MHPM is described by Langmuir and Freundlich isotherm models, and Langmuir isotherm provides the best fit. From the above results, we can conclude that this novel adsorbent MnO-coated HPM removes Mn ions from water resources efficiently.

Footnotes

Data Availability

All data is available in the manuscript.

Conflicts of Interest

The authors declare no competing financial or personal interests.

Acknowledgments

The authors would like to acknowledge the kind help of the director of DRL and the director and testing team of DMSRDE, Kanpur, along with the director and faculty members of NIT Nagaland for the guidance and help in this research work.

References

Williams-Johnson

World Health Organization Geneva 1999

World Health Organization (WHO)

Bethesda

HSDB (2001) Manganese compounds

2001 http://toxnet.nlm.nih.gov/cgi-bin/sis/htmlgen?HSDB

Lynam

D. R.

Roos

J. W.

Pfeifer

G. D.

Fort

B. F.

Pullin

T. G.

Environmental effects and exposures to manganese from use of methylcyclopentadienyl manganese tricarbonyl (MMT) in gasoline

Neurotoxicology 1999 20

5597299

2–3 145 150

10385878

Dietrich

A. M.

Burlingame

G. A.

Critical review and rethinking of USEPA secondary standards for maintaining organoleptic quality of drinking water

Environmental Science & Technology 2015 49

5597299

2 708 720

10.1021/es504403t

2-s2.0-84921818303

25517292

Tay

C. K.

Hayford

Levels, source determination and health implications of trace metals in groundwater within the Lower Pra Basin, Ghana

Environmental Earth Sciences 2016 75

5597299

10.1007/s12665-016-6034-2

2-s2.0-84987619224

World Health Organization

Guidelines for Drinking-Water Quality, Fourth Edition

2011 2020, https://www.who.int

Esfandiar

Nasernejad

Ebadi

Removal of Mn(II) from groundwater by sugarcane bagasse and activated carbon (a comparative study): application of response surface methodology (RSM)

Journal of Industrial and Engineering Chemistry 2014 20

5597299

5 3726 3736

10.1016/j.jiec.2013.12.072

2-s2.0-84903531736

2000 A

Toxicological profile for manganese

ATSDR's Toxicological Profiles 2002

Agency for Toxic Substances and Disease Registry, U.S. Department of Health and Human Services (Public Health Service)

10.1201/9781420061888_ch108

Freeland-Graves

J. H.

Bales

C. W. B. F.

Kies

Manganese requirements of humans

Nutritional bioavailability of manganese 1987

ACS Publications

90 104

10.

Han

Zhao

Z.-W.

Gao

Cui

F.-Y.

Study on the factors affecting simultaneous removal of ammonia and manganese by pilot-scale biological aerated filter (BAF) for drinking water pre- treatment

Bioresource Technology 2013 145

5597299

17 24

10.1016/j.biortech.2013.02.101

2-s2.0-84883452103

23578601

11.

Elsner

R. J. F.

Spangler

J. G.

Neurotoxicity of inhaled manganese: public health danger in the shower?

Medical Hypotheses 2005 65

5597299

3 607 616

10.1016/j.mehy.2005.01.043

2-s2.0-21844450900

15913899

12.

Roels

H. A.

Bowler

R. M.

Kim

Henn

B. C.

Mergler

Hoet

Gocheva

V. V.

Bellinger

D. C.

Wright

R. O.

Harris

M. G.

Chang

Manganese exposure and cognitive deficits: a growing concern for manganese neurotoxicity

NeuroToxicology 2012 33

5597299

4 872 880

10.1016/j.neuro.2012.03.009

2-s2.0-84864522618

22498092

13.

Griffin

A. E.

Significance and removal of manganese in water supplies

Journal - American Water Works Association 1960 52

5597299

10 1326 1334

10.1002/j.1551-8833.1960.tb00599.x

14.

Bean

E. L.

Potable water-quality goals

Journal - American Water Works Association 1974 66

5597299

4 221 230

10.1002/j.1551-8833.1974.tb02011.x

15.

Sri Abirami Saraswathi

Divya

Selvapandian

Mohan

Rana

Nagendran

Permeation and antifouling performance of poly (ether imide) composite ultrafiltration membranes customized with manganese dioxide nanospheres

Materials Chemistry and Physics 2019 231

5597299

159 167

10.1016/j.matchemphys.2019.04.023

2-s2.0-85064430746

16.

Funes

de Vicente

Cruz-Pizarro

de Vicente

The influence of pH on manganese removal by magnetic microparticles in solution

Water Research 2014 53

5597299

110 122

10.1016/j.watres.2014.01.029

2-s2.0-84893518362

24509345

17.

Tekerlekopoulou

A. G.

Vasiliadou

I. A.

Vayenas

D. V.

Biological manganese removal from potable water using trickling filters

Biochemical Engineering Journal 2008 38

5597299

3 292 301

10.1016/j.bej.2007.07.016

2-s2.0-38849176873

18.

Wagloehner

Nitzer-Noski

Kureti

Oxidation of soot on manganese oxide catalysts

Chemical Engineering Journal 2015 259

5597299

492 504

10.1016/j.cej.2014.08.021

2-s2.0-84906841712

19.

Yang

Chu

Ren

Zhang

Distribution and genetic diversity of the microorganisms in the biofilter for the simultaneous removal of arsenic, iron and manganese from simulated groundwater

Bioresource Technology 2014 156

5597299

384 388

10.1016/j.biortech.2014.01.067

2-s2.0-84899085718

24507582

20.

Jia

Liu

Zhong

Zhang

Gong

Manganese oxide coated river sand for Mn(II) removal from groundwater

Journal of Chemical Technology & Biotechnology 2015 90

5597299

9 1727 1734

10.1002/jctb.4524

2-s2.0-84937780172

21.

Abdel Salam

Synthesis and characterization of novel manganese oxide nanocorals and their application for the removal of methylene blue from aqueous solution

Chemical Engineering Journal 2015 270

5597299

50 57

10.1016/j.cej.2015.02.022

2-s2.0-84922985710

22.

Goher

M. E.

Hassan

A. M.

Abdel-Moniem

I. A.

Fahmy

A. H.

Abdo

M. H.

El-Sayed

S. M.

Removal of aluminum, iron and manganese ions from industrial wastes using granular activated carbon and Amberlite IR-120H

Egyptian Journal of Aquatic Research 2015 41

5597299

2 155 164

10.1016/j.ejar.2015.04.002

2-s2.0-84930574846

23.

S.-L.

Jeng

H.-T.

Lai

C.-H.

Characteristics and adsorption properties of iron-coated sand

Water Science and Technology 1997 35

5597299

7 63 70

10.2166/wst.1997.0261

24.

Luo

Tian

Yang

Zhang

Yan

Manganese dioxide/iron oxide/acid oxidized multi-walled carbon nanotube magnetic nanocomposite for enhanced hexavalent chromium removal

Chemical Engineering Journal 2013 234

5597299

256 265

10.1016/j.cej.2013.08.084

2-s2.0-84884369351

25.

Kan

C. C.

Aganon

M. C.

Futalan

C. M.

Dalida

M. L. P.

Adsorption of Mn2+ from aqueous solution using Fe and Mn oxide-coated sand

Journal of Environmental Sciences (China) 2013 25

5597299

7 1483 1491

10.1016/S1001-0742(12)60188-0

2-s2.0-84880718405

26.

Khan

Wahab

Rashid

Khan

Khatoon

Musarrat

A Al-Khedhairy

The use of carbonaceous nanomembrane filter for organic waste removal

Application of Nanotechnology in Water Research 2014

Wiley Blackwell

115 152

10.1002/9781118939314.ch6

2-s2.0-84927137544

27.

Mishra

A. K.

Mishraed

A. K.

Application of Nanotechnology in Water Research 2014

Hoboken, NJ, USA

John Wiley & Sons, Inc.

10.1002/9781118939314

2-s2.0-84927145272

28.

Tobiason

J. E.

Bazilio

Goodwill

Mai

Nguyen

Manganese removal from drinking water sources

Current Pollution Reports 2016 2

5597299

3 168 177

10.1007/s40726-016-0036-2

2-s2.0-84975719950

29.

Zodrow

Brunet

Mahendra

Zhang

Alvarez

P. J.

Polysulfone ultrafiltration membranes impregnated with silver nanoparticles show improved biofouling resistance and virus removal

Water Research 2009 43

5597299

3 715 723

10.1016/j.watres.2008.11.014

2-s2.0-59349116749

30.

Ates

Role of modification of natural zeolite in removal of manganese from aqueous solutions

Powder Technology 2014 264

5597299

86 95

10.1016/j.powtec.2014.05.023

2-s2.0-84901682948

31.

Han

Zou

Daotong

Shi

Jiujun

Removal of copper(II) and lead(II) from aqueous solution by manganese oxide coated sand: II. Equilibrium study and competitive adsorption

Journal of Hazardous Materials 2006 137

5597299

1 480 488

10.1016/j.jhazmat.2006.02.018

2-s2.0-33748040430

16631305

32.

Zhang

Lin

Zhang

Removal of fluoride from groundwater by adsorption onto La(III)- Al(III) loaded scoria adsorbent

Applied Surface Science 2014 303

5597299

1 5

10.1016/j.apsusc.2014.01.169

2-s2.0-84898005487

33.

Babel

Kurniawan

T. A.

Low-cost adsorbents for heavy metals uptake from contaminated water: a review

Journal of Hazardous Materials 2003 97

5597299

1-3 219 243

10.1016/S0304-3894(02)00263-7

2-s2.0-0037469982

12573840

34.

Doula

M. K.

Removal of Mn2+ ions from drinking water by using clinoptilolite and a clinoptilolite-Fe oxide system

Water Research 2006 40

5597299

17 3167 3176

10.1016/j.watres.2006.07.013

2-s2.0-33748297439

16938338

35.

Erdem

Karapinar

Donat

The removal of heavy metal cations by natural zeolites

Journal of Colloid and Interface Science 2004 280

5597299

2 309 314

10.1016/j.jcis.2004.08.028

2-s2.0-7944220988

15533402

36.

J. C.

Chen

Huang

X. F.

G. M.

Liu

Yang

Gao

S. N.

Iron and manganese removal by using manganese ore constructed wetlands in the reclamation of steel wastewater

Journal of Hazardous Materials 2009 169

5597299

1–3 309 317

10.1016/j.jhazmat.2009.03.074

2-s2.0-67649819400

19443107

37.

Alexander

J. A.

Zaini

M. A. A.

Abdulsalam

Aliyu El-Nafaty

Aroke

U. O.

Isotherm studies of lead(II), manganese(II), and cadmium(II) adsorption by Nigerian bentonite clay in single and multimetal solutions

Particulate Science and Technology 2019 37

5597299

4 403 413

10.1080/02726351.2017.1404514

2-s2.0-85042220368

38.

Jeż-Walkowiak

Dymaczewski

Szuster-Janiaczyk

Nowicka

Szybowicz

Efficiency of Mn removal of different filtration materials for groundwater treatment linking chemical and physical properties

Water (Switzerland) 2017 9

5597299

7 498

10.3390/w9070498

2-s2.0-85025439211

39.

Mohamadreza

Maryam

Absorption isotherm study of Mn2+ on MnO2 and FeO - coated zeolite from aqueous solution

International Journal of Advanced Science and Technology 2014 72

5597299

63 72

10.14257/ijast.2014.72.06

40.

Taffarel

S. R.

Rubio

Removal of Mn2+ from aqueous solution by manganese oxide coated zeolite

Minerals Engineering 2010 23

5597299

14 1131 1138

10.1016/j.mineng.2010.07.007

2-s2.0-77958091759

41.

Itodo

Gafar

Estimation of specific surface area using Langmuir isotherm method

Journal of Applied Sciences and Environmental Management 2011 14

5597299

4 1 5

10.4314/jasem.v14i4.63287

42.

Zohreh

Abedi

Efficacy of impregnated active carbon in manganese removal from aqueous solutions

Journal of Advances in Environmental Health Research 2019 7

5597299

113 121

10.22102/JAEHR.2019.128628.1074

43.

Rose

Hager

Glas

Rehmann

Hofmann

Coating techniques for glass beads as filter media for removal of manganese from water

Water Science and Technology: Water Supply 2017 17

5597299

1 95 106

10.2166/ws.2016.116

2-s2.0-85013297998

44.

Dubey

Bag

D. S.

Varadan

V. K.

Lal

Mathur

G. N.

Polyaniline coating on glass and PMMA microspheres

Reactive and Functional Polymers 2006 66

5597299

4 441 445

10.1016/j.reactfunctpolym.2005.09.006

2-s2.0-33645046161

45.

ICCD

The International Centre for Diffraction Data

2012 2020, http://www.icdd.com

46.

Khadse

G. K.

Patni

P. M.

Labhasetwar

P. K.

Removal of iron and manganese from drinking water supply

Sustainable Water Resources Management 2006 1

5597299

2 157 165

10.1016/j.watres.2006.07.013

2-s2.0-33748297439

47.

Okeola

O. F.

Odebunmi

E. O.

Ameen

O. M.

Adebayo

G. B.

Adsorption of iron (III) ion from aqueous solution using Jatropha adsorption of iron (III) ion from aqueous solution using Jatropha curcas seed coat activated carbon 2011

Ilorin, Nigeria

KWA, Chemistry Department, University of Ilorin

48.

Al-Ghouti

M. A.

Da’ana

D. A.

Guidelines for the use and interpretation of adsorption isotherm models: a review

Journal of Hazardous Materials 2020 393

5597299

10.1016/j.jhazmat.2020.122383

122383

49.

Chen

Modeling of experimental adsorption isotherm data

Information 2015 6

5597299

1 14 22

10.3390/info6010014

2-s2.0-84924626392

50.

Foo

K. Y.

Hameed

B. H.

Insights into the modeling of adsorption isotherm systems

Chemical Engineering Journal 2010 156

5597299

1 2 10

10.1016/j.cej.2009.09.013

2-s2.0-72049131056

51.

Gholizadeh

Kermani

Gholami

Farzadkia

Kinetic and isotherm studies of adsorption and biosorption processes in the removal of phenolic compounds from aqueous solutions: comparative study

Journal of Environmental Health Science and Engineering 2013 11

5597299

1 1 10

10.1186/2052-336X-11-29

2-s2.0-84910092337

24355013

52.

Freundlich

H. M. F.

Over the adsorption in solution

Journal of Physical Chemistry 1906 57

5597299

385 471