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Mining, Metallurgy & Exploration

, Volume 36, Issue 1, pp 215–225 | Cite as

Effect of Anions on the Solubility of Rare Earth Element-Bearing Minerals in Acids

  • Kenneth N. HanEmail author
Article
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Abstract

The effect of anions such as sulfate, chloride, and nitrate on the solubility of rare earth element-bearing phosphates and fluoro-carbonates has been examined using relevant thermodynamic data. The study has found that these anions have significant influences on the overall dissolution of rare earth elements (REEs) from various sources in the aqueous media. The thermodynamic calculations in this study have indicated that the speciation of REEs with the anions tends to increase the overall solubility of the REE-bearing minerals. However, in some cases, strong precipitation of REEs with some of these anions causes an adverse effect on the solubility of these elements. It has been found that sulfate ion has the most pronounced effect on the solubility compared to nitrate and chloride, in which the latter two acids exhibited almost identical results. The calculated results have indicated that REE-oxide (hydrated) is the easiest compound to dissolve, followed by carbonate, phosphate, fluoride, sulfate, fluoro-carbonate, and oxalate in that order.

Keywords

Speciation Rare earth elements Cut-off-pH Solubility Sulfate precipitation Leaching 

1 Introduction

Rare earth elements (REEs) have become an important group of metals used in various high-tech industries. The use of REEs is vast including high-strength magnets, plasma TVs, various military applications, and clean and efficient green energy industries.

In general, there are two types of ore deposits containing REEs, one being the “bastnaesite-type” as in the case of the Mountain Pass deposit in California and another, the “monazite/xenotime type,” which frequently occurs in beach sands around the world. Bastnaesite contains primarily cerium and lanthanum, forming a fluoro-carbonate. On the other hand, monazite and xenotime are REE-phosphates, where monazite contains mainly cerium, lanthanum, neodymium, thorium, and yttrium, and xenotime comprises mainly of yttrium as the major constituent along with some heavy REEs as minor ones.

Leaching of REE-bearing ores is usually carried out using acids such as sulfuric, hydrochloric, and nitric. There are a few instances where the direct acid leaching without further treatments has been attempted [1, 2]. However, leaching of most REE-bearing minerals under normal conditions is difficult especially when these minerals are bastnaesite, monazite, or xenotime type. As a result, these ores are treated, in general, by sulfuric acid cracking or alkaline roasting at high temperatures prior to water or acid leaching [3, 4, 5, 6].

However, thermodynamic calculations often indicate that these minerals are soluble in strong acids. The question is what acidity concentration is required to dissolve these minerals. It is also questionable whether different acids would yield different degrees of dissolution of these minerals. Furthermore, it is unclear as to the roles of anions such as sulfate, chloride, and nitrate in the overall solubility of REE-bearing minerals.

In this paper, the solubility of REE-phosphates and fluoro-carbonates in three acids, i.e., sulfuric, hydrochloric, and nitric, has been examined in terms of thermodynamic principles with a particular emphasis on the effect of speciation of REEs with these anions on the overall solubility.

There are more than 250 rare earth-bearing minerals, from which REEs can be extracted [7, 8]. However, most of these minerals are scarce in quantities, and furthermore, the content of REEs is very low to be treated as commercial applications. More than 80% of REEs extracted commercially are extracted from either REE-phosphates (monazite or xenotime) or REE fluoro-carbonates (bastnaesite).

Leaching of these minerals is usually carried out using acids such as sulfuric, hydrochloric, and nitric. Unfortunately, REE phosphates and REE-bearing fluoro-carbonates are in general sparingly soluble in acids under normal conditions, since these compounds are thermodynamically very stable as solids.

2 Thermodynamic Data

Some salient thermodynamic data for various REE compounds have been tabulated in the literature [9, 10]. The Gibbs standard free energy formation values for various REE compounds is shown in Fig. 1 [11, 12, 13, 14, 15, 16, 17, 18, 19, 20]. It should be noted that the majority of the thermodynamic data used for the calculations of speciation of REE were taken from either the HSC Database [11], the NBS Table by Wagman [21] or Dean [22].
Fig. 1

Gibbs standard free energy formation values in KJ/mol for various REE compounds [Oct-hyd-S: Octa-hydrated Sulfate]

It is noted from Fig. 1 that there is a general trend in the Gibbs standard free energy formation of the REE compounds, because the chemical properties of these compounds are very similar. The Gibbs standard free energy formation values of bastnaesite were obtained from the literature value of bastnaesite-Ce [18] and by drawing the line parallel to that of hydrated oxides, which will be termed as the “parallel line concept.” The Gibbs standard free energy formation values for yttrium-compounds are consistent with the general understanding of the chemistry of REEs, which states that the chemistry of yttrium lies between holmium and erbium.

Using the same line of approach described above as the parallel line concept, some of the thermodynamic data that are absent from the literature have been estimated by inter or extrapolation of these lines and the results are shown in Table 1 and Fig. 1. It should be noted that the data estimated in this study are listed in italics in the table.
Table 1

Gibbs free energy formation values in KJ/mol for various REE compounds

 

Oxide (hydrate)

Carbonate

Phosphate

Sulfate

Fluoride

Oxalate

F-Carbonate

Oct-hyd-S*

La

− 1909

− 3215

− 1904

− 3734

− 1672

− 3649

− 1722

− 7207

Ce

− 1896

− 3185

− 1904

− 3701

− 1660

− 3632

− 1709

− 7152

Pr

− 1880

− 3198

− 1886

− 3679

− 1649

− 3600

− 1698

− 7105

Nd

− 1876

− 3186

− 1877

− 3670

− 1638

− 3605

− 1686

− 7056

Sm

− 1871

− 3178

− 1872

− 3676

− 1641

− 3598

− 1690

− 7072

Eu

− 1870

− 3177

− 1869

− 3662

− 1636

− 3594

− 1682

− 7039

Gd

− 1865

− 3174

− 1864

− 3657

− 1633

− 3588

− 1676

− 7014

Tb

− 1864

− 3167

− 1860

− 3655

− 1626

− 3585

− 1684

− 7047

Dy

− 1844

− 3140

− 1850

− 3635

− 1614

− 3581

− 1650

− 6904

Ho

− 1833

− 3113

− 1842

− 3603

− 1602

− 3554

− 1641

− 6865

Er

− 1822

− 3112

− 1837

− 3581

− 1593

− 3522

− 1630

− 6822

Tm

− 1818

− 3083

− 1833

− 3610

− 1588

− 3542

− 1674

− 7007

Yb

− 1813

− 3077

− 1830

− 3604

− 1583

− 3512

− 1662

− 6956

Lu

− 1807

− 3073

− 1823

− 3598

− 1576

− 3520

− 1635

− 6843

Y

− 1858

− 3137

− 1853

− 3626

− 1629

− 3562

− 1623

− 6792

Sc

− 1744

− 2990

− 1749

− 3536

− 1512

− 3465

− 1560

− 6528

*Oct-hyd-S: Octa-hydrated sulfates. Italicized values are estimated

2.1 Cut-off pH

When REE-bearing compounds, such as REE phosphates are subjected to dissolution in acids, the chemical reactions involved in such cases can be presented by Eqs. 1 and 2.
$$ \left\langle {\mathrm{RePO}}_4\right\rangle +3\left\{{\mathrm{H}}^{+}\right\}=\left\{{\operatorname{Re}}^{3+}\right\}+\left\{{\mathrm{H}}_3{\mathrm{PO}}_4\right\} $$
(1)
$$ \left\langle {\mathrm{ReCO}}_3\mathrm{F}\right\rangle +3\left\{{\mathrm{H}}^{+}\right\}=\left\{{\operatorname{Re}}^{3+}\right\}+\left\{{\mathrm{H}}_2{\mathrm{CO}}_3\right\}+\left\{\mathrm{H}\mathrm{F}\right\} $$
(2)

[The symbols, < > and { } represent the solid and liquid phase, respectively.]

It is important to realize that the solubility of REE-phosphates and fluoro-carbonates cannot be accurately determined by the simple Eqs. 1 and 2, since the reaction products of these equations will be subjected to dissociation or speciation in solution, which often dramatically changes the ultimate solubility of these solids. However, these equations can be useful in examining the relative solubility of various REE-bearing compounds under given chemical conditions.

The pH of the solution will be uniquely determined for a given amount of Re3+ in solution when the concentration of H3PO4 is known in Eq. 1. Such pH is henceforth defined as the cut-off pH for the given conditions [10]. The cut-off pH is similarly defined for fluoro-carbonates in Eq. 2. For monazite, Re in RePO4 could be Ce, La, Nd, and Sm, since primarily four kinds of monazite exist in nature. They are often referred to as Ce-rich, La-rich, Nd-rich, and Sm-rich monazite respectively. Therefore, the solubility of monazite could be very much dependent upon the type of REEs associated with the mineral and therefore, the cut-off pH will also vary for the given composition. On the other hand, the typical chemical formula for bastnaesite, being a fluoro-carbonate, is (Ce, La, Y)CO3F, and it also contains minor amounts of REEs such as Gd, Dy, Er, Yb, Tb, Ho, Tm, and Lu.

The cut-off pH is illustrated in Fig. 2 for various REE compounds when the concentration of REEs in solution is known, for example, a value of 10−3 M is used in the paper unless otherwise specified. Equations 1 to 7 represent the chemical reactions considered in this case. In this figure, the concentrations of other reaction products such as {H3PO4}, {H2CO3}, and {HF} are assumed to be the stoichiometric amounts as given by these reactions.
$$ \left\langle {\operatorname{Re}}_2{\mathrm{O}}_3\right\rangle +6\left\{{\mathrm{H}}^{+}\right\}=2\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{{\mathrm{H}}_2\mathrm{O}\right\} $$
(3)
$$ \left\langle {\operatorname{Re}}_2{\left(C{O}_3\right)}_3\right\rangle +6\left\{{\mathrm{H}}^{+}\right\}=2\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{{\mathrm{H}}_2{\mathrm{CO}}_3\right\} $$
(4)
$$ \left\langle {\mathrm{RePO}}_4\right\rangle +3\left\{{\mathrm{H}}^{+}\right\}=\left\{{\operatorname{Re}}^{3+}\right\}+\left\{{\mathrm{H}}_3{\mathrm{PO}}_4\right\} $$
(5)
$$ \left\langle \operatorname{Re}{\mathrm{F}}_3\right\rangle +3\left\{{\mathrm{H}}^{+}\right\}=\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{\mathrm{H}\mathrm{F}\right\} $$
(6)
$$ \left\langle {\operatorname{Re}}_2{\left({C}_2{O}_4\right)}_3\right\rangle +6\left\{{\mathrm{H}}^{+}\right\}=2\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{{\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4\right\} $$
(7)
Fig. 2

Cut-off pH of various REE compounds for the concentration of Re3+ in the solution being 10−3 M and the stoichiometric amounts of other products in the absence of further reactions of these reaction products

As mentioned earlier and also will be discussed in detail in the following section, such figure shows only the relative solubility in relation to various rare earth elements and does not give the true cut-off pH of these elements. Figure 2 seems to indicate that Re-oxide (hydrated) is the easiest compound to dissolve, followed by carbonate, phosphate, fluoride, sulfate, fluoro-carbonate, and oxalate in that order.

3 Composition Effect

In minerals occurring in nature, it is quite often that they are contaminated by components that are not part of the crystal structure, and hence, the dissolution behavior of such contaminated minerals can be quite unpredictable. In addition, irregularity of chemical composition of a given mineral can behave in an irrational fashion. For example, there are many different compositions of bastnaesite occurring in nature, and the leaching behavior of this mineral in acids could be greatly affected by the chemical composition, particularly due to a relative abundance of fluoride or carbonate. As an illustration of such behavior, in Fig. 3, the cut-off pH lines were established for bastnaesite containing three different compositions, in which the ratios of F:CO3 are 55:45, 50:50, and 33.5:66.5, designated as 45%, 50%, and 66.5% CO3.
Fig. 3

Cut-off pH for three different compositions of bastnaesite

As expected, when bastnaesite consists of more fluoride than carbonate, the resulting cut-off pH is reduced. This implies that when more fluoride is present, the dissolution is more difficult.

A similar behavior of the leaching of bastnaesite with HCl was observed in the literature. For example, Bian et al. [23] observed two different leaching patterns from bastnaesite. The leaching behavior of Re2.F(CO3)3 was found to be 89.6% and that of ReF3 was 1.5% at 90 °C in 90 min using 6 mol/L HCl.

4 Effect of Anions on the Solubility of REE Compounds

4.1 REE-Phosphate System

As discussed earlier, the cut-off pH is defined by the assumption that the reaction products of the dissolution of REE-bearing minerals are only the dissolved REE ions and phosphoric acid in the case of Eq. 1, and the dissolved REE ions, carbonic acid, and hydrofluoric acid in Eq. 2. For example, if cerium phosphate, <CePO4> is placed in a solution containing an acid, the products will be free cerium ion, Ce3+, and H3PO4 as indicated by Eq. 1. It is true that as soon as H3PO4 is produced in the solution, this species is subjected to dissociation to various species such as H2PO4, HPO4=, PO43−, etc. and the distribution of these species varies as the pH of the solution changes. However, when the pH of the solution is less than 2, H3PO4 is the predominant species [12] and the range of pH considered in this discussion, which is 0–2, under which only H3PO4 can be considered without introducing significant errors.

However, Ce3+, the other product in the above equation, is expected to associate with anions in the solution such as SO4=, NO3, and Cl depending upon the kinds of acids used to adjust the pH of the system. The possible association of OH can be ignored in this discussion, simply because such interaction will not be significant in high acidity. The cerium-bearing species considered in this discussion are {CeSO4+}, {Ce(SO4)2}, and {Ce2(SO4)3} with sulfuric acid, and {Ce(NO3)2+} and {Ce(NO3)3} with nitric acid, and {CeCl2+}, {CeCl2+}, {CeCl3}, and {CeCl4} with hydrochloric acid.

As soon as Re3+ is produced by the reaction with acids as shown in Eq. 1, the produced Re3+ will interact with anions of the acids used and combine with these anions to produce many other soluble complexes as given in Eqs. 812, when HCl is used. Similar chemical reactions can be formulated for HNO3.
$$ {\displaystyle \begin{array}{c}\left\{{\operatorname{Re}}^{3+}\right\}+\left\{\mathrm{C}{1}^{-}\right\}=\left\{\operatorname{Re}{\mathrm{C}1}^{2+}\right\}\\ {}{\mathrm{K}}_1=1.24\times {10}^1\end{array}} $$
(8)
$$ {\displaystyle \begin{array}{c}\left\{{\operatorname{Re}}^{3+}\right\}+2\left\{\mathrm{C}{1}^{-}\right\}=\left\{\operatorname{Re}{{\mathrm{C}1}_2}^{+}\right\}\\ {}{\mathrm{K}}_2=6.49\end{array}} $$
(9)
$$ {\displaystyle \begin{array}{c}\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{\mathrm{C}{1}^{-}\right\}=\left\{\operatorname{Re}{\mathrm{C}1}_3\right\}\\ {}{\mathrm{K}}_3=2.99\end{array}} $$
(10)
$$ {\displaystyle \begin{array}{c}\left\{{\operatorname{Re}}^{3+}\right\}+4\left\{\mathrm{C}{1}^{-}\right\}=\left\{\operatorname{Re}{{\mathrm{C}1}_4}^{-}\right\}\\ {}{\mathrm{K}}_4=1.20\end{array}} $$
(11)
$$ {\displaystyle \begin{array}{l}\begin{array}{c}\left\{{\operatorname{Re}}^{3+}\right\}+3\left\{\mathrm{C}{1}^{-}\right\}=\left\langle \operatorname{Re}\mathrm{C}{1}_3\right\rangle \\ {}{\mathrm{K}}_{\mathrm{sp}}=2.76\times {10}^{15}\end{array}\\ {}\left[K\ values\ are\ for\ cerium\right]\end{array}} $$
(12)

However, when sulfuric acid is used, Eqs. 1317 become the relevant chemical reactions.

$$ {\displaystyle \begin{array}{c}\left\{{\mathrm{Ce}}^{3+}\right\}+\left\{{{\mathrm{SO}}_4}^{=}\right\}=\left\{{{\mathrm{Ce}\mathrm{SO}}_4}^{+}\right\}\\ {}{\mathrm{K}}_1=5.85\times {10}^4\end{array}} $$
(13)
$$ {\displaystyle \begin{array}{c}\left\{{\mathrm{Ce}}^{3+}\right\}+2\left\{{{\mathrm{SO}}_4}^{=}\right\}=\left\{\mathrm{Ce}{{\left({\mathrm{SO}}_4\right)}_2}^{-}\right\}\\ {}{\mathrm{K}}_2=1.47\times {10}^5\end{array}} $$
(14)
$$ {\displaystyle \begin{array}{c}2\left\{{\mathrm{Ce}}^{3+}\right\}+3\left\{{{\mathrm{SO}}_4}^{=}\right\}=\left\{{\mathrm{Ce}}_2{\left({\mathrm{SO}}_4\right)}_3\right\}\\ {}{\mathrm{K}}_3=1.03\times {10}^3\end{array}} $$
(15)
$$ {\displaystyle \begin{array}{c}2\left\{{\mathrm{Ce}}^{3+}\right\}+3\left\{{{\mathrm{SO}}_4}^{=}\right\}=\left\langle {\mathrm{Ce}}_2{\left({\mathrm{SO}}_4\right)}_3\right\rangle \\ {}{\mathrm{K}}_{\mathrm{spl}}=4.56\times {10}^5\end{array}} $$
(16)
$$ {\displaystyle \begin{array}{c}2\left\{{\mathrm{Ce}}^{3+}\right\}+3\left\{{{\mathrm{SO}}_4}^{=}\right\}=8\left\{{\mathrm{H}}_2\mathrm{O}\right\}=\left\langle \mathrm{Ce}{\left({\mathrm{SO}}_4\right)}_38{\mathrm{H}}_3\mathrm{O}\right\rangle \\ {}{\mathrm{K}}_{\mathrm{sp}2}=3.33\times {10}^{24}\end{array}} $$
(17)
$$ {\displaystyle \begin{array}{c}\left\{{{\mathrm{H}\mathrm{SO}}_4}^{-}\right\}+\left\{{\mathrm{H}}^{+}\right\}+\left\{{{\mathrm{SO}}_4}^{=}\right\}\\ {}{\mathrm{K}}_0=1.24\times {10}^{-2}\end{array}} $$
(18)
Table 2 gives the total concentration of REEs including all of the relevant dissolved species when HCl, HNO3, and H2SO4 are used, while keeping the concentration Re3+ at 10−3 M. As noted from the table, the total dissolved REEs are more than expected. The total concentration of REEs is about 3.4 times that of Re3+ when HCl is used, while it is about 1.58 times with HNO3 but the total concentration of dissolved REEs is more than 17,000 times with sulfuric acid. This includes the adverse effect of REEs being precipitated as octa-hydrated sulfate. In summary, among the three acids, sulfuric acid is by far the best acid that maximizes REE dissolution. The potential precipitation of metal ions with chloride, nitrate, and sulfate (anhydrate) was also considered but these precipitations were irrelevant under the conditions studied in this investigation. It should be noted that all the calculations in this study have been done manually using an Excel spread sheet.
Table 2

Concentration of the total REEs dissolved for three different acids when the concentration Re3+ is kept at 10−3 M

 

HCl

HNO3

H2SO4

La

1.57E−03

2.07E−03

8.84E−04

Ce

2.31E−03

3.91E−03

2.06E+01

Pr

1.25E−03

1.57E−03

3.70E−04

Nd

1.26E−03

1.70E−03

8.24E−03

Sm

1.23E−03

1.66E−03

7.94E+01

Eu

1.22E−03

1.80E−03

1.34E+00

Gd

1.43E−03

1.65E−03

6.87E−02

Tb

1.20E−03

1.36E−03

1.16E+00

Dy

1.21E−03

1.16E−03

1.01E+02

Ho

1.20E−03

1.42E−03

5.04E+00

Er

1.21E−03

1.15E−03

1.95E+01

Tm

1.18E−03

1.17E−03

1.51E+00

Yb

1.15E−03

1.22E−03

3.90E+01

Lu

1.10E−03

1.43E−03

9.68E−01

Y

3.12E−02

1.01E−03

8.48E+00

Sc

4.75E−03

1.00E−03

2.52E−05

Avg

3.40E−03

1.58E−03

1.74E+01

The italicized values represent the total concentration of REEs limited by precipitation as REE-octa-hydrated sulfate

The equilibrium concentrations of the free cerium ion, Ce3+ at pH values of 2, 1, and 0 in contact with the solid phase, <RePO4>, can be calculated using Eq. 1, in which sulfuric, nitric, and hydrochloric acid are used to dissolve the compound. The results are shown in Table 3. When there is no influence of anions supplied by the acids, the concentration of the free cerium ion is the same for all three acids. However, when the cerium ion interacts with the anions produced by the acids, the total concentration of cerium-bearing species is introduced and the concentration of the total dissolved cerium species is much more than the free cerium ion alone. Figure 4 demonstrates the concentrations of the various cerium-bearing species for pH 2, 1, and 0.
Table 3

Concentration in mol/L of Ce3+ and the total cerium dissolved in various acids at 25 °C

 

Ce3+ in mol/L

Total cerium species in mol/L

pH

2

1

0

2

1

0

H2SO4

3.34E−11

1.06E−09

3.34E−08

2.45E−08

2.18E−05

5.12E−02

HNO3

3.34E−11

1.06E−09

3.34E−08

4.31E−11

4.13E−09

1.39E−06

HCl

3.34E−11

1.06E−09

3.34E−08

3.8E−11

2.44E−09

8.07E−07

Fig. 4

Concentration of cerium species at pH of 2, 1, and 0, when sulfuric acid is used

In these calculations, Eq.1 together with Eqs. 13 through 17 were used. It is striking to note that the total concentration of cerium-bearing species is more than 700-fold of the free cerium ion, Ce3+ at pH 2, because of the speciation of cerium into Ce(SO4)+, Ce(CO4)2 and Ce2(SO4)3. The results are much more pronounced at pH 1 and 0. It is noted that the increase in the total concentration of Ce-species due to the speciation with nitric and hydrochloric acid is far less than those with sulfuric acid. It is assumed that the concentration of SO4= is determined by Eq. 18 for a given pH.

It is noted that the results of nitric and hydrochloric acids are almost identical as seen in Table 3 as well as Figs. 5 and 6, which is a significant contrast to sulfuric acid as shown in Fig. 4. Such differences are even more remarkable with other REE ions such as lanthanum, praseodymium, neodymium, gadolinium, and dysprosium due to low solubility of Re2(SO4)3.8(H2O), where Re represents Pr, Nd, Gd, Dy, etc.
Fig. 5

Concentration of cerium species at pH of 0, 1, and 2, when nitric acid is used

Fig. 6

Concentration of cerium species at pH of 0, 1, and 2, when hydrochloric acid is used

Unlike the case of La, Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc, the equilibrium concentration of Ce3+ due to Eq. 1 is always less than that given by Eq. 17. As soon as Ce3+ is released in the solution containing sulfuric acid for example, Ce3+ will interact with sulfate forming various sulfate species as shown by Eqs.13–17.

As in the case of La, Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc, the equilibrium concentration of dissolved REE ions given by Eq. 17 is such that the concentration of the dissolved REE species in equilibrium with their hydrated sulfate is much lower than that given by Eq. 1, and as a result, the total REE ions in the solution can be far less than that given by Eq. 1. For example, in the case of La3+, the concentration given by Eq.1 at pH = 0, is 3.31E−07, while it is 1.18E−09 at the same pH governed by Eq. 17. Such behavior of La is shown in Table 4 along with the values for Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc. As long as the solid phase of <RePO4> is present in the system, the concentration of free Re3+ ion will try to satisfy Eq. 1 and also Eq. 17.
Table 4

Equilibrium concentration of free REE ions and the total REE species at pH = 0 with sulfuric acid in the absence of hydrated sulfate precipitation and presence of precipitation

 

In the absence of ppt

In the presence of ppt

 

Re3+

Re-total*

Re3+

Re-total*

La

3.31E−07

2.20E−02

1.18E−09

2.21E−03

Pr

6.48E−06

7.37E+00

9.44E−13

1.07E−06

Nd

8.36E−06

6.78E+00

2.32E−05

2.86E−11

Eu

6.22E−03

8.16E+03

6.48E−05

8.49E+01

Gd

1.64E−05

9.16E+00

2.92E−10

1.63E−04

Dy

3.77E−04

4.32E+02

3.40E−08

3.89E−02

Ho

1.82E−02

7.05E+03

3.40E−08

1.29E−02

Y

1.49E−02

6.78E+03

1.81E−03

8.24E+02

Sc

3.77E−02

2.26E+10

1.33E−20

7.98E−09

*Re-total refers to the total dissolved Re-bearing species

The concentration of free Re3+ and its hydrated sulfate complexes will be governed by both reactions trying to sustain the dynamic equilibrium governed by both Eqs. 1 and 17. However, when the Re-phosphate solid disappears from the system, the final concentrations of these ions will be thoroughly governed by Eq. 17. For example, if part of the solution is taken out of the solution for chemical analysis, Re3+ is expected to decrease gradually due to precipitation governed by Eq. 17. This will occur in the absence of the solid RePO4.

On the other hand, for cerium, samarium, terbium, erbium, thulium, ytterbium, and lutetium, Eq. 1 prevails over Eq. 17, and there will be no precipitation of Re-hydrated sulfate precipitation until all of the solid Re-phosphate disappears. The resulting concentrations of free REE ions are shown in Table 5 under this condition.
Table 5

Equilibrium concentration of free REE species and the concentration of the total dissolved REEs at pH = 0 with sulfuric acid, in which there is no hydrated sulfate precipitation

 

In the absence of ppt

 

Re3+

Re-total

Ce

1.34E−07

2.04E−01

Sm

6.00E−06

4.69E+01

Tb

9.66E−05

1.42E+00

Er

1.31E−02

2.44E+04

Tm

3.18E−02

1.98E+03

Yb

1.60E−04

6.08E+02

Lu

1.69E−01

1.63E+03

Referring to Table 5, it can be seen that the concentration of the free cerium ion, Ce3+ at pH 0 is calculated to be 1.34 × 10−7 M using Eq. 1, while the total dissolved cerium species produced in the solution at this pH would be 2.04 × 10−1!

As discussed earlier, in the case of La, Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc, the precipitation reaction given by Eq. 17 prevails and therefore, the effect of Re-hydrated sulfate precipitation influencing the distribution of the dissolved Re-ion species is quite different from the case with Ce, Sm, Tb, Er, Tm, Yb, and Lu. The results are presented in Figs. 7 and 8 for praseodymium as an example. Such precipitation of light REEs (LREEs) has been reported in the literature, in which LREEs (La-Sm) were precipitated as double salt sulfates by adding Na2SO4 into the solution containing these ions [24, 25].
Fig. 7

Concentration of praseodymium species at pH of 0, 1, and 2, when sulfuric acid is used using Eqs. 1 and 13 through 16. Precipitation of hydrated sulfate is missing

Fig. 8

Concentration of praseodymium species at pH of 0, 1, and 2, when sulfuric acid is used using Eqs. 1 and 13 through 17, when precipitation of hydrated sulfated is present

The results when Eq. 17 is introduced are shown in Fig. 7. It is noted that the concentrations of Pr-bearing species are almost identical for three different pHs. This is an unexpected phenomenon but can be explained by the fact that the hydrated sulfate precipitation is greater at higher acidity, because the concentration of sulfate ion, SO4= is higher at low pH. This result can explain at least partially the anomaly observed by Kim et al. [1], in which the total recovery REEs from REE-bearing ores increased with the concentration of sulfuric acid until 5 mol/L, at which point the total recovery dropped significantly.

4.2 REE Fluoro-carbonate System

The effect of speciation of REEs with anions released from REE fluoro-carbonates has also been examined. In this study, the solubility of Ce-bastnaesite has been considered and the results are compared with those of CePO4.

As seen in Table 6, the concentrations of total cerium species dissolved in three different acids from the Ce-bastnaesite are, in general, more than those from Ce-phosphates. This is consistent with the notion that bastnaesite is generally easier to dissolve in acid than monazite and xenotime.
Table 6

Concentrations of total Ce-species in mol/L from CePO4 and CeFCO3 in three different acids at 25 °C

 

From CeFCO3

From CePO4

pH

0

1

2

0

1

2

H2SO4

2.04E−01

2.75E−04

9.80E−07

5.12E−02

1.06E−09

3.34E−11

HNO3

5.56E−06

5.22E−08

1.72E−09

1.39E−06

4.13E−09

4.31E−11

HCl

3.22E−06

3.09E−08

1.576E−09

8.07E−07

2.44E−09

3.76E−11

4.3 Effect of Concentration of Anions

The effect of anions in hydrometallurgical processes has been discussed elsewhere [26, 27]. It is apparent that when the concentration of anions responsible in making complexation increases, the concentration of the resulting metal-complex will also increase. For example, as the concentration of the chloride ion in Eqs. 811 increases, the concentrations of the REE-chloride complexes will increase.

Figures 9, 10, and 11 represent the concentration of REE-complexes as the concentration of chloride, nitrate, or sulfate increases from 0.001 to 5 mol/L at pH 1.
Fig. 9

Concentration of cerium-chloride complexes as a function of the concentration of chloride ion at pH 1

Fig. 10

Concentration of cerium-nitrate complexes as a function of the concentration of nitrate ion at pH 1

Fig. 11

Concentration of cerium-sulfate complexes as a function of the concentration of sulfate ion at pH 1

It should be noted that the increase of the total dissolved cerium concentration is about 1330, 1560, and 6.08 × 105 times for chloride, nitrate, and sulfate respectively, as the concentration of these anions increase from 0.001 to 5 M at pH 1. Furthermore, as seen in Table 7, the overall concentration of dissolved cerium with 5 M of sulfate at pH 1 is nearly same as that at pH 0 with 1 M of sulfate. Such behavior illustrates that a wise selection of anions can lead to a practical implication in the production of metal ions in hydrometallurgy. For example, it could be practical advantages to obtain a better recovery of REEs with one tenth of the consumption of acids as shown in the above example.
Table 7

Concentrations of Ce-species in mol/L at pH 1 and 0 as the concentration of SO4= increases from 0.001 to 5 mol/L

pH = 1 with H2SO4

 SO4=

0.001

0.01

0.1

1

5

 Ce3+

1.06E−09

1.06E−09

1.06E−09

1.06E−09

1.06E−09

 Ce(SO4)^+

6.19E−08

6.19E−07

6.19E−06

6.19E−05

3.10E−04

 Ce(SO4)2^−

1.56E−09

1.56E−07

1.56E−05

1.56E−03

3.89E−02

 Ce2(SO4)3

1.15E−24

1.15E−21

1.15E−18

1.15E−15

1.44E−13

 Ce Tot

6.45E−08

7.76E−07

2.18E−05

1.62E−03

3.92E−02

pH = 0 with H2SO4

 SO4=

0.001

0.01

0.1

1

5

 Ce3+

3.34E−08

3.34E−08

3.34E−08

3.34E−08

3.34E−08

 Ce(SO4)^+

1.96E−06

1.96E−05

1.96E−04

1.96E−03

9.79E−03

 Ce(SO4)2^−

4.92E−08

4.92E−06

4.92E−04

4.92E−02

1.23E+00

 Ce2(SO4)3

1.15E−21

1.15E−18

1.15E−15

1.15E−12

1.44E−10

 Ce Tot

2.04E−06

2.45E−05

6.88E−04

5.12E−02

1.24E+00

It should be noted that it is practically impossible to maintain the concentration of sulfate more than 0.01 M at pH 0, since the ration of SO4= to HSO4 is 1, 0.1, and 0.01 as the pH of the solution is 2, 1, and 0.

It should be noted, however, that when a similar calculation is made for lanthanum, the advantage of gaining a higher recovery by adding sulfate is not as much as in the case of cerium due to precipitation of La-octa-hydrated sulfates. This is shown in Table 8. It is also noted that Sm, Tb, Er, Tm, Yb, and Lu follow a similar behavior as Ce, while Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc follow the trend of La, although the extent of precipitation varies with different REEs.
Table 8

Concentrations of La-species in mol/L at pH 1 and 0 as the concentration of SO4= increases from 0.001 to 5 mol/L

pH = 1 with H2SO4

 SO4=

0.001

0.01

0.1

1

5

 La3+

1.05E−08

1.05E−08

1.05E−08

1.18E−09

1.06E−10

 La(SO4)^+

3.47E−07

3.47E−06

3.47E−05

3.92E−05

1.76E−05

 La(SO4)2^−

1.83E−08

1.83E−06

1.83E−04

2.06E−03

4.61E−03

 La2(SO4)3

8.38E−12

8.38E−09

8.38E−06

1.07E−04

1.07E−04

 La Tot

3.76E−07

5.32E−06

2.26E−04

2.21E−03

4.74E−03

pH = 0 with H2SO4

 SO4=

0.001

0.01

0.1

1

5

 La3+

3.31E−07

3.31E−07

3.74E−08

1.18E−09

1.06E−10

 La(SO4)^+

1.10E−05

1.10E−04

1.24E−04

3.92E−05

1.76E−05

 La(SO4)2^−

5.78E−07

5.78E−05

6.53E−04

2.06E−03

4.61E−03

 La2(SO4)3

8.38E−09

8.38E−06

1.07E−04

1.07E−04

1.07E−04

 La Tot

1.19E−05

1.76E−04

8.84E−04

2.21E−03

4.74E−03

The italicized values represent the concentration of La-species limited by precipitation as La-octa-hydrated sulfate

5 Summary

The physics and chemistry of REEs are very similar reflecting a generally consistent trend in their thermodynamic values.

The cut-off pH calculations have indicated that the solubility of various REE compounds follows the order anhydrated oxides > hydrated oxides > carbonates > sulfates > fluorides > phosphates > oxalates.

The equilibrium concentrations of REEs from REE-bearing minerals are influenced greatly by various anions provided by different acids, namely, sulfate, nitrate, and chloride. In general, sulfuric acid offers higher equilibrium concentrations of REEs, while the behaviors of nitric and hydrochloric acids are almost identical offering lower concentrations of REE-bearing species. In the case of Re-phosphates, some REEs such as La, Pr, Nd, Eu, Gd, Dy, Ho, Y, and Sc are greatly affected by precipitation of hydrated sulfates, Ce, Sm, Tb, Er, Tm, Yb, and Lu are not significantly affected by precipitation.

As the concentration of anions supplied by the three acids increases, the overall dissolved REE concentration increases dramatically indicating the significant role of anions in the overall leaching behavior of REE compounds.

These trends based on thermodynamic calculations deserve a close examination by a series of tests in laboratories and operating plants, which will undoubtedly help assist in better understanding of the leaching behavior of various REE minerals in acidic media.

Notes

Acknowledgements

The author would like to express his thanks to Dr. Jim Gebhardt for his valuable comments on the paper and also to Dr. Rina Kim for sharing some of thermodynamic data.

Compliance with Ethical Standards

Conflict of Interest

The author declares that there is no conflict of interest.

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Copyright information

© The Society for Mining, Metallurgy & Exploration 2018

Authors and Affiliations

  1. 1.Department of Materials and Metallurgical EngSouth Dakota School of Mines and TechnologyRapid CityUSA

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