Adsorption of Lithium from Shale Gas Produced Water Using Titanium

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Applied Chemistry

Adsorption of lithium from shale gas produced water using titanium based adsorbent YunJai Jang, and Eunhyea Chung Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.8b00805 • Publication Date (Web): 05 Jun 2018 Downloaded from http://pubs.acs.org on June 5, 2018

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Industrial & Engineering Chemistry Research

Adsorption of lithium from shale gas produced water using titanium based adsorbent Yunjai Jang,† Eunhyea Chung‡,*

†

Department of Energy Systems Engineering, Seoul National University, Seoul, 08826, Republic of Korea; ‡

Department of Energy Resources Engineering · Research Institute of Energy and Resources, Seoul National University, Seoul, 08826, Republic of Korea

*Corresponding author. Email: [email protected]; Tel: +82-2-880-7225

Abstract

Lithium is a valuable metal that has been recovered from synthetic shale gas produced water using the titanium-based adsorbent H2TiO3 in this study. The maximum adsorptive capacity of lithium obtained was 2.58 mmol/g after adsorption-desorption tests, and the recovery rate of Li+ was much higher than those of the other cations in the produced water when used with a pH buffer. To enhance the adsorptive capacity and selectivity of lithium, the precipitation process using sodium carbonate was applied prior to adsorption. More than 96% of the divalent cations

ACS Paragon Plus Environment

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Industrial & Engineering Chemistry Research 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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such as barium, calcium and strontium were precipitated. From the precipitation-adsorptiondesorption test, the lithium desorption capacity obtained for the produced water was 3.61 mmol/g, with a higher selectivity compared to the other cations. Thus, adsorption using a titanium-based adsorbent might be a promising method for lithium recovery from shale gas produced water, when used in conjunction with precipitation.

1. Introduction Lithium, which is one of the most valuable energy resources, is mainly used in ceramics, glasses, and lubricating greases.1 In particular, it is an essential element in a wide range of applications such as energy-storage systems for electronic devices. The demand for lithium is expected to increase up to 160% until 2020 owing to the increasing production and usage of electric devices, including vehicles.2 It is therefore important to produce lithium from various sources to satisfy the increasing demand, since the global lithium reserves are limited. The main source of lithium includes minerals such as pegmatites and brines. For example, lithium is only produced from brines in the United States.1 Highly concentrated lithium is also found in the salt lake brines in Bolivia, Chile, and Argentina.2 The lithium concentration of seawater is known to be 0.17 mg/L on average,3 so the ocean is sometimes considered as another reserve of lithium owing to its large volume and consequently high availability of the metal. However, it might not be an ideal source of lithium considering its low concentration.4 Shale gas produced water can be suggested as another possible reserve of lithium. The shale gas produced water is one of the byproducts of shale gas extraction and is a mixture of fracturing fluid and formation water. It is known to have high total dissolved solid (TDS) levels, with high and


2

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varying lithium concentrations. In the case of Marcellus shale gas produced water, the median concentrations of TDS and lithium are about 157,000 mg/L and 95 mg/L, respectively.5 The development of hydraulic fracturing and horizontal drilling techniques rendered the massive production of shale gas feasible, and the amount of shale gas produced water has been increasing, with recent studies highlighting shale gas as an unconventional natural gas resource. In particular, the production of shale gas and tight oil is expected to double from 2015 to 2040.6 Considering the amount of shale gas produced water and its lithium concentration, it might be a promising source of lithium if the lithium could be selectively recovered. Therefore, in this study, shale gas produced water has been chosen as a target solution for lithium recovery. Lithium can be recovered by several methods, and those commonly used include solvent extraction, electrochemical extraction, and adsorption. Solvent extraction involves a simple procedure and its operation time is relatively short compared to the other methods. However, several divalent cations have shown stronger affinities for the solvent than lithium, therefore, it is not a suitable method for lithium recovery when the target solution contains significant amounts of divalent cations.7 The electrochemical method showed a high selectivity for the lithium ion, however, it appeared that its uptake capacity was lower than that of the adsorption method.8 For lithium adsorption from aqueous solutions, manganese- and titanium-based adsorbents are mainly used. Manganese oxide adsorbents have shown high adsorptive capacity and selectivity for the lithium contained in seawater.9,10 There have also been studies on lithium adsorption from brines using titanium-based adsorbents, and high adsorptive capacity and selectivity for lithium have been observed.11 The titanium-based adsorbent has been chosen for lithium adsorption from shale gas produced water because it has a relatively stable structure of metal loss during acid treatment, while exhibiting a great selectivity for lithium.12 It is reported that the H2TiO3


3


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adsorbent can adsorb lithium through the ion-exchange method.11,12 The lithium ion in the solution is exchanged with the proton in the adsorbent and ions having larger radii than lithium, such as sodium ions, cannot interact with the exchange site of the adsorbent due to steric effects.13 Therefore, these types of adsorbents typically exhibit high selectivity for the lithium ion, so H2TiO3 was chosen as the lithium adsorbent for the produced water used in this study. There has been little research on the adsorption of lithium from shale gas produced water. Therefore, in this study, we present lithium adsorption using H2TiO3 combined with precipitation to enhance the adsorptive capacity of lithium with high selectivity.

2. Materials and methods 2.1 Preparation of synthetic shale gas produced water and adsorbent Shale gas produced water was synthesized using the ion concentrations and pH values from the reference of Marcellus shale gas produced water.5 The produced water is a mixture of fracturing fluid and formation water and, here, they were assumed to be mixed in a 1:1 volumetric ratio. In this study, the majority of organic compounds in the produced water were not considered, but the organic compounds that exist in the fracturing fluid were only considered during the synthesis of the produced water. The fracturing fluid was synthesized using the main components and their volumetric compositions according to the reference (Table S1 in Supporting Information).14,15 The chemicals used for synthesis of fracturing fluid include polyacrylamide, isopropanol, guar gum and sands. The pH level and concentrations of the inorganic ions such as lithium, barium, calcium, strontium, magnesium, sodium, and chloride in the produced water were adjusted based on the reference values.5 Table 1 shows the concentrations of the major ions in the synthetically produced water. Since there are significant amounts of suspended materials, including sand, in


4

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the solution, they were allowed to settle down for one day, and the supernatant was only used for the precipitation and adsorption experiment. Table 1. Concentrations of synthetic and real shale gas produced waters from Marcellus Synthetically produced water (mg/L)

Flowback water (mg/L)*

Total dissolved solids

98800

3,010–228,000

Ca

8,580–11,500

204-14,800

K

369-398

8-1,010

Li

70–140

4–202

Mg

813-853

22–1,800

Na

23,800-25,300

1,100–44,100

Sr

2,110-2240

46–5,350

Ba

1,670-1,740

76–13,600

*

Haluszczak et al. (2013)5

To synthesize the H2TiO3 adsorbent, powders of Li2CO3 and TiO2 were well mixed in a 2:1 molar ratio of Li:Ti and heated to 700 °C in a furnace for 4 h. Then, the Li2TiO3 powder was stirred in 0.2 M HCl for 24 hrs at room temperature to form H2TiO3. The H2TiO3 was then washed with deionized water, filtered, and dried at room temperature. As part of characterization, the X-ray diffraction (XRD; Rigaku, high-resolution X-ray diffractometer and powder X-ray diffractometer) patterns of the synthesized Li2TiO3 and H2TiO3 samples were analyzed to confirm their chemical compositions. The morphologies of Li2TiO3 and H2TiO3 were observed using a field-emission scanning electron microscope (Carl Zeiss, Sigma)

2.2 Lithium adsorption and desorption in pH buffered produced water


5


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From the preliminary tests, a rapid decline in pH was observed after the adsorbent was added to the solution due to lithium ion exchange with hydrogen. Therefore, KHCO3 was used as a pH buffer in the produced water, so that the effect of the change in pH could be diminished. Kinetic studies of the H2TiO3 adsorbent in the pH buffered produced water were carried out and the adsorptive capacity of lithium was tested for 48 hrs. In addition, various molar ratios of hydrogen-to-titanium in the adsorbent were considered. The adsorbents were synthesized using different molar ratios of Li: Ti for Li2CO3 and TiO2, from 1.8: 1 to 2.3: 1. Then, the adsorptive capacity of lithium in the buffered solution was investigated for 48 hrs. All the adsorption experiments were conducted at 30 °C with 150 rpm using a shaking incubator (SH Scientific, SH-BSI-16R) with 0.03 g of the adsorbent and 10 mL produced water. Each sample was filtered with a 0.45 µm syringe-driven filter before chemical analysis. The concentrations of cations, including lithium, in the initial produced water and equilibrium solutions were determined by inductively coupled plasma optical emission spectrometer (ICP-OES; Perkin Elmer, Optima 8300). To observe the effect of adsorbent dosage on adsorptive capacity, an isothermal study was performed using 10 mL of the buffered solution with 0.01, 0.03, 0.05, 0.1, and 0.2 g of adsorbents, and the lithium adsorptive capacity after 24 hrs was analyzed. The adsorptive capacity was calculated using the following equation. =

− ×

1

: adsorptive capacity (mmol/g),

: volume of produced water (L), : weight of adsorbent (g), : initial concentration (mmol/L), and


6

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: concentration at equilibrium (mmol/L). Additionally, the removal rate (%) of each cation was calculated using the following equation. Removal rate =

− × 100% 2

The distribution coefficient, separation factor, and concentration factor were calculated using the following equations. =

−

× 3

=

& =

!" 4 #$

#$ 5 #$

For the desorption test, the adsorbents used in the studies with 10 mL of solution with 0.03 g of adsorbents were washed several times with deionized water and stirred in 40 mL of 0.2 M HCl for a day before analyzing the cation concentrations. The recovery rate (%) and desorption capacity (mmol/g) of the cations were calculated using the following equations. Desorption capacity =

/ 6

/ : Amount of cations desorbed (mmol) Recovery rate =

/ × 100% 7 /12345 26 785"249 "4 :;2

2.3 Precipitation-adsorption-desorption test In order to eliminate the inhibitive effect of divalent cations on the adsorption of lithium from the produced water, Na2CO3 was used as a precipitation reagent. The required amount of


7


carbonate was calculated based on the sum of the moles of the divalent cations (Ca2+, Ba2+, Mg2+, and Sr2+) that are present in the shale gas produced water sample. Various molar ratios of the carbonate to the sum of divalent cations (Ba, Ca, Mg, and Sr) from 0.8:1 to 1.4:1 were considered to observe the effect of Na2CO3 dose on precipitation efficiency. Further, different reaction times, up to 1 h, were tested to observe the precipitation rates of the divalent cations over time. After removing the divalent cations by precipitation, the adsorption and desorption test was performed with the precipitated produced water in the same experimental conditions as that used for the buffered produced water.

3. Results and discussion

10

30

) (

(200)

(202)

Characterization of Li2TiO3 and H2TiO3 (020) (002)

3.1

Intensity (arbitrary unit)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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50

70

2-theta Figure 1. XRD patterns of Li2TiO3 (upper) and H2TiO3 (lower) The XRD patterns of Li2TiO3 and H2TiO3 are shown in Figure 1. The peaks in the upper pattern indicate that Li2TiO3 was synthesized from Li2CO3 and TiO2. In addition, as indicated by


8

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the lower pattern, H2TiO3 was formed through ion-exchange between the lithium in Li2TiO3 and the protons of the acid during HCl treatment. The peaks of Li2TiO3 and H2TiO3 showed similar patterns, with identical (002), (020), and (200) peaks. However, H2TiO3 exhibited a lower @ ) peaks of Li2TiO3 were intensity for the (002) peak than Li2TiO3, while the (202) and (33 1

absent due to the insertion of H+ instead of Li+ during HCl treatment, which was also reported in other studies.11,12,16

(b)

(a)

Figure 2. SEM images of (a) Li2TiO3 and (b) H2TiO3 The SEM images of Li2TiO3 and H2TiO3 in Figure 2 show that the particles were clustered and had a relatively uniform size of about 100–200 nm, with H2TiO3 revealing somewhat smaller sizes compared to Li2TiO3.

3.2

Adsorption characteristics of H2TiO3 in pH buffered produced water

Since the extraction of lithium from the raw produced water in the preliminary test revealed a low adsorption efficiency due to the decrease in pH, KHCO3 was used as a buffer to fix the pH level of the solution. When this buffer was not used, the pH of the produced water after adsorption decreased to 4.5 and minimal lithium was adsorbed. However, by using KHCO3 as


9


the buffer, the pH of the produced water before adsorption was 7.0, and after the adsorption, it became 6.4; the adsorptive capacity also increased significantly. 3.2.1 Kinetics of lithium adsorption in pH buffered produced water The lithium adsorptive capacity over 48 hrs has been plotted in Figure 3. The adsorptive capacity significantly increased for the first 8 h and reached equilibrium after about 24 hrs. The adsorptive capacity after 24 hrs was 2.16 mmol/g, which became 2.36 mmol/g after 48 hrs. 2.5 Adsorptive capacity (mmol/g)

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2.0 1.5 1.0 0.5 0.0 0

20

40

60

Time (h)

Figure 3. Change in the lithium adsorptive capacity with time for buffered produced water

3.2.2 Isotherms of lithium in pH buffered produced water Figure 4 describes the change in the equilibrium adsorptive capacity for different adsorbent doses of 1, 3, 5, 10, and 20 g/L. As the H2TiO3 dosage increased, the rate of lithium removal (%) from the produced water increased as well. The highest removal rate of 92.7% and the lowest adsorptive capacity of 0.638 mmol/g were observed when 20 g/L of the adsorbent dose was used. The removal rates seem to be stabilized at 20 g/L of the adsorbent dose, which implied that lithium in the solution is depleted. The adsorptive capacity of lithium increased as a smaller


10

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amount of H2TiO3 was used, reaching the maximum value of 2.64 mmol/g for 1 g/L of

3

100 90 80 70 60 50 40 30 20 10 0

2.5 2 1.5 1 0.5 0 0

10 20 Adsorbent dose (g/L)

Removal rate (%)

adsorbent.

Adsorptive capacity (mmol/g)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Adsorptive capacity (mmol/g) Removal rate (%)

30

Figure 4. Lithium adsorptive capacity and removal rate for buffered produced water In this study, Langmuir and Freundlich isotherm models were used to fit the experimental data. The isotherm equations are described below. 1 1 = + 7 A AB AB : equilibrium concentration of adsorbate (mmol/L), A : amount of metal adsorbed per gram of adsorbent at equilibrium (mmol/g), AB : maximum monolayer coverage capacity (mmol/g), and : Langmuir isotherm constant (L/mmol). log A =

1 log + log E 8 4


11


4: adsorption intensity, and E : Freundlich isotherm constant.

Figure 5 shows the Li+ adsorption isotherms of the pH buffered produced water. As shown in Figure 5, the fit of the Langmuir isotherm model displayed a R2 value of 0.989, with theoretical maximum adsorptive capacity of 3.78 mmol/g and =0.200. The fitting of the Freundlich isotherm model resulted in the R2 value of 0.981, with 4 =1.72 and E =1.55. The R2 values of both the isotherms are relatively high, though the value for the Langmuir model is slightly higher than that of the Freundlich model. The highest adsorptive capacity at equilibrium observed in Figure 4 is less than the obtained AB value, which might indicate that the maximum adsorptive capacity of the adsorbent was not reached during the experiments. However, the result showing a better fit to the Langmuir isotherm implies that lithium adsorption to H2TiO3 occurs in the form of homogeneous monolayer adsorption. 5.0

0.5

(a)

(b)

0.4

4.0

0.3 R² = 0.9886

3.0

log $

Ce/$

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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2.0

0.2 0.1 0 -0.1 0

1.0

R² = 0.9812

0.5

1

1.5

-0.2 0.0

-0.3 0

5

10

15 log Ce

Ce

Figure 5. Fits of Langmuir (a) and Freundlich (b) adsorption isotherms

3.2.3 Adsorbents with different Li: Ti molar ratios


12

Page 13 of 31

It was found that the Li: Ti molar ratio used for synthesis of the adsorbent would affect the adsorptive capacity of lithium.12 In this study, therefore, different ratios of lithium to titanium were considered during the synthesis of the adsorbent to observe the effect of the molar ratio.

2.5 Adsorptive capacity (mmol/g)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


2.0

1.8:1

1.5

1.9:1 2.0:1

1.0

2.2:1 2.3:1

0.5

0.0 0

10

20 30 Time (h)

40

50

Figure 6. Lithium adsorptive capacities of H2TiO3 for different molar ratios of Li:Ti during synthesis Figure 6 indicates that 2:1 molar ratio resulted in the highest adsorptive capacity during the entire reaction period. There does not seem to be much difference in the adsorptive capacity of lithium between the adsorbents with different molar ratios in the first two hours of the kinetic study. However, the adsorptive capacity after 24 hrs showed differences for each adsorbent, depending on the ratio. The lithium adsorptive capacity after 24 hrs followed the order 2.0:1 > 2.2:1 > 1.8:1 > 2.3:1 > 1.9:1 of the Li: Ti molar ratio used for the synthesis of the adsorbent. It has been reported in other study that as the molar ratio of lithium to titanium increases from 2.0:1


13


Page 14 of 31

to 2.16:1, the adsorptive capacity of lithium decreases.12 This study showed similar result where the adsorptive capacity of lithium decreases as the ratio of lithium to titanium increases from 2.0:1 to 2.3:1. After 48 hrs, the adsorptive capacity of the adsorbent with the molar ratio of 1.9:1 was the least and it was 20% lower than that of the adsorbent with 2.0:1 molar ratio. These results may indicate that the Li: Ti molar ratio in the adsorbent synthesis could affect the lithium adsorptive capacity.

3.2.4 Adsorption-desorption of lithium in pH buffered produced water The equilibrium concentrations of cations, including lithium, were determined after 24 hrs of the adsorption reaction, and, accordingly, the distribution coefficient, concentration factor and separation factors were calculated, as listed in Table 2. The adsorptive capacity of lithium at equilibrium was 2.43 mmol/g. The concentration changes of several cations such as Ba, Ca, and K during the reaction were negligible, which indicates that those ions were not adsorbed on to H2TiO3. The distribution coefficients ( ) and concentration factors (&) of the cations in the buffered produced water are in the order Li+ >> Sr2+ > Mg2+, and the separation factors for lithium over Mg2+ and Sr2+ were significantly high. This high selectivity of lithium over magnesium and strontium is due to the ion-exchange property of H2TiO3, allowing only the lithium ions to be chemically adsorbed onto the adsorption sites present in the adsorbent. Lithium has an ionic radius smaller than those of Na+, K+, and other divalent cations that are present in the produced water, therefore, those cations cannot enter the exchange sites of the adsorbent. This indicates that Li+ has high selectivity compared to the interfering divalent cations during the adsorption process. Table 2. Adsorption selectivity of buffered produced water


14

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(mmol/L)

(mmol/L)

(mmol/g)

& (×103)

Ba2+

12.4

12.4±0.1

-

-

-

-

Ca2+

135

135±20

-

-

-

-

K+

199

199±47

-

-

-

-

Li+

13.3

6.01

2.43

404

183

1.00

Mg2+

36.4

35.5

0.293

8.25

8.05

48.9

Na+

889

889±181

-

-

-

-

Sr2+

25.3

24.7

0.211

8.52

8.31

47.4

The desorption of cations from the adsorbents was investigated after the solution reaches equilibrium. In the desorption test, the adsorbents were rinsed with deionized water before placement in HCl, so that the adsorbed cations could be replaced by the protons. The desorption test results shown in Table 3 indicate that the recovery rates of the cations were in the order Li+ >> Sr2+ > Mg2+. Table 3. Desorption capacity of ions from buffered produced water Desorption capacity (mmol/g)

Recovery rate (%)

Li+

2.58

58.3

Mg2+

0.234

1.93

Sr2+

0.174

2.06

Na+

0.245

0.083

The final rate of lithium recovery from the produced water after the adsorption-desorption process was 58.3%, which was significantly higher than those of the other cations. However, the amounts of Mg2+ and Sr2+ recovered along with lithium were not negligible, because the initial concentrations of those cations in the synthetic produced water were much higher than the


15


Page 16 of 31

lithium concentration. In addition, during the rinsing step before desorption, large amounts of those cations were observed in the rinsed water, therefore the removal of such divalent cations from the produced water before the adsorption would be necessary for selective lithium recovery. Therefore, the precipitation test using sodium carbonate was conducted prior to the adsorptiondesorption process.

3.3

Precipitation-adsorption-desorption of lithium from produced water

In Table 3, the desorption capacity values of several cations including Mg2+ and Na+ were not negligible. It indicates that a certain amount of those ions is successively removed by H2TiO3 adsorbent through ion exchange reaction. Considering that their concentrations in the original solution were significantly higher than the lithium concentration, the removal of competitive cations is required to enhance the selectivity of lithium over other ions. Sodium carbonate was used as the precipitating reagent in the shale gas produced water for the selective recovery of lithium over divalent cations. The optimal amount of the precipitant was determined after testing various molar ratios of the carbonate to the sum of divalent cations, including Ba, Ca, Mg, and Sr. Table 4 shows the removal rate (%) of the divalent cations in the produced water when sodium carbonate is used with different doses and reaction times. Table 4. Removal rates (%) of divalent cations for different Na2CO3 doses and reaction times Na2CO3 dosage ratio

Reaction time (min)

0.8:1

1:1

1.2:1

1.4:1

10

20

30

60

Ba2+

46.5

96.9

99.9

100.0

98.8

100.0

99.9

99.9

Ca2+

85.1

98.5

99.6

99.6

98.4

99.6

99.7

99.6


16

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Mg2+

41.1

42.1

67.6

75.6

55.2

54.0

67.2

67.6

Sr2+

65.9

77.9

99.1

99.3

96.1

98.7

99.0

99.1

Li+

0.34

-3.86

-1.60

4.26

3.26

2.60

6.33

-1.60

Ninety nine percent of Ba, Ca, and Sr were precipitated when a large amount of Na2CO3 (molar ratio of 1.2:1 or higher) was used for 60 min and the change in lithium concentration due to precipitation was insignificant. Further, to determine the appropriate reaction time, various times was considered for the precipitant, whose molar ratio was 1.2:1. As shown in Table 4, a duration of 10 min might be sufficient to remove over 95% of Ba, Ca, and Sr ions. Therefore, a molar ratio of 1.2:1 and reaction duration of 10 min were determined to be optimal for the precipitation reaction followed by adsorption-desorption involving the synthetic shale gas produced water. The pH of the produced water after precipitation increased to 10.2, which was much higher than those of the initial and pH buffered produced water, due to the addition of sodium carbonate. Table 5 showed the adsorption results of four cations including K+, Li+, Mg2+, and Na+. Over 96% of Ba, Ca, and Sr ions in the produced water were removed during the precipitation step, so the adsorption of these ions was negligible. The pH of the solution decreased from 10.2 before adsorption to 9.7 after adsorption. As shown in Table 5, the equilibrium adsorptive capacity of lithium was 3.30 mmol/g which is higher than that in the buffered produced water (2.43mmol/g). The increase in the adsorptive capacity might be partially due to the pH increase of the solution. It has been reported in other studies that the adsorptive capacity of H2TiO3 for lithium increases with the pH level of the solution.11,16 However, our previous study revealed that when the pH of the produced water was increased using KOH to similar level of the precipitated produced water (pH 10.1), the adsorptive capacity for lithium was 1.40 mmol/g.17 Therefore, the removal of


17


Page 18 of 31

competitive cations as well as the pH increase could play a significant role on the increase of the adsorption efficiency when the precipitation process was applied. Table 5. Adsorption of cations after precipitation

(mmol/L)

(mmol/L)

(mmol/g)

& (×103)

K+

12.3

12.2

0.04

0.004

0.09

106

Li+

18.7

8.74

3.30

0.38

25.5

1.00

Mg2+

12.9

1.02

3.97

3.89

12.6

0.10

Na+

1870

1840

10.7

0.01

0.25

64.6

The separation factors of lithium over monovalent cations in precipitated produced water was high and the distribution coefficients and concentration factors of the monovalent ions except lithium were significantly low, which indicates that the lithium ion has a higher selectivity than Na+ and K+. However, significantly high values for Mg2+ and Na+ were also observed, which indicates that significant amounts of these ions were removed from the produced water by the adsorbent. In particular, the high distribution coefficients and concentration factors of Mg2+, as well as the low separation factor of Li+ relative to Mg2+, imply that lithium does not have a significantly higher selectivity than magnesium during adsorption. The adsorption selectivity calculated with mass unit are shown in Table S2 (in Supporting Information) and the adsorptive capacity values of K+, Li+, Mg2+, and Na+ are 1.71, 22.9, 96.4, 247 mg/g, respectively. Factors such as , &, and had the same value as in Table 5.

It seems that the significant amount of cations including Mg2+ and Na+ were removed from the solution. However, as previously mentioned, the adsorbents were rinsed before the desorption reaction. Those cations had been rinsed away with deionized water during the rinsing process, and thus, significantly high concentrations of Mg2+ and Na+ were found in the aqueous solution


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after rinsing the adsorbents. It would indicate that these ions were weakly adhered to the adsorbent surface which did not occur through an ion exchange reaction, so they could be easily detached from the adsorbent with rinsing. Table 6 shows the desorption capacity and recovery rate of each ion, as obtained from the precipitation-adsorption-desorption tests. Table 6. Desorption capacity of precipitated produced water Desorption capacity (mmol/g)

Recovery rate (%)

K+

0.01

0.19

Li+

3.61

54.4

Mg2+

0.02

0.17

Na+

0.24

0.06

From the result of the adsorption and desorption studies, it can be seen that the cations that seemed to be adsorbed in much larger amounts than lithium were actually desorbed in much smaller amounts compared to lithium. For example, the Na+ adsorptive capacity calculated from the adsorption test was 10.7 mmol/g, however, this decreases significantly to 0.24 mmol/g based on the desorption test, indicating that most of the Na+ were adsorbed weakly to the surface of the adsorbent and were subsequently rinsed away with deionized water. The result implies that, during the adsorption reaction, lithium was adsorbed to the H2TiO3 in the produced water through the ion-exchange reaction, with a high selectivity. The lithium desorption capacity of 3.61 mmol/g was found to be higher than that obtained from the adsorption-desorption test using buffered solution (2.58 mmol/g). However, the desorption capacities of the divalent cations decreased after the application of the precipitation process. For example, the desorption capacity of Mg2+ in the precipitated produced water (0.02


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mmol/g) is much lower than that of buffered produced water (0.234 mmol/g). The desorption capacity of Na+ did not change after the precipitation is introduced. This indicates that the Na+ concentration increase in the solution due to the use of Na2CO3 does not significantly influence on the Na+ removal through the ion exchange reaction. Therefore, it could be concluded that the precipitation process enhanced the selectivity of lithium adsorption using H2TiO3. The lithium adsorptive and desorption capacities in the produced water were lower compared with the lithium adsorptive capacity reported in other studies, e.g., a value of 4.7 mmol/g was reported for lithium adsorption from salt lake brine.11 The desorption capacities of the ions followed the order Li+ >> Na+ > Mg2+ > K+. Finally, 54.4% of lithium in the synthetic shale gas produced water was recovered through the precipitation-adsorption-desorption process. Compared with the results of the adsorption-desorption test of buffered produced water, the recovery rate of lithium from precipitated produced water was slightly lower, though the recovery rates of other cations were significantly decreased (

Adsorption of Lithium from Shale Gas Produced Water Using Titanium

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