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Separation and recovery of iron impurities from a complex oxalic acid solution containing vanadium by K3Fe(C2O4)3·3H2O crystallization
Separation and Purification Technology 232 (2020) 115970
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Separation and recovery of iron impurities from a complex oxalic acid solution containing vanadium by K3Fe(C2O4)3·3H2O crystallization
T
Zishuai Liua,b,c,d, Jing Huanga,b,c,d, , Yimin Zhanga,b,c,d,e, , Tao Liua,b,c,d, Pengcheng Hua,b,c,d, Hong Liua,b,c,d, Qiushi Zhenga,b,c,d ⁎
⁎
a
School of Resource and Environmental Engineering, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China State Environmental Protection Key Laboratory of Mineral Metallurgical Resources Utilization and Pollution Control, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China c Hubei Collaborative Innovation Center for High Efficient Utilization of Vanadium Resources, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China d Hubei Provincial Engineering Technology Research Center of High Efficient Cleaning Utilization for Shale Vanadium Resource, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China e School of Resource and Environmental Engineering, Wuhan University of Technology, Wuhan 430070, Hubei Province, China b
ARTICLE INFO
ABSTRACT
Keywords: Iron impurities Recovery K3Fe(C2O4)3·3H2O Crystallization Oxalic acid Vanadium
Vanadium and iron have been effectively separated from the acid leachate of vanadium-bearing shale by solvent extraction or ion exchange; however, in these methods, the iron was not recovered, and the comprehensive utilization rates of the resources were low. In this paper, the recovery of iron impurities from a complex oxalic acid solution containing vanadium via the crystallization of K3Fe(C2O4)3·3H2O was investigated. Using the optimum conditions, 90.7% of the Fe was recovered, and only 1.4% of the V and 5.0% of the Al were lost. The concentration of Fe decreased from 5366 mg/L before crystallization to 500 mg/L after crystallization. The crystallized product was then further purified by recrystallization to obtain K3Fe(C2O4)3·3H2O in 99.0% purity, which was characterized using FT-IR, XRD, and SEM. Based on the speciation of vanadium and iron in the oxalic acid system, the mechanism by which the iron impurities were recovered was determined. The mechanism consisted of five main steps: complexation, oxidization, coordination, nucleation, and crystal growth. The recovery of iron impurities from a complex oxalic acid solution containing vanadium via the crystallization of K3Fe (C2O4)3·3H2O is a novel and efficient method, and this process provides a new direction for the comprehensive utilization of vanadium resources.
1. Introduction Vanadium-bearing shale is an important vanadium resource in China [1]. With the increasing market demand for vanadium, the efficient and clean extraction of vanadium from vanadium-bearing shale has become an important task [2]. To achieve this goal, oxalic acid, a clean and eco-friendly organic acid, has been used successfully to extract vanadium from vanadium-bearing shale [3–5]. The leaching of the raw shale using oxalic acid leaching selectively extracts V over Fe impurities, but for roasted shale, the iron is easily dissolved along with the vanadium [5], resulting in a high iron concentration in the leachate. However, iron impurities affect the purification and enrichment of vanadium, for instance, by increasing the amount of extractant required, reducing the purity of the vanadium-rich liquid, and influencing
the quality of the resulting vanadium products. Therefore, proper treatment of the iron in vanadium-bearing shale during the vanadium extraction process is very important. In recent decades, the main methods of treating iron in different vanadium- containing solutions have included solvent extraction, ion exchange, and chemical precipitation [6–7]. Solvent extraction has been most widely used to separate vanadium and iron, and various extractants have been reported, including di-2-ethylhexyl phosphoric acid (D2EHPA) [8], 2-ethylhexyl phosphonic acid mono-2-ethylhexyl ester (PC-88A) [9], and trialkylamine (Alamine 336) [10]. However, solvent extraction methods have presented some drawbacks: (1) The valuable metal iron is not recovered; (2) the extractant cost is high; and (3) the extraction process is long. Ion exchange is a more environmentally friendly method for the separation of vanadium from
⁎ Corresponding authors at: School of Resources and Environmental Engineering, Wuhan University of Science and Technology, Wuhan, 430081, Hubei Province, China. E-mail addresses: [email protected] (J. Huang), [email protected] (Y. Zhang).
https://doi.org/10.1016/j.seppur.2019.115970 Received 10 June 2019; Received in revised form 9 August 2019; Accepted 21 August 2019 Available online 22 August 2019 1383-5866/ © 2019 Elsevier B.V. All rights reserved.
Separation and Purification Technology 232 (2020) 115970
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iron, and a variety of resins have been used, including D201 [11], D314 [12], and AG1-x8 [13]. However, ion exchange also has some disadvantages: (1) Iron and other valuable metals are not recovered; (2) the resin can easily be polluted by the impurity ions; and (3) the adsorption capacity of resins is limited. Chemical precipitation is an effective method for the separation and recovery of iron, and numerous researchers have precipitated iron as FeC2O4 from stone coal [4] and red mud [14]. The disadvantages of chemical precipitation include high vanadium losses and poor selectivity [15]. In general, the separation of vanadium and iron from various vanadium-containing solutions can be achieved using the above methods, but the valuable metal iron is not recycled and the comprehensive utilization rate of the vanadiumbearing shale is low. Based on the above analysis, the current processes of treating iron from vanadium-containing solution could be improved in three ways: (1) Developing a process with simple operation and a short flowsheet; (2) recovering iron as high value-added products; and (3) enhancing the selectivity towards iron and decreasing the loss of vanadium. A variety of complexes have been reported to form between oxalate and iron ions in aqueous solution, including Fe(C2O4)33−, Fe(C2O4)22−, FeC2O4+, and others [16]. However, Fe(C2O4)33− readily decomposes into Fe(C2O4)22− or FeC2O4 when the solutions are exposed to strong sunlight [17–18]. This feature can be exploited through use of potassium tris(oxalato)ferrate(III) trihydrate (K3Fe(C2O4)3·3H2O) as a promoter for the photodegradation of textile dyes in water [19]. K3Fe (C2O4)3·3H2O has been extensively synthesized and researched, and has also been used to fabricate functional metal-organic frameworks and nanostructured materials in recent years [20]. A previous report [21] summarized the three most typical methods used to synthesize K3Fe (C2O4)3·3H2O: (1) The oxidation of FeC2O4 by H2O2 in the presence of H2C2O4 and K2C2O4; (2) the direct combination of FeCl3 and K2C2O4·H2O solutions; and (3) the digestion of a mixture of Fe2(SO4)3 and K2C2O4 over a steam bath. The reactions corresponding to the three synthesis methods are given below:
Table 1 Chemical composition of the initial solution (mg/L).
FeCl3(aq) + 3K2C2 O4 ·H2 O(aq) = K3Fe(C2 O4 )3 ·3H2 O(aq) + 3KCl (aq)
Fe(II)
Al
Mg
Ca
K
Na
2064
5366
2756
20,920
2430
1260
2225
728
the shale with 6 mol/kg H2C2O4 and 5% CaF2 at 95 °C for 4 h [3–5]. The vanadium-bearing shale was supplied by Hubei Tongshan Mining Co., Ltd. The pH of the initial solution was 0.45, and its chemical composition is given in Table 1. The initial solution was a complex system due to its high acidity, various impurities, and the organic acid. Deionized water was used in all experiments, and all reagents were of analytical grade. 2.2. Experimental procedures First, 50 mL of the initial solution was placed in a 100 mL beaker and adjusted to the required pH value using KOH. Next, a given amount of K2C2O4 was added to the resulting solution, which was stirred in a water bath at 60 °C for 30 min. The obtained mixed solution was then held at room temperature for 2 h. After measuring its pH value, the solution was filtered to obtain the treated solution. At this time, a cotton thread 1 mm in diameter that had been pretreated by soaking in a saturated aqueous solution of K3Fe(C2O4)3·3H2O for one day and allowed to dry naturally was added to the treated solution. The solution was then allowed to crystallize at low temperature for the desired time. After the crystallization was completed, the crystalline product and the purified solution were separated using filtration. In order to obtain the product in higher purity, the crystalline product was then recrystallized. Finally, the purified solution was analyzed, and the crystallization efficiency (E) and separation factor were calculated as follows:
E =
(1) (2)
=
Fe2 (SO4 )3(aq) + 3K2C2 O4 ·H2 O(aq) + 3BaC2 O4(s) = 2K3Fe(C2 O4 )3 ·3H2 O(aq) + 3BaSO4(s)
TFe*
* Note: TFe stands for total Fe.
2FeC2 O4(s) + H2 C2 O4 ·2H2 O(aq) + 3K2C2 O4 · H2 O(aq) + H2 O2(aq) = 2K3Fe(C2 O4 )3·3H2 O(aq) + H2 O(aq)
V
C0 V0 C1 V1 × 100% C0 V0
Da Db
(4) (5)
where C0 and C1 are the ion concentration in the initial solution and purified solution (mg/L), respectively, and V0 and V1 are the volume of the initial solution and the purified solution (L), respectively. β is the separation factor of the components a and b, and Da and Db are the ratios of the contents of a and b in the crystal to the amount in the purified solution, respectively. Strong light was avoided during the crystallization processes, and the processes were performed one time.
(3)
According to recent reports [3], the concentration of iron in the oxalic acid leachate of shale is 2–3 times that of vanadium, and the iron is present as Fe2+ and Fe3+ ions. Therefore, this iron could be recovered as K3Fe(C2O4)3·3H2O from vanadium-containing oxalic acid solutions. Unfortunately, until now, the recovery of iron via its crystallization as K3Fe(C2O4)3·3H2O from complex, highly acidic solutions containing various ions has not been reported. Therefore, in this paper, a novel and efficient process for the recovery of iron impurities from a complex oxalic acid solution containing vanadium was studied, with the aim of recovering the iron in the form of K3Fe(C2O4)3·3H2O via a crystallization process. The effects of different crystallizing agents and their concentration, the initial pH, and the crystallization time on the crystallization efficiency of iron were investigated. Moreover, the product K3Fe(C2O4)3·3H2O was characterized, and the mechanism of iron recovery was determined.
2.3. Analysis methods Solution pH values were measured using a pH meter (PHS-3C, INESA Scientific Instrument Co., Ltd). The concentrations of all elements were analyzed using inductively coupled plasma-atomic emission spectroscopy (ICP-AES, Optima-4300DV, PerkinElmer, Boston, MA, USA) in addition to ferrous iron (Fe (II)), ferric iron (Fe(III)), and oxalate concentration. The Fe(II) concentration was measured using the 1, 10-phenanthroline spectrophotometry using a spectrophotometer (UV5500, Metash, Shanghai, China) [22–23], and the Fe(III) concentration was calculated from the mass balance. The oxalate concentration of K3Fe(C2O4)3·3H2O was determined by titration with potassium permanganate [24]. The functional groups of the product were analyzed using Fourier transform infrared spectroscopy (FTIR, Thermo Nicolet NEXUS, USA). The phase compositions were determined using X-ray diffraction (XRD, D/MAX 2500PC, Rigaku, Tokyo, Japan).
2. Experimental 2.1. Materials A complex oxalic acid solution containing vanadium (which will be referred to from here on as the initial solution) was prepared by roasting vanadium-bearing shale at 850 °C for 1 h, followed by leaching 2
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given K2C2O4 concentration increased, but even at a K2C2O4 concentration of 1.0 mol/L, the crystallization efficiency of Fe was still not close to equilibrium. As shown in Fig. 2(c), when the initial pH of the solution was 1.5, the crystallization efficiencies of Fe, V, and Al increased from 0 to 91.8%, from 0 to 2.3%, and from 0 to 6.3%, respectively, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, and the equilibrium pH rose from 1.5 to 4.1. At a K2C2O4 concentration of 0.8 mol/L, the crystallization efficiency of Fe was 90.4%. Further increasing the K2C2O4 concentration increased the crystallization efficiency of Fe only slightly and increased the losses of V and Al. Fig. 2(d) shows that when the initial pH of the solution was 2.0, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiencies of Fe, V, and Al increased from 0 to 92.9%, from 0 to 3.2%, and from 0 to 7.8%, respectively, and the equilibrium pH rose from 2.0 to 4.8. At a K2C2O4 concentration of 0.6 mol/L, the equilibrium pH was 3.4, but the crystallization efficiency of Fe was only 83.3%. However, when the K2C2O4 concentration exceeded 0.8 mol/L, the crystallization efficiency of Fe was > 91%, but the losses of V and Al showed obvious increases. Based on the above data, the crystallization efficiency of Fe was determined by the equilibrium pH and K2C2O4 concentration simultaneously. Moreover, a comparison of the four sets of results in Fig. 2 shows that the crystallization efficiency of Fe3+ was always higher than that of Fe2+, which indicated that Fe3+ crystallized more readily than Fe2+. In addition, as the equilibrium pH increased, the crystallization efficiency of V and Al also increased. Based on overall considerations, an initial pH of 1.5 and a K2C2O4 concentration of 0.8 mol/L were selected as the optimal conditions; the corresponding equilibrium pH was 3.0.
Crystallization efficiency of Fe (%)
100 90
85.1
80 70 60 50 40 30 20 10 0
8.5
5.6
6.8
7.3
KCl
K2SO4
KOH
KNO3
K2 C 2 O4
Fig. 1. Effect of the crystallizing agents. (K+: 1.6 mol/L, equilibrium pH value: 3.0, crystallization time: 48 h, crystallization temperature: 5 °C).
3. Results and discussion 3.1. Crystallization of K3Fe(C2O4)3·3H2O 3.1.1. Effect of the crystallizing agents on the separation of V and Fe The effect of different crystallizing agents was studied using a fixed K+ concentration of 1.6 mol/L, an equilibrium pH value of 3.0, a crystallization time of 48 h, and crystallization temperature of 5 °C. Fig. 1 clearly shows that the crystallization efficiency of Fe was the highest (85.1%) when K2C2O4 was used as the crystallizing agent, while that of the other crystallizing agents were all lower than 10%, indicating that K2C2O4 was more likely to form complexes with Fe3+. Piro et al. [21] summarized the three most commonly used synthetic methods for K3Fe(C2O4)3·3H2O, and K2C2O4 was required for all three. This further indicates that K2C2O4 easily complexes with Fe3+ to form K3Fe(C2O4)3·3H2O. Thus, K2C2O4 was selected as the optimum crystallizing agent.
3.1.3. Effect of the crystallization time on the separation of V and Fe Crystallization at low temperatures results in slow crystal growth processes, and the rate of spontaneous crystallization is usually very slow. However, the addition of a cotton thread can accelerate crystallization [25]. In order to determine the optimal crystallization time, experiments using crystallization times from 0 to 112 h and with or without a pretreated cotton thread were conducted. In all experiments, the initial pH and K2C2O4 concentration were 1.5 and 0.8 mol/L, respectively. As can be seen from Fig. 3, when a pretreated cotton thread was used, the crystallization efficiency of Fe rapidly increased from 0 to 56.81% as the crystallization time was increased from 0 to 12 h. This indicated that the initial stage of crystallization involved rapid nucleation. The pretreated cotton thread increased the specific surface area available for crystallization, and the K3Fe(C2O4)3·3H2O adsorbed on the pretreated cotton thread induced crystallization [26]. At longer crystallization times, the crystallization efficiency of Fe increased more gradually, indicating that the crystallization had entered the slow crystal growth phase due to the decreased iron concentration and decreased solution saturation [27]. After 56 h, the crystallization efficiency of Fe had reached 90.69%; when the crystallization time was extended further, the crystallization efficiency of Fe tended to equilibrium. When no cotton thread was used, all phases of the crystallization process were very slow, mainly because crystals do not form quickly without the inductive effect [26]. Without the cotton thread, 88 h of crystallization were required to reach a Fe crystallization efficiency of 90.06%; that is, the crystallization time with the cotton thread was 32 h shorter than that without the cotton thread. Therefore, the cotton thread was utilized in further crystallization experiments, and 56 h was selected as the optimum crystallization time.
3.1.2. Effect of the K2C2O4 concentration and initial pH on the separation of V and Fe K2C2O4 is the salt of a strong alkali and a weak acid; therefore, aqueous solutions of K2C2O4 are weakly alkaline. In exploratory experiments, we found that the pH value of the solution and the concentration of K2C2O4 both affected the crystallization efficiency of iron. Therefore, experiments with different K2C2O4 concentrations at different initial pH values were carried out using a crystallization time of 60 h and crystallization temperature of 5 °C. As shown in Fig. 2(a), for an initial solution pH of 0.5, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiency of Fe increased from 0 to 56.1%, the crystallization efficiency of V remained lower than 1%, the crystallization efficiency of Al increased from 0 to 2.8%, and the equilibrium pH rose from 0.5 to 1.8. At K2C2O4 concentrations of 0–0.6 mol/L, the crystallization efficiency of Fe was lower than 14%, and the equilibrium pH ranged from 0.5 to 1.25. When the K2C2O4 concentration was increased to 0.8–1.0 mol/L, the crystallization efficiency of Fe increased to 47.5%-57.0%, and the equilibrium pH rose to 1.6–1.8. This indicated that the crystallization efficiency of Fe was low regardless of the K2C2O4 concentration when the equilibrium pH was below 1.8. As can be seen from Fig. 2(b), for an initial pH of 1.0, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiencies of Fe, V, and Al increased from 0 to 87.7%, from 0 to 1.4%, and from 0 to 4.6%, respectively, and the equilibrium pH rose from 1.0 to 3.0. Compared to Fig. 2(a), the crystallization efficiency of Fe at a
3.1.4. Separation of V and Fe As discussed above, 90.7% of the Fe was recovered with V and Al losses of only 1.4% and 5.0% using the optimum conditions (initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization time: 56 h, 3
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40
1.5
30 1.0
20 10
0.5
0 0.0
0.2
0.4
0.6
0.8
1.0
80 70 60 50
0.0
80 70 60
20
1.5
10
1.0
(c)
3.5 3.0
40
2.0
30
1.5
20 10 0.8
0.6
0.8
1.0
1.0
(d)
0.5
90 80 70 60
5.0 4.5 4.0 3.5
50 40
3.0
30 20
2.5
10
1.0
0 0.6
0.4
Fe Fe2+ Fe3+ V Al pH
100
4.5 4.0
0.4
0.2
K2C2O4 concerntration (mol/L)
2.5
0.2
2.0
110
5.0
50
0.0
2.5
30
0.0
Crystallization efficiency (%)
Crystallization efficiency (%)
90
3.0
0
Equilibrium pH
Fe Fe2+ Fe3+ V Al pH
100
3.5
40
K2C2O4 concerntration (mol/L) 110
4.0
2.0
0
0.5
Equilibrium pH
50
2.0
(b)
Fe Fe2+ Fe3+ V Al pH
90 2.5
Crystallization ratio (%)
60
4.5
Equilibrium pH
(a)
Fe Fe2+ Fe3+ V Al pH
70
100
3.0
Equilibrium pH
Crystallization efficiency (%)
80
0.0
K2C2O4 concerntration (mol/L)
0.2
0.4
0.6
0.8
1.0
K2C2O4 concerntration (mol/L)
Fig. 2. Effect of the K2C2O4 concentration and initial pH. (a) pH = 0.5, (b) pH = 1.0, (c) pH = 1.5, (d) pH = 2.0. (Crystallization time: 60 h, crystallization temperature: 5 °C). Table 2 Separation of V and Fe.
Crystallization efficiency of Fe (%)
100 90
CI (mg/L) Fe V 5366 2064
80 70
With cotton thread Without cotton thread
40
R Initial 2.60
Purified 0.25
Separation factor (βFe/V) 875
Note: CI is the concentration of the initial solution, CP is the concentration of the purified solution, and R is the concentration ratio of Fe and V.
60 50
CP (mg/L) Fe V 500 2036
separation factor for Fe and V was 875, indicating that V and Fe were selectively separated. The concentration of Fe in the purified solution decreased significantly after crystallization, which was conducive to the extraction of vanadium via solvent extraction or ion exchange.
30 20
3.2. Recrystallization and product characterization
10 0 0
3.2.1. Recrystallization The crystallized product was characterized using XRD. The results shown in Fig. 4 indicate that the main phase was K3Fe(C2O4)3·3H2O, but impurities such as KH3(C2O4)2·2H2O were also present in small amounts. The purity of the crystallized product was determined to be only 94.3% using ICP-AES analysis. Therefore, further purification of the crystallized product via recrystallization was required. Deionized water was added to the crystallized product in a liquid–solid ratio (L/kg) of 2:1. The mixture was heated to 60 °C, and after the crystals had dissolved completely, the crystal slurry was cooled to room temperature, filtered, and recrystallized at 5 °C. After
8 16 24 32 40 48 56 64 72 80 88 96 104 112
Crystallization time (h) Fig. 3. Effect of the crystallization time. (Initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization temperature: 5 °C).
crystallization temperature: 5 °C; equilibrium pH: 3.0). Table 2 lists the concentrations of V and Fe in the initial solution (before crystallization) and purified solution (after crystallization). The Fe/V concentration ratio decreased from 2.60 to 0.25 after crystallization, and the 4
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1
1
1
1
1 2 1
2 2 1 2 1
5
10
1 1
1
15
1
20
1
1
25
1
1
1
1 1
30
2θ (°)
1
35
1
1 1
40
45
Fig. 4. XRD pattern of the crystallized product.
the recrystallization was complete, the recrystallized product and filtrate were obtained via filtration, and the recrystallized product was washed using 5 °C deionized water and dried at 50 °C. The filtrate from the recrystallization process was analyzed and found to be suitable for reuse in the leaching process based on its remaining level of oxalate and low impurity content. The recrystallization efficiency of Fe was 96.5%. Therefore, the overall recovery of Fe from the vanadium-containing oxalic acid solution in the crystallization process was 87.5%.
Counts
Al
Ca
Mg
Na
23.62
11.29
0.05
53.25
0.025
0.008
0.003
0.012
VO2 + + HC2 O4 = VOC2 O4 + H+
(6)
VO2 + + 2HC2 O4 = VO(C2 O4 )22 + 2H+
(7)
VO+2
(8)
+ HC2 O4 = VO2 C2 O4 + H+
However, the complex anions of can be further reduced to VO (C2O4)22− by oxalate. The reactions can be expressed as in Eqs. (10)–(15) [33]:
H + + VO2 (C2 O4 )
HVO2 (C2 O4 ) +
1
1
H+ + VO2 (C2 O4 )32
1
9000
1 6000
1 1
1 1
1
1 1 1
1 1
1 1 1
1
HVO2 (C2 O4 )22
1
20
25
30
35
VO(C2 O4 ) + H2 O fast
(12) (13)
HVO2 (C2 O4 )22 HVO2 (C2 O4 )32
(11)
+ HC2 O4
VO(C2 O4 )22 + H2 O fast
(14) (15)
Using the above analysis, the speciation of VO2+ in the oxalic acid system was determined using data from the free software Medusa [34], and is shown in Fig. 7(a). V was determined to exist stably as the complex anion VO(C2O4)22− over a wide pH range of 0–6. As shown in Fig. 2, the vanadium loss was low under all conditions, which indicated that the vanadium loss was independent of the initial pH and K2C2O4 concentration, and that the vanadium in the initial solution was not involved in crystallization process of K3Fe(C2O4)3·3H2O.
1
40
HVO2 (C2 O4 ) + HC2 O4
+ HC2 O4
HVO2 (C2 O4 )32 + H+
11 1 1 11
3000 15
H+
(10)
HVO2 C2 O4
HVO2 (C2 O4 ) + HC2 O4
1 1
(9)
VO2+
1
10
C2O42−
VO+2 + 2HC2 O4 = VO2 (C2 O4 )32 + 2H+
1——K3Fe(C2O4)·3H2O
5
Fe(II)
3.3.1. Speciation of V and Fe in a complex oxalic acid solution containing vanadium The speciation of V has been previously investigated by researchers. V exists mainly as VO2+ and VO2+ in inorganic acid solutions with a pH value of less than 1 [31]. However, in oxalic acid solution, vanadium exists as a complex anion due to reduction and complexation with oxalic acid. VO2+ forms the complexes VOC2O4 and VO(C2O4)22−, and VO2+ forms the complexes VO2C2O4− and VO2(C2O4)23− via the reactions given in Eqs. (6)-(9) [32]:
18000
12000
TFe
3.3. Mechanism of the recovery of the iron impurities
3.2.2. Product characterization In order to confirm the recrystallized product, XRD analysis was first used; the results are shown in Fig. 5. All the diffraction peaks in the XRD pattern could be perfectly indexed to the structure of K3Fe (C2O4)3·3H2O from the PDF card 96–152–8614, indicating that K3Fe (C2O4)3·3H2O was of very high purity. Thus, its chemical composition was then analyzed. The data presented in Table 3 clearly show that the contents of impurities in K3Fe(C2O4)3·3H2O were very low, and that it had a purity of 99.0%. Notably, the Fe2+ content was only 0.05%, which indicated that most of the Fe2+ in the initial solution had been oxidized to Fe3+, and thus, a redox reaction may have occurred before or during crystallization. In order to determine its structure, the recrystallized product was
15000
K
characterized using FT-IR. As shown in Fig. 6(a), O–H stretching vibrations corresponding to water molecules were observed at 3567 cm−1 and 3431 cm−1 [28]; these characteristic peaks may have originated from the crystal water of K3Fe(C2O4)3·3H2O. Additionally, the peaks at 1714 cm−1, 1684 cm−1, and 1646 cm−1 were attributed to asymmetrical deformation vibrations νas(C]O) [20,28]; those at 1393 cm−1, 1273 cm−1, 1256 cm−1, and 893 cm−1 corresponded to symmetrical deformation vibrations νs(CeO) [20,28]; and that at 805 cm−1 originated from a bending vibration δ(OCO) [20,28]. The presence of these peaks indicated that the product contained oxalate functional groups. Moreover, the peaks at 582 cm−1, 533 cm−1, and 499 cm−1 were assigned to ν(FeeO) [28–29]. The assignments of the vibrations in FTIR spectra of the K3[Fe(C2O4]3·3H2O are given in Table 4. The above analysis provided further confirmation that the recrystallized product was K3Fe(C2O4)3·3H2O. To understand the spatial arrangement of the atoms, the crystal structure of the K3Fe(C2O4)3·3H2O unit cell was determined, and is shown in Fig. 6(b). Octahedral coordination was observed between the three oxalate ions and iron [30].
1
1 1
1 1
1—K3Fe(C2O4)3·3H2O 2—KH3(C2O4)2·2H2O
1 11
1
Table 3 Chemical composition of K3Fe(C2O4)3·3H2O (wt.%).
45
2θ (°) Fig. 5. XRD pattern of the recrystallized product. 5
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(b)
(a)
582 893
3567
1256 533 499 805 1273
3431
1646 1393
1714
1684 500
1000
1500
2000
2500
Wavenumber (cm-1)
3000
3500
4000
Fig. 6. FTIR spectra and crystal structure of the K3[Fe(C2O4]3·3H2O. Table 4 Assignment of the vibrations in the FTIR spectra of the K3[Fe(C2O4]3·3H2O. Wavenumber (cm−1)
Assignment
References
3567, 3431 1714, 1684, 1646 1393 1273, 1256, 893 805 582, 533, 499
νa(OeH) νa(C]O) νs(CeO) + ν(C]O) νs(CeO) + δ(OeC]O) δ(OeC]O) ν(FeeO)
[28] [20,28] [20,28] [20,28] [20,28] [28–29]
Fe2 + + HC2 O4 = FeC2 O4 + H +
(16)
Fe2 + + 2HC2 O4 = Fe(C2 O4 )22 + 2H+
(17)
Fe3 + + HC2 O24 - = FeC2 O4+ + H+
(18)
Fe3 + + 2HC2 O4 = Fe(C2 O4 )2 + 2H +
(19)
2+
100
(a)
80
60
40
C2O42-
60 50 40 30
VOC2O4 0
1
3
pH
4
5
0
6
Fe(C2O4)22-
FeC2O4
10 2
Fe(C2O4)33-
Fe(C2O4)2-
20
20
0
HC2O4-
H 2 C 2 O4
70
Fractionn (%)
Fraction(%)
(b)
90
VO(C2O4)22-
80
3+
Based on the above analysis, the speciation of Fe and Fe in the oxalic acid system was determined using data from the free software Medusa [34]; the results are shown in Fig. 7(b). The speciation of Fe in the oxalic acid system was very complex. Fe2+ existed mainly as FeC2O4, FeC2O4 + Fe(C2O4)22−, and Fe(C2O4)22− in the pH ranges 0–1.0, 1.0–2.0, and > 2.0, respectively. Fe3+ existed mainly in the form of Fe(C2O4)2− and Fe(C2O4)33− at pH values of 0–1.0 and > 1, respectively. In this study, the initial pH value tested for the initial solution was
Compared to that of V, the speciation of Fe in oxalic acid solution is more complex. Fe has been reported to exist as complex anions in oxalic acid solution. The main complexes of Fe2+ and C2O42− are FeC2O4 and Fe(C2O4)22−, and the complexes of Fe3+ and C2O42− include mainly FeC2O4+, Fe(C2O4)2−, and Fe(C2O4)33− [35–36]. The corresponding reactions are given as Eqs. (16)–(20) [36]:
100
(20)
Fe3 + + 3HC2 O4 = Fe(C2 O4 )33 + 3H+
0
1
2
3
pH
4
(a) 0.03 mol/L VO2+ and 2mol/L C2O42-, (b) 0.05mol/L Fe2+, 0.05mol/L Fe3+ and 2mol/L C2O42Fig. 7. The speciation of VO2+ and Fe in oxalic acid system.
6
5
6
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Z. Liu, et al.
0.45. As shown in Fig. 7(b), Fe2+ may have existed in the form of FeC2O4, while Fe3+ may have existed as Fe(C2O4)2−. When the pH value of the initial solution was increased to 1.5, Fe2+ may have existed as soluble FeC2O4 and Fe(C2O4)22−, while Fe3+ may have existed as Fe (C2O4)33−. At a solution pH of 3.0, Fe2+ and Fe3+ may have existed in the form of Fe(C2O4)22− and Fe(C2O4)33−, respectively. According to the results presented in Fig. 2, the crystallization efficiency of Fe increased with increasing initial pH and K2C2O4 concentration. Under the optimal conditions (initial pH: 1.50, K2C2O4 concentration: 0.8 mol/L, equilibrium pH: 3.0) the iron was almost completely transferred to the complex anions Fe(C2O4)33− and Fe(C2O4)22−.
(3) Coordination Since most of the Fe ions were converted to Fe(C2O4)33− and a large amount of K2C2O4 was added, K2C2O4 and Fe(C2O4)33− coordinated to form K3Fe(C2O4)3. In this work, the concentration of iron was only 5366 mg/L (0.096 mol/L); however, crystallization requires a supersaturated solution. The addition of excess K2C2O4 released large amounts of K+ and C2O42− and created a supersaturated solution for Fe (C2O4)33−, which promoted the formation of K3Fe(C2O4)3·3H2O. The coordination reaction can be expressed as follows:
3K2 C2 O4 ·H2 O + 2Fe(C2 O4 )33 + 3H2 O = 2K3Fe(C2 O4 )3 ·3H2 O + 3C2 O24 (27)
3.3.2. Mechanism of the recovery of the iron impurities According to the above analysis and the combined results of the various experiments, the mechanism of the recovery of the iron impurities was concluded to involve the following steps.
(4) Nucleation
H2 C2 O4
HC2 O4 + H +
(21)
The cotton thread was pretreated with K3Fe(C2O4)3·3H2O. When the pretreated cotton thread was added to the treated solution, which was crystallized at low temperature, K3Fe(C2O4)3·3H2O rapidly nucleated and was adsorbed on the treated cotton thread. As shown in Fig. 3, the crystallization efficiency of iron increased slowly when no cotton thread was added, but when the cotton thread was added to the treated solution, the crystallization efficiency of iron improved to reach 56% at 12 h. This enhancement was due to the increased specific surface area for the crystallization with the addition of cotton thread and the inductive effect of the K3Fe(C2O4)3·3H2O adsorbed on the treated cotton thread, resulting in the rapid crystallization and nucleation of iron during the initial stage of crystallization.
HC2 O4
C2 O24 + H +
(22)
(5) Crystal growth
(1) Complexation As the initial pH of the solution was increased from 0.45 to 1.5, a large number of H+ were neutralized by OH−, and free H2C2O4 molecules were ionized to HC2O4− and C2O42−. Subsequently, the speciation state of the Fe-C2O4 complex anions in the solution changed, with a fraction of the soluble FeC2O4 being complexed to form Fe (C2O4)22−, and Fe(C2O4)2− forming Fe(C2O4)33−. These reactions can be expressed as follows [37]:
FeC2 O4 + C2 O24 = Fe(C2 O4 )22
(23)
Fe(C2 O4 )2 + C2 O24 = Fe(C2 O4 )33
(24)
With increasing crystallization time, the iron concentration in the solution decreased gradually, and the solution changed from a supersaturated state to saturated state and then an unsaturated state. The crystal nuclei of K3Fe(C2O4)3·3H2O that had been formed grew slowly along the core planes of the surface layer of the crystal nuclei until the end of crystallization. Based on the above discussion, a conversion diagram for the Fe(II)Fe(III)-oxalate complexes during the iron recovery process was determined, and is shown in Fig. 8.
(2) Oxidization After the addition of K2C2O4, the solution was stirred in a water bath at 60 °C for 30 min. After this time, the equilibrium pH value of the mixed solution had risen to 3.0, and the speciation of the complex anions in the solution had changed further. The FeC2O4 was completely converted to Fe(C2O4)22−. However, the stability constant (lgβ) of the complex ion Fe(C2O4)22− was only 4.52 [38]; thus, its chemical properties were not stable and it was readily oxidized by the O2 in the air to form the stable complex ion Fe(C2O4)33− (lgβ = 20.2 [38]) under heating [39–40]. When Fe2+ is added to an oxygen-saturated aqueous solution of oxalic acid, the Fe(II)-oxalate complexes are oxidized to Fe (III)-oxalate complexes [39], with the oxidation rate of Fe2+ being determined by the reaction temperature and time [40]. The ratio of Fe2+ to Fe3+ in the solution was measured before and after heating. As shown in Table 5, the percentage of Fe3+ increased from 48.64% to 91.56%, while that of Fe2+ decreased from 51.36% to 8.44% after heating, which indicates that most of the Fe2+ was oxidized to Fe3+ during heating. This result was consistent with previous reports in the literature [39–40]. Therefore, the oxidization reaction can be expressed as follows:
Fe(C2 O4 )22 + O2 = Fe(C2 O4 )2 + 2O2
(25)
Fe(C2 O4 )2 + C2 O24 = Fe(C2 O4 )33
(26)
4. Conclusions (1) Iron was recovered in the form of K3Fe(C2O4)3·3H2O from a complex oxalic acid solution containing vanadium by a crystallization process for the first time. 90.7% of the Fe was recovered and only 1.4% of V and 5.0% of Al were lost under the optimum conditions (initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization time: 56 h, crystallization temperature: 5 °C). The crystalline product was then further purified by recrystallization to obtain the product K3Fe(C2O4)3·3H2O in 99.0% purity. (2) The vanadium in the initial solution did not participate in the crystallization reaction, as it was present as the stable complex VO (C2O4)22− in the oxalic acid system. The recovery mechanism of the iron impurity proceeded via five main steps: (I) Complexation: the complex anions of Fe2+ and Fe3+ formed Fe(C2O4)22− and Fe (C2O4)33−. (II) Oxidization: the unstable complex ion Fe(C2O4)22− was oxidized to the stable complex ion Fe(C2O4)33−. (III) Coordination: K2C2O4 and Fe(C2O4)33− coordinated to form K3[Fe (C2O4)3]. (IV) Nucleation: K3Fe(C2O4)3·3H2O crystals nucleated rapidly. (V) Growth: The formed K3Fe(C2O4)3·3H2O crystals grew slowly. (3) The crystallization process not only separated V and Fe, but also allowed the recovery of Fe as the value-added product K3Fe (C2O4)3·3H2O. Therefore, this process represents a new direction for the comprehensive utilization of vanadium resources.
Table 5 Percentages of Fe2+ and Fe3+ in the solution before and after heating (%).
Before heating After heating
Fe2+
Fe3+
TFe
51.36 8.44
48.64 91.56
100.00 100.00
7
Separation and Purification Technology 232 (2020) 115970
Z. Liu, et al.
FeC2O4
C2O42Fe(C2O4)22(23)
Fe(C2O4)2-
C2O42(24)
O2 (25)
Fe(C2O4)2-
C2O42(26)
Fe(C2O4)33-
K2C2O4 (27)
Fe(C2O4)33-
Acknowledgements This study was financially supported by Project of National Natural Science Foundation of China (No. 51774215 and No. 51774216) and Project of Hubei Province Science Foundation of China (No. 2018CFA068).
[17]
Appendix A. Supplementary material
[20]
Supplementary data to this article can be found online at https:// doi.org/10.1016/j.seppur.2019.115970.
[21]
[18] [19]
[22]
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8
K3Fe(C2O4)3·3H2O
Fig. 8. Conversion diagram for the Fe(II)-Fe (III)-oxalate complexes during the iron recovery process.
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Separation and recovery of iron impurities from a complex oxalic acid solution containing vanadium by K3Fe(C2O4)3·3H2O crystallization
T
Zishuai Liua,b,c,d, Jing Huanga,b,c,d, , Yimin Zhanga,b,c,d,e, , Tao Liua,b,c,d, Pengcheng Hua,b,c,d, Hong Liua,b,c,d, Qiushi Zhenga,b,c,d ⁎
⁎
a
School of Resource and Environmental Engineering, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China State Environmental Protection Key Laboratory of Mineral Metallurgical Resources Utilization and Pollution Control, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China c Hubei Collaborative Innovation Center for High Efficient Utilization of Vanadium Resources, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China d Hubei Provincial Engineering Technology Research Center of High Efficient Cleaning Utilization for Shale Vanadium Resource, Wuhan University of Science and Technology, Wuhan 430081, Hubei Province, China e School of Resource and Environmental Engineering, Wuhan University of Technology, Wuhan 430070, Hubei Province, China b
ARTICLE INFO
ABSTRACT
Keywords: Iron impurities Recovery K3Fe(C2O4)3·3H2O Crystallization Oxalic acid Vanadium
Vanadium and iron have been effectively separated from the acid leachate of vanadium-bearing shale by solvent extraction or ion exchange; however, in these methods, the iron was not recovered, and the comprehensive utilization rates of the resources were low. In this paper, the recovery of iron impurities from a complex oxalic acid solution containing vanadium via the crystallization of K3Fe(C2O4)3·3H2O was investigated. Using the optimum conditions, 90.7% of the Fe was recovered, and only 1.4% of the V and 5.0% of the Al were lost. The concentration of Fe decreased from 5366 mg/L before crystallization to 500 mg/L after crystallization. The crystallized product was then further purified by recrystallization to obtain K3Fe(C2O4)3·3H2O in 99.0% purity, which was characterized using FT-IR, XRD, and SEM. Based on the speciation of vanadium and iron in the oxalic acid system, the mechanism by which the iron impurities were recovered was determined. The mechanism consisted of five main steps: complexation, oxidization, coordination, nucleation, and crystal growth. The recovery of iron impurities from a complex oxalic acid solution containing vanadium via the crystallization of K3Fe (C2O4)3·3H2O is a novel and efficient method, and this process provides a new direction for the comprehensive utilization of vanadium resources.
1. Introduction Vanadium-bearing shale is an important vanadium resource in China [1]. With the increasing market demand for vanadium, the efficient and clean extraction of vanadium from vanadium-bearing shale has become an important task [2]. To achieve this goal, oxalic acid, a clean and eco-friendly organic acid, has been used successfully to extract vanadium from vanadium-bearing shale [3–5]. The leaching of the raw shale using oxalic acid leaching selectively extracts V over Fe impurities, but for roasted shale, the iron is easily dissolved along with the vanadium [5], resulting in a high iron concentration in the leachate. However, iron impurities affect the purification and enrichment of vanadium, for instance, by increasing the amount of extractant required, reducing the purity of the vanadium-rich liquid, and influencing
the quality of the resulting vanadium products. Therefore, proper treatment of the iron in vanadium-bearing shale during the vanadium extraction process is very important. In recent decades, the main methods of treating iron in different vanadium- containing solutions have included solvent extraction, ion exchange, and chemical precipitation [6–7]. Solvent extraction has been most widely used to separate vanadium and iron, and various extractants have been reported, including di-2-ethylhexyl phosphoric acid (D2EHPA) [8], 2-ethylhexyl phosphonic acid mono-2-ethylhexyl ester (PC-88A) [9], and trialkylamine (Alamine 336) [10]. However, solvent extraction methods have presented some drawbacks: (1) The valuable metal iron is not recovered; (2) the extractant cost is high; and (3) the extraction process is long. Ion exchange is a more environmentally friendly method for the separation of vanadium from
⁎ Corresponding authors at: School of Resources and Environmental Engineering, Wuhan University of Science and Technology, Wuhan, 430081, Hubei Province, China. E-mail addresses: [email protected] (J. Huang), [email protected] (Y. Zhang).
https://doi.org/10.1016/j.seppur.2019.115970 Received 10 June 2019; Received in revised form 9 August 2019; Accepted 21 August 2019 Available online 22 August 2019 1383-5866/ © 2019 Elsevier B.V. All rights reserved.
Separation and Purification Technology 232 (2020) 115970
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iron, and a variety of resins have been used, including D201 [11], D314 [12], and AG1-x8 [13]. However, ion exchange also has some disadvantages: (1) Iron and other valuable metals are not recovered; (2) the resin can easily be polluted by the impurity ions; and (3) the adsorption capacity of resins is limited. Chemical precipitation is an effective method for the separation and recovery of iron, and numerous researchers have precipitated iron as FeC2O4 from stone coal [4] and red mud [14]. The disadvantages of chemical precipitation include high vanadium losses and poor selectivity [15]. In general, the separation of vanadium and iron from various vanadium-containing solutions can be achieved using the above methods, but the valuable metal iron is not recycled and the comprehensive utilization rate of the vanadiumbearing shale is low. Based on the above analysis, the current processes of treating iron from vanadium-containing solution could be improved in three ways: (1) Developing a process with simple operation and a short flowsheet; (2) recovering iron as high value-added products; and (3) enhancing the selectivity towards iron and decreasing the loss of vanadium. A variety of complexes have been reported to form between oxalate and iron ions in aqueous solution, including Fe(C2O4)33−, Fe(C2O4)22−, FeC2O4+, and others [16]. However, Fe(C2O4)33− readily decomposes into Fe(C2O4)22− or FeC2O4 when the solutions are exposed to strong sunlight [17–18]. This feature can be exploited through use of potassium tris(oxalato)ferrate(III) trihydrate (K3Fe(C2O4)3·3H2O) as a promoter for the photodegradation of textile dyes in water [19]. K3Fe (C2O4)3·3H2O has been extensively synthesized and researched, and has also been used to fabricate functional metal-organic frameworks and nanostructured materials in recent years [20]. A previous report [21] summarized the three most typical methods used to synthesize K3Fe (C2O4)3·3H2O: (1) The oxidation of FeC2O4 by H2O2 in the presence of H2C2O4 and K2C2O4; (2) the direct combination of FeCl3 and K2C2O4·H2O solutions; and (3) the digestion of a mixture of Fe2(SO4)3 and K2C2O4 over a steam bath. The reactions corresponding to the three synthesis methods are given below:
Table 1 Chemical composition of the initial solution (mg/L).
FeCl3(aq) + 3K2C2 O4 ·H2 O(aq) = K3Fe(C2 O4 )3 ·3H2 O(aq) + 3KCl (aq)
Fe(II)
Al
Mg
Ca
K
Na
2064
5366
2756
20,920
2430
1260
2225
728
the shale with 6 mol/kg H2C2O4 and 5% CaF2 at 95 °C for 4 h [3–5]. The vanadium-bearing shale was supplied by Hubei Tongshan Mining Co., Ltd. The pH of the initial solution was 0.45, and its chemical composition is given in Table 1. The initial solution was a complex system due to its high acidity, various impurities, and the organic acid. Deionized water was used in all experiments, and all reagents were of analytical grade. 2.2. Experimental procedures First, 50 mL of the initial solution was placed in a 100 mL beaker and adjusted to the required pH value using KOH. Next, a given amount of K2C2O4 was added to the resulting solution, which was stirred in a water bath at 60 °C for 30 min. The obtained mixed solution was then held at room temperature for 2 h. After measuring its pH value, the solution was filtered to obtain the treated solution. At this time, a cotton thread 1 mm in diameter that had been pretreated by soaking in a saturated aqueous solution of K3Fe(C2O4)3·3H2O for one day and allowed to dry naturally was added to the treated solution. The solution was then allowed to crystallize at low temperature for the desired time. After the crystallization was completed, the crystalline product and the purified solution were separated using filtration. In order to obtain the product in higher purity, the crystalline product was then recrystallized. Finally, the purified solution was analyzed, and the crystallization efficiency (E) and separation factor were calculated as follows:
E =
(1) (2)
=
Fe2 (SO4 )3(aq) + 3K2C2 O4 ·H2 O(aq) + 3BaC2 O4(s) = 2K3Fe(C2 O4 )3 ·3H2 O(aq) + 3BaSO4(s)
TFe*
* Note: TFe stands for total Fe.
2FeC2 O4(s) + H2 C2 O4 ·2H2 O(aq) + 3K2C2 O4 · H2 O(aq) + H2 O2(aq) = 2K3Fe(C2 O4 )3·3H2 O(aq) + H2 O(aq)
V
C0 V0 C1 V1 × 100% C0 V0
Da Db
(4) (5)
where C0 and C1 are the ion concentration in the initial solution and purified solution (mg/L), respectively, and V0 and V1 are the volume of the initial solution and the purified solution (L), respectively. β is the separation factor of the components a and b, and Da and Db are the ratios of the contents of a and b in the crystal to the amount in the purified solution, respectively. Strong light was avoided during the crystallization processes, and the processes were performed one time.
(3)
According to recent reports [3], the concentration of iron in the oxalic acid leachate of shale is 2–3 times that of vanadium, and the iron is present as Fe2+ and Fe3+ ions. Therefore, this iron could be recovered as K3Fe(C2O4)3·3H2O from vanadium-containing oxalic acid solutions. Unfortunately, until now, the recovery of iron via its crystallization as K3Fe(C2O4)3·3H2O from complex, highly acidic solutions containing various ions has not been reported. Therefore, in this paper, a novel and efficient process for the recovery of iron impurities from a complex oxalic acid solution containing vanadium was studied, with the aim of recovering the iron in the form of K3Fe(C2O4)3·3H2O via a crystallization process. The effects of different crystallizing agents and their concentration, the initial pH, and the crystallization time on the crystallization efficiency of iron were investigated. Moreover, the product K3Fe(C2O4)3·3H2O was characterized, and the mechanism of iron recovery was determined.
2.3. Analysis methods Solution pH values were measured using a pH meter (PHS-3C, INESA Scientific Instrument Co., Ltd). The concentrations of all elements were analyzed using inductively coupled plasma-atomic emission spectroscopy (ICP-AES, Optima-4300DV, PerkinElmer, Boston, MA, USA) in addition to ferrous iron (Fe (II)), ferric iron (Fe(III)), and oxalate concentration. The Fe(II) concentration was measured using the 1, 10-phenanthroline spectrophotometry using a spectrophotometer (UV5500, Metash, Shanghai, China) [22–23], and the Fe(III) concentration was calculated from the mass balance. The oxalate concentration of K3Fe(C2O4)3·3H2O was determined by titration with potassium permanganate [24]. The functional groups of the product were analyzed using Fourier transform infrared spectroscopy (FTIR, Thermo Nicolet NEXUS, USA). The phase compositions were determined using X-ray diffraction (XRD, D/MAX 2500PC, Rigaku, Tokyo, Japan).
2. Experimental 2.1. Materials A complex oxalic acid solution containing vanadium (which will be referred to from here on as the initial solution) was prepared by roasting vanadium-bearing shale at 850 °C for 1 h, followed by leaching 2
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given K2C2O4 concentration increased, but even at a K2C2O4 concentration of 1.0 mol/L, the crystallization efficiency of Fe was still not close to equilibrium. As shown in Fig. 2(c), when the initial pH of the solution was 1.5, the crystallization efficiencies of Fe, V, and Al increased from 0 to 91.8%, from 0 to 2.3%, and from 0 to 6.3%, respectively, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, and the equilibrium pH rose from 1.5 to 4.1. At a K2C2O4 concentration of 0.8 mol/L, the crystallization efficiency of Fe was 90.4%. Further increasing the K2C2O4 concentration increased the crystallization efficiency of Fe only slightly and increased the losses of V and Al. Fig. 2(d) shows that when the initial pH of the solution was 2.0, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiencies of Fe, V, and Al increased from 0 to 92.9%, from 0 to 3.2%, and from 0 to 7.8%, respectively, and the equilibrium pH rose from 2.0 to 4.8. At a K2C2O4 concentration of 0.6 mol/L, the equilibrium pH was 3.4, but the crystallization efficiency of Fe was only 83.3%. However, when the K2C2O4 concentration exceeded 0.8 mol/L, the crystallization efficiency of Fe was > 91%, but the losses of V and Al showed obvious increases. Based on the above data, the crystallization efficiency of Fe was determined by the equilibrium pH and K2C2O4 concentration simultaneously. Moreover, a comparison of the four sets of results in Fig. 2 shows that the crystallization efficiency of Fe3+ was always higher than that of Fe2+, which indicated that Fe3+ crystallized more readily than Fe2+. In addition, as the equilibrium pH increased, the crystallization efficiency of V and Al also increased. Based on overall considerations, an initial pH of 1.5 and a K2C2O4 concentration of 0.8 mol/L were selected as the optimal conditions; the corresponding equilibrium pH was 3.0.
Crystallization efficiency of Fe (%)
100 90
85.1
80 70 60 50 40 30 20 10 0
8.5
5.6
6.8
7.3
KCl
K2SO4
KOH
KNO3
K2 C 2 O4
Fig. 1. Effect of the crystallizing agents. (K+: 1.6 mol/L, equilibrium pH value: 3.0, crystallization time: 48 h, crystallization temperature: 5 °C).
3. Results and discussion 3.1. Crystallization of K3Fe(C2O4)3·3H2O 3.1.1. Effect of the crystallizing agents on the separation of V and Fe The effect of different crystallizing agents was studied using a fixed K+ concentration of 1.6 mol/L, an equilibrium pH value of 3.0, a crystallization time of 48 h, and crystallization temperature of 5 °C. Fig. 1 clearly shows that the crystallization efficiency of Fe was the highest (85.1%) when K2C2O4 was used as the crystallizing agent, while that of the other crystallizing agents were all lower than 10%, indicating that K2C2O4 was more likely to form complexes with Fe3+. Piro et al. [21] summarized the three most commonly used synthetic methods for K3Fe(C2O4)3·3H2O, and K2C2O4 was required for all three. This further indicates that K2C2O4 easily complexes with Fe3+ to form K3Fe(C2O4)3·3H2O. Thus, K2C2O4 was selected as the optimum crystallizing agent.
3.1.3. Effect of the crystallization time on the separation of V and Fe Crystallization at low temperatures results in slow crystal growth processes, and the rate of spontaneous crystallization is usually very slow. However, the addition of a cotton thread can accelerate crystallization [25]. In order to determine the optimal crystallization time, experiments using crystallization times from 0 to 112 h and with or without a pretreated cotton thread were conducted. In all experiments, the initial pH and K2C2O4 concentration were 1.5 and 0.8 mol/L, respectively. As can be seen from Fig. 3, when a pretreated cotton thread was used, the crystallization efficiency of Fe rapidly increased from 0 to 56.81% as the crystallization time was increased from 0 to 12 h. This indicated that the initial stage of crystallization involved rapid nucleation. The pretreated cotton thread increased the specific surface area available for crystallization, and the K3Fe(C2O4)3·3H2O adsorbed on the pretreated cotton thread induced crystallization [26]. At longer crystallization times, the crystallization efficiency of Fe increased more gradually, indicating that the crystallization had entered the slow crystal growth phase due to the decreased iron concentration and decreased solution saturation [27]. After 56 h, the crystallization efficiency of Fe had reached 90.69%; when the crystallization time was extended further, the crystallization efficiency of Fe tended to equilibrium. When no cotton thread was used, all phases of the crystallization process were very slow, mainly because crystals do not form quickly without the inductive effect [26]. Without the cotton thread, 88 h of crystallization were required to reach a Fe crystallization efficiency of 90.06%; that is, the crystallization time with the cotton thread was 32 h shorter than that without the cotton thread. Therefore, the cotton thread was utilized in further crystallization experiments, and 56 h was selected as the optimum crystallization time.
3.1.2. Effect of the K2C2O4 concentration and initial pH on the separation of V and Fe K2C2O4 is the salt of a strong alkali and a weak acid; therefore, aqueous solutions of K2C2O4 are weakly alkaline. In exploratory experiments, we found that the pH value of the solution and the concentration of K2C2O4 both affected the crystallization efficiency of iron. Therefore, experiments with different K2C2O4 concentrations at different initial pH values were carried out using a crystallization time of 60 h and crystallization temperature of 5 °C. As shown in Fig. 2(a), for an initial solution pH of 0.5, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiency of Fe increased from 0 to 56.1%, the crystallization efficiency of V remained lower than 1%, the crystallization efficiency of Al increased from 0 to 2.8%, and the equilibrium pH rose from 0.5 to 1.8. At K2C2O4 concentrations of 0–0.6 mol/L, the crystallization efficiency of Fe was lower than 14%, and the equilibrium pH ranged from 0.5 to 1.25. When the K2C2O4 concentration was increased to 0.8–1.0 mol/L, the crystallization efficiency of Fe increased to 47.5%-57.0%, and the equilibrium pH rose to 1.6–1.8. This indicated that the crystallization efficiency of Fe was low regardless of the K2C2O4 concentration when the equilibrium pH was below 1.8. As can be seen from Fig. 2(b), for an initial pH of 1.0, as the K2C2O4 concentration was increased from 0 to 1.0 mol/L, the crystallization efficiencies of Fe, V, and Al increased from 0 to 87.7%, from 0 to 1.4%, and from 0 to 4.6%, respectively, and the equilibrium pH rose from 1.0 to 3.0. Compared to Fig. 2(a), the crystallization efficiency of Fe at a
3.1.4. Separation of V and Fe As discussed above, 90.7% of the Fe was recovered with V and Al losses of only 1.4% and 5.0% using the optimum conditions (initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization time: 56 h, 3
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40
1.5
30 1.0
20 10
0.5
0 0.0
0.2
0.4
0.6
0.8
1.0
80 70 60 50
0.0
80 70 60
20
1.5
10
1.0
(c)
3.5 3.0
40
2.0
30
1.5
20 10 0.8
0.6
0.8
1.0
1.0
(d)
0.5
90 80 70 60
5.0 4.5 4.0 3.5
50 40
3.0
30 20
2.5
10
1.0
0 0.6
0.4
Fe Fe2+ Fe3+ V Al pH
100
4.5 4.0
0.4
0.2
K2C2O4 concerntration (mol/L)
2.5
0.2
2.0
110
5.0
50
0.0
2.5
30
0.0
Crystallization efficiency (%)
Crystallization efficiency (%)
90
3.0
0
Equilibrium pH
Fe Fe2+ Fe3+ V Al pH
100
3.5
40
K2C2O4 concerntration (mol/L) 110
4.0
2.0
0
0.5
Equilibrium pH
50
2.0
(b)
Fe Fe2+ Fe3+ V Al pH
90 2.5
Crystallization ratio (%)
60
4.5
Equilibrium pH
(a)
Fe Fe2+ Fe3+ V Al pH
70
100
3.0
Equilibrium pH
Crystallization efficiency (%)
80
0.0
K2C2O4 concerntration (mol/L)
0.2
0.4
0.6
0.8
1.0
K2C2O4 concerntration (mol/L)
Fig. 2. Effect of the K2C2O4 concentration and initial pH. (a) pH = 0.5, (b) pH = 1.0, (c) pH = 1.5, (d) pH = 2.0. (Crystallization time: 60 h, crystallization temperature: 5 °C). Table 2 Separation of V and Fe.
Crystallization efficiency of Fe (%)
100 90
CI (mg/L) Fe V 5366 2064
80 70
With cotton thread Without cotton thread
40
R Initial 2.60
Purified 0.25
Separation factor (βFe/V) 875
Note: CI is the concentration of the initial solution, CP is the concentration of the purified solution, and R is the concentration ratio of Fe and V.
60 50
CP (mg/L) Fe V 500 2036
separation factor for Fe and V was 875, indicating that V and Fe were selectively separated. The concentration of Fe in the purified solution decreased significantly after crystallization, which was conducive to the extraction of vanadium via solvent extraction or ion exchange.
30 20
3.2. Recrystallization and product characterization
10 0 0
3.2.1. Recrystallization The crystallized product was characterized using XRD. The results shown in Fig. 4 indicate that the main phase was K3Fe(C2O4)3·3H2O, but impurities such as KH3(C2O4)2·2H2O were also present in small amounts. The purity of the crystallized product was determined to be only 94.3% using ICP-AES analysis. Therefore, further purification of the crystallized product via recrystallization was required. Deionized water was added to the crystallized product in a liquid–solid ratio (L/kg) of 2:1. The mixture was heated to 60 °C, and after the crystals had dissolved completely, the crystal slurry was cooled to room temperature, filtered, and recrystallized at 5 °C. After
8 16 24 32 40 48 56 64 72 80 88 96 104 112
Crystallization time (h) Fig. 3. Effect of the crystallization time. (Initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization temperature: 5 °C).
crystallization temperature: 5 °C; equilibrium pH: 3.0). Table 2 lists the concentrations of V and Fe in the initial solution (before crystallization) and purified solution (after crystallization). The Fe/V concentration ratio decreased from 2.60 to 0.25 after crystallization, and the 4
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1
1
1
1
1 2 1
2 2 1 2 1
5
10
1 1
1
15
1
20
1
1
25
1
1
1
1 1
30
2θ (°)
1
35
1
1 1
40
45
Fig. 4. XRD pattern of the crystallized product.
the recrystallization was complete, the recrystallized product and filtrate were obtained via filtration, and the recrystallized product was washed using 5 °C deionized water and dried at 50 °C. The filtrate from the recrystallization process was analyzed and found to be suitable for reuse in the leaching process based on its remaining level of oxalate and low impurity content. The recrystallization efficiency of Fe was 96.5%. Therefore, the overall recovery of Fe from the vanadium-containing oxalic acid solution in the crystallization process was 87.5%.
Counts
Al
Ca
Mg
Na
23.62
11.29
0.05
53.25
0.025
0.008
0.003
0.012
VO2 + + HC2 O4 = VOC2 O4 + H+
(6)
VO2 + + 2HC2 O4 = VO(C2 O4 )22 + 2H+
(7)
VO+2
(8)
+ HC2 O4 = VO2 C2 O4 + H+
However, the complex anions of can be further reduced to VO (C2O4)22− by oxalate. The reactions can be expressed as in Eqs. (10)–(15) [33]:
H + + VO2 (C2 O4 )
HVO2 (C2 O4 ) +
1
1
H+ + VO2 (C2 O4 )32
1
9000
1 6000
1 1
1 1
1
1 1 1
1 1
1 1 1
1
HVO2 (C2 O4 )22
1
20
25
30
35
VO(C2 O4 ) + H2 O fast
(12) (13)
HVO2 (C2 O4 )22 HVO2 (C2 O4 )32
(11)
+ HC2 O4
VO(C2 O4 )22 + H2 O fast
(14) (15)
Using the above analysis, the speciation of VO2+ in the oxalic acid system was determined using data from the free software Medusa [34], and is shown in Fig. 7(a). V was determined to exist stably as the complex anion VO(C2O4)22− over a wide pH range of 0–6. As shown in Fig. 2, the vanadium loss was low under all conditions, which indicated that the vanadium loss was independent of the initial pH and K2C2O4 concentration, and that the vanadium in the initial solution was not involved in crystallization process of K3Fe(C2O4)3·3H2O.
1
40
HVO2 (C2 O4 ) + HC2 O4
+ HC2 O4
HVO2 (C2 O4 )32 + H+
11 1 1 11
3000 15
H+
(10)
HVO2 C2 O4
HVO2 (C2 O4 ) + HC2 O4
1 1
(9)
VO2+
1
10
C2O42−
VO+2 + 2HC2 O4 = VO2 (C2 O4 )32 + 2H+
1——K3Fe(C2O4)·3H2O
5
Fe(II)
3.3.1. Speciation of V and Fe in a complex oxalic acid solution containing vanadium The speciation of V has been previously investigated by researchers. V exists mainly as VO2+ and VO2+ in inorganic acid solutions with a pH value of less than 1 [31]. However, in oxalic acid solution, vanadium exists as a complex anion due to reduction and complexation with oxalic acid. VO2+ forms the complexes VOC2O4 and VO(C2O4)22−, and VO2+ forms the complexes VO2C2O4− and VO2(C2O4)23− via the reactions given in Eqs. (6)-(9) [32]:
18000
12000
TFe
3.3. Mechanism of the recovery of the iron impurities
3.2.2. Product characterization In order to confirm the recrystallized product, XRD analysis was first used; the results are shown in Fig. 5. All the diffraction peaks in the XRD pattern could be perfectly indexed to the structure of K3Fe (C2O4)3·3H2O from the PDF card 96–152–8614, indicating that K3Fe (C2O4)3·3H2O was of very high purity. Thus, its chemical composition was then analyzed. The data presented in Table 3 clearly show that the contents of impurities in K3Fe(C2O4)3·3H2O were very low, and that it had a purity of 99.0%. Notably, the Fe2+ content was only 0.05%, which indicated that most of the Fe2+ in the initial solution had been oxidized to Fe3+, and thus, a redox reaction may have occurred before or during crystallization. In order to determine its structure, the recrystallized product was
15000
K
characterized using FT-IR. As shown in Fig. 6(a), O–H stretching vibrations corresponding to water molecules were observed at 3567 cm−1 and 3431 cm−1 [28]; these characteristic peaks may have originated from the crystal water of K3Fe(C2O4)3·3H2O. Additionally, the peaks at 1714 cm−1, 1684 cm−1, and 1646 cm−1 were attributed to asymmetrical deformation vibrations νas(C]O) [20,28]; those at 1393 cm−1, 1273 cm−1, 1256 cm−1, and 893 cm−1 corresponded to symmetrical deformation vibrations νs(CeO) [20,28]; and that at 805 cm−1 originated from a bending vibration δ(OCO) [20,28]. The presence of these peaks indicated that the product contained oxalate functional groups. Moreover, the peaks at 582 cm−1, 533 cm−1, and 499 cm−1 were assigned to ν(FeeO) [28–29]. The assignments of the vibrations in FTIR spectra of the K3[Fe(C2O4]3·3H2O are given in Table 4. The above analysis provided further confirmation that the recrystallized product was K3Fe(C2O4)3·3H2O. To understand the spatial arrangement of the atoms, the crystal structure of the K3Fe(C2O4)3·3H2O unit cell was determined, and is shown in Fig. 6(b). Octahedral coordination was observed between the three oxalate ions and iron [30].
1
1 1
1 1
1—K3Fe(C2O4)3·3H2O 2—KH3(C2O4)2·2H2O
1 11
1
Table 3 Chemical composition of K3Fe(C2O4)3·3H2O (wt.%).
45
2θ (°) Fig. 5. XRD pattern of the recrystallized product. 5
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(b)
(a)
582 893
3567
1256 533 499 805 1273
3431
1646 1393
1714
1684 500
1000
1500
2000
2500
Wavenumber (cm-1)
3000
3500
4000
Fig. 6. FTIR spectra and crystal structure of the K3[Fe(C2O4]3·3H2O. Table 4 Assignment of the vibrations in the FTIR spectra of the K3[Fe(C2O4]3·3H2O. Wavenumber (cm−1)
Assignment
References
3567, 3431 1714, 1684, 1646 1393 1273, 1256, 893 805 582, 533, 499
νa(OeH) νa(C]O) νs(CeO) + ν(C]O) νs(CeO) + δ(OeC]O) δ(OeC]O) ν(FeeO)
[28] [20,28] [20,28] [20,28] [20,28] [28–29]
Fe2 + + HC2 O4 = FeC2 O4 + H +
(16)
Fe2 + + 2HC2 O4 = Fe(C2 O4 )22 + 2H+
(17)
Fe3 + + HC2 O24 - = FeC2 O4+ + H+
(18)
Fe3 + + 2HC2 O4 = Fe(C2 O4 )2 + 2H +
(19)
2+
100
(a)
80
60
40
C2O42-
60 50 40 30
VOC2O4 0
1
3
pH
4
5
0
6
Fe(C2O4)22-
FeC2O4
10 2
Fe(C2O4)33-
Fe(C2O4)2-
20
20
0
HC2O4-
H 2 C 2 O4
70
Fractionn (%)
Fraction(%)
(b)
90
VO(C2O4)22-
80
3+
Based on the above analysis, the speciation of Fe and Fe in the oxalic acid system was determined using data from the free software Medusa [34]; the results are shown in Fig. 7(b). The speciation of Fe in the oxalic acid system was very complex. Fe2+ existed mainly as FeC2O4, FeC2O4 + Fe(C2O4)22−, and Fe(C2O4)22− in the pH ranges 0–1.0, 1.0–2.0, and > 2.0, respectively. Fe3+ existed mainly in the form of Fe(C2O4)2− and Fe(C2O4)33− at pH values of 0–1.0 and > 1, respectively. In this study, the initial pH value tested for the initial solution was
Compared to that of V, the speciation of Fe in oxalic acid solution is more complex. Fe has been reported to exist as complex anions in oxalic acid solution. The main complexes of Fe2+ and C2O42− are FeC2O4 and Fe(C2O4)22−, and the complexes of Fe3+ and C2O42− include mainly FeC2O4+, Fe(C2O4)2−, and Fe(C2O4)33− [35–36]. The corresponding reactions are given as Eqs. (16)–(20) [36]:
100
(20)
Fe3 + + 3HC2 O4 = Fe(C2 O4 )33 + 3H+
0
1
2
3
pH
4
(a) 0.03 mol/L VO2+ and 2mol/L C2O42-, (b) 0.05mol/L Fe2+, 0.05mol/L Fe3+ and 2mol/L C2O42Fig. 7. The speciation of VO2+ and Fe in oxalic acid system.
6
5
6
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0.45. As shown in Fig. 7(b), Fe2+ may have existed in the form of FeC2O4, while Fe3+ may have existed as Fe(C2O4)2−. When the pH value of the initial solution was increased to 1.5, Fe2+ may have existed as soluble FeC2O4 and Fe(C2O4)22−, while Fe3+ may have existed as Fe (C2O4)33−. At a solution pH of 3.0, Fe2+ and Fe3+ may have existed in the form of Fe(C2O4)22− and Fe(C2O4)33−, respectively. According to the results presented in Fig. 2, the crystallization efficiency of Fe increased with increasing initial pH and K2C2O4 concentration. Under the optimal conditions (initial pH: 1.50, K2C2O4 concentration: 0.8 mol/L, equilibrium pH: 3.0) the iron was almost completely transferred to the complex anions Fe(C2O4)33− and Fe(C2O4)22−.
(3) Coordination Since most of the Fe ions were converted to Fe(C2O4)33− and a large amount of K2C2O4 was added, K2C2O4 and Fe(C2O4)33− coordinated to form K3Fe(C2O4)3. In this work, the concentration of iron was only 5366 mg/L (0.096 mol/L); however, crystallization requires a supersaturated solution. The addition of excess K2C2O4 released large amounts of K+ and C2O42− and created a supersaturated solution for Fe (C2O4)33−, which promoted the formation of K3Fe(C2O4)3·3H2O. The coordination reaction can be expressed as follows:
3K2 C2 O4 ·H2 O + 2Fe(C2 O4 )33 + 3H2 O = 2K3Fe(C2 O4 )3 ·3H2 O + 3C2 O24 (27)
3.3.2. Mechanism of the recovery of the iron impurities According to the above analysis and the combined results of the various experiments, the mechanism of the recovery of the iron impurities was concluded to involve the following steps.
(4) Nucleation
H2 C2 O4
HC2 O4 + H +
(21)
The cotton thread was pretreated with K3Fe(C2O4)3·3H2O. When the pretreated cotton thread was added to the treated solution, which was crystallized at low temperature, K3Fe(C2O4)3·3H2O rapidly nucleated and was adsorbed on the treated cotton thread. As shown in Fig. 3, the crystallization efficiency of iron increased slowly when no cotton thread was added, but when the cotton thread was added to the treated solution, the crystallization efficiency of iron improved to reach 56% at 12 h. This enhancement was due to the increased specific surface area for the crystallization with the addition of cotton thread and the inductive effect of the K3Fe(C2O4)3·3H2O adsorbed on the treated cotton thread, resulting in the rapid crystallization and nucleation of iron during the initial stage of crystallization.
HC2 O4
C2 O24 + H +
(22)
(5) Crystal growth
(1) Complexation As the initial pH of the solution was increased from 0.45 to 1.5, a large number of H+ were neutralized by OH−, and free H2C2O4 molecules were ionized to HC2O4− and C2O42−. Subsequently, the speciation state of the Fe-C2O4 complex anions in the solution changed, with a fraction of the soluble FeC2O4 being complexed to form Fe (C2O4)22−, and Fe(C2O4)2− forming Fe(C2O4)33−. These reactions can be expressed as follows [37]:
FeC2 O4 + C2 O24 = Fe(C2 O4 )22
(23)
Fe(C2 O4 )2 + C2 O24 = Fe(C2 O4 )33
(24)
With increasing crystallization time, the iron concentration in the solution decreased gradually, and the solution changed from a supersaturated state to saturated state and then an unsaturated state. The crystal nuclei of K3Fe(C2O4)3·3H2O that had been formed grew slowly along the core planes of the surface layer of the crystal nuclei until the end of crystallization. Based on the above discussion, a conversion diagram for the Fe(II)Fe(III)-oxalate complexes during the iron recovery process was determined, and is shown in Fig. 8.
(2) Oxidization After the addition of K2C2O4, the solution was stirred in a water bath at 60 °C for 30 min. After this time, the equilibrium pH value of the mixed solution had risen to 3.0, and the speciation of the complex anions in the solution had changed further. The FeC2O4 was completely converted to Fe(C2O4)22−. However, the stability constant (lgβ) of the complex ion Fe(C2O4)22− was only 4.52 [38]; thus, its chemical properties were not stable and it was readily oxidized by the O2 in the air to form the stable complex ion Fe(C2O4)33− (lgβ = 20.2 [38]) under heating [39–40]. When Fe2+ is added to an oxygen-saturated aqueous solution of oxalic acid, the Fe(II)-oxalate complexes are oxidized to Fe (III)-oxalate complexes [39], with the oxidation rate of Fe2+ being determined by the reaction temperature and time [40]. The ratio of Fe2+ to Fe3+ in the solution was measured before and after heating. As shown in Table 5, the percentage of Fe3+ increased from 48.64% to 91.56%, while that of Fe2+ decreased from 51.36% to 8.44% after heating, which indicates that most of the Fe2+ was oxidized to Fe3+ during heating. This result was consistent with previous reports in the literature [39–40]. Therefore, the oxidization reaction can be expressed as follows:
Fe(C2 O4 )22 + O2 = Fe(C2 O4 )2 + 2O2
(25)
Fe(C2 O4 )2 + C2 O24 = Fe(C2 O4 )33
(26)
4. Conclusions (1) Iron was recovered in the form of K3Fe(C2O4)3·3H2O from a complex oxalic acid solution containing vanadium by a crystallization process for the first time. 90.7% of the Fe was recovered and only 1.4% of V and 5.0% of Al were lost under the optimum conditions (initial pH: 1.5, K2C2O4 concentration: 0.8 mol/L, crystallization time: 56 h, crystallization temperature: 5 °C). The crystalline product was then further purified by recrystallization to obtain the product K3Fe(C2O4)3·3H2O in 99.0% purity. (2) The vanadium in the initial solution did not participate in the crystallization reaction, as it was present as the stable complex VO (C2O4)22− in the oxalic acid system. The recovery mechanism of the iron impurity proceeded via five main steps: (I) Complexation: the complex anions of Fe2+ and Fe3+ formed Fe(C2O4)22− and Fe (C2O4)33−. (II) Oxidization: the unstable complex ion Fe(C2O4)22− was oxidized to the stable complex ion Fe(C2O4)33−. (III) Coordination: K2C2O4 and Fe(C2O4)33− coordinated to form K3[Fe (C2O4)3]. (IV) Nucleation: K3Fe(C2O4)3·3H2O crystals nucleated rapidly. (V) Growth: The formed K3Fe(C2O4)3·3H2O crystals grew slowly. (3) The crystallization process not only separated V and Fe, but also allowed the recovery of Fe as the value-added product K3Fe (C2O4)3·3H2O. Therefore, this process represents a new direction for the comprehensive utilization of vanadium resources.
Table 5 Percentages of Fe2+ and Fe3+ in the solution before and after heating (%).
Before heating After heating
Fe2+
Fe3+
TFe
51.36 8.44
48.64 91.56
100.00 100.00
7
Separation and Purification Technology 232 (2020) 115970
Z. Liu, et al.
FeC2O4
C2O42Fe(C2O4)22(23)
Fe(C2O4)2-
C2O42(24)
O2 (25)
Fe(C2O4)2-
C2O42(26)
Fe(C2O4)33-
K2C2O4 (27)
Fe(C2O4)33-
Acknowledgements This study was financially supported by Project of National Natural Science Foundation of China (No. 51774215 and No. 51774216) and Project of Hubei Province Science Foundation of China (No. 2018CFA068).
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Appendix A. Supplementary material
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Supplementary data to this article can be found online at https:// doi.org/10.1016/j.seppur.2019.115970.
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K3Fe(C2O4)3·3H2O
Fig. 8. Conversion diagram for the Fe(II)-Fe (III)-oxalate complexes during the iron recovery process.
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