The Synthesis and Characterization of Rare-Earth Fluocarbonates

Abstract

A procedure is described for synthesizing cerium fluorcarbonate. This synthetic bastnaesite was characterized through X-Ray diffraction, chemical analysis and thermogravimetric studies. The solubility product of Ce-bastnaesite was determined to be 10^−10.1. The stability constant pK CeHCO₃²⁺ complex was experimentally determined to be 3.04 ± 0.21. Based on the thermodynamic data available in the literature, a speciation diagram has been prepared for the synthetic cerium bastnaesite.

Introduction

Owing to their unique magnetic and optical properties, the world demand for rare-earths has been steadily increasing in the latter half of the twentieth century and sharply since the 1990s. The discovery of high-temperature superconductivity in mixed oxides involving rare-earths opened up yet another market avenue for them. A recent report prepared by the US Department of Energy highlights the critical role of rare-earth metals in the clean energy economy¹⁾. The need to exploit all available rare-earth resources is expected to grow rapidly. Rare-earths occur in nature mainly as salt minerals, most commonly as complex fluocarbonates and phosphates. Bastnaesite, a fluocarbonate of the cerium group of rare-earth metals. (RE)FCO₃, was first discovered in 1818 at Bastnas, Sweden²⁾. The mineral contains about 75 percent light rare-earth (RE) oxides distributed as follows: 50.0% cerium, 34.0% lanthanum, 11.0% neodymium, 4.0% praseodymium, 0.5 % samarium, 0.2% gadolinium, 0.1% europium, and 0.2% others. The theoretical composition of bastnaesite (CeFCO₃ synthetic or (Ce, La)FCO₃ natural) is 74.77% Ce₂O₃, 20.17% CO₃, and 8.73% F. The elemental analysis of natural bastnaesite is generally close to this composition.

Major occurrences of bastnaesite have been reported in several locations, but by far the two largest deposits known are those at Mountain Pass in California and in Inner Mongolia and Kanshu Province, China. The largest single producer of rare-earths was the Molycorp operation at Mountain Pass^3–5) until low-cost Chinese production caused it to be shut down in the late 1990s. Extensive efforts are underway to reopen the Mountain Pass mine since today 97 percent of the world’s production comes from China.

Being a relatively rare mineral, little data are available in the literature about the various physico-chemical properties of bastnaesite. To provide information about pure particles for study of this important source of rare-earth metals, this paper summarizes some of our results on the synthesis and characterization of pure bastnaesite.

Crystal structure of bastnaesite

Bastnaesite belongs to a family of rare-earth/alkaline-earth fluocarbonates (those of the calcium and barium varieties) and is similar in structure to calcium fluorcarbonate. In 1956, Donnay and Donnay⁶⁾ established the following isostructural series of calcium fluocarbonates, with different ratios of CeFCO₃ to CaCO₃:

Bastnaesite	CeFCO3	100% CeFCO3
Parisite	2CeFCO3 · CaCO3	2/3 CeFCO3
Rontgenite	3CeFCO3 · 2CaCO3	3/5 CeFCO3
Synchisite	CeFCO3 · CaCO3	1/2 CeFCO3

Semenov⁷⁾ has suggested that additional members of this series with even higher calcium contents might exist:

2CeFCO3 · CaCO3	2/5 CeFCO3
2CeFCO3 · 2CaCO3	1/3 CeFCO3
Vaterite	0 % CeFCO3

McConnel⁷⁾ later showed that calcite and vaterite are isostructural, vaterite being a special trigonal polymorphous modification of CaCO₃ that differs from calcite. The minerals of this group have a characteristic layer structure, in which the layers of Ce and F ions are parallel to the (0001) plane and alternate with the layer of Ca ions. Plane CO₃ triangles are arranged between the Ce-F and Ca layers. Details of the crystal structures of these minerals are available in the excellent paper by Donnay and Donnay⁶⁾.

The structural scheme of bastnaesite and related minerals described above was first predicted by Oftedal⁶⁾ in 1931 and subsequently confirmed by the Donnays⁶⁾. The structure of bastnaesite is trigonal with a_o = 7.16A, c_o = 9.79A, with the c_o/a_o ratio being 1.367. Cerium and fluorine ions alternate at the vertices of regular hexagons with carbonate groups interspersed between the Ce-F layers. There are six formula units of CeFCO₃ per cell.

Synthesis of Bastnaesite

In our laboratory, we tried to separate some pure bastnaesite crystals from a sample of Mountain Pass ore through physical separation under ultra violet light⁵⁾. However, this was done with only limited success and, therefore, for very precise characterization of this mineral, it was necessary to synthesize pure bastnaesite in the laboratory. Only one previous attempt in this direction was reported in the literature by Jansen et al.⁸⁾, who were able to successfully synthesize a few milligrams of bastnaesite in the early 1950’s.

We modified the Jansen method in order to synthesize about one gram of bastnaesite in each run. The reaction involved in the synthesis is the following:

Cerium Bastnaesite:   

$$\text{Ce}{\left({\text{CO}}_3\right)}_3+2\text{HF}=2\hspace{0.17em}{\text{CeFCO}}_3+{\text{CO}}_2+{\text{H}}_2\text{O}$$(1)

or

Lanthanum Bastnaesite:   

$${\text{La}}_2{\left({\text{CO}}_3\right)}_3\cdot 8{\text{H}}_2\text{O}+2\text{HF}=2\hspace{0.17em}{\text{LaFCO}}_3+{\text{CO}}_2+9{\text{H}}_2\text{O}$$(2)

The apparatus used for the synthesis (Fig. 1) consisted of a cylindrical container (11 cm × 18 cm) with a capacity of about 1 liter. It was constructed of Monel (resistant to hydrofluoric acid at high temperature), This was surrounded by a heating jacket (Glas-Col Heating Mantle for Griffin type beakers, made up of corrosion-resistant glass cloth with Nichrome heating elements) capable of providing controlled temperatures up to 450 °C through a Variac. The container was closed with a Teflon lid, through with a Teflon stirrer and Teflon inlet and outlet tubes (0.6 cm diameter) for CO₂ gas were provided. An alternative method for stirring could be provided by bubbling the CO₂ gas through the solution, with a porous Teflon tube introduced into the solution. The outgoing gases and vapors were condensed through a condenser attached to the lid as shown in the figure and supported by a stand. Bone-dry CO₂ gas was bubbled through the solution so that the reaction was carried out in a CO₂ - saturated solution. Hydrofluoric acid was introduced through a vertical Teflon tube drop by drop over a period of a few minutes so that it was thoroughly mixed with the contents of the container. Basically, the procedure involved dispersing about 1.150 g of cerium carbonate or (1.505 g of lanthanum hydrous carbonate) in 700 cc of preheated distilled water. Then 50 cc of 0.1 M HF acid (100 mg of HF) was introduced slowly into the solution, and the reaction was allowed to take place at around 80 – 90 °C in a CO₂ atmosphere for about 24 hours. Digestion of this product overnight at this temperature was found sufficient for the reaction to go to completion, after which the product was cooled, filtered, and air dried. Under unfiltered UV light, a characteristic light green color of bastnaesite was visible in the product.

Fig. 1

Schematic drawing of the apparatus used for the laboratory synthesis of bastnaesite in the laboratory.

The synthesis of LaFCO₃ would essentially involve the same procedure. Bastnaesite of the composition as found in nature (having both La and Ce as major constituents) can also be synthesized starting with the stoichiometric amounts of respective carbonates and hydrofluoric acid as reactants.

It is also important to point out that the limitation of this method for producing larger quantities of bastnaesite is due to the limited capacity (l liter) of the container. Since the solid/liquid ratio was found to affect the purity of the sample, it is impossible to work at higher concentrations.

Characterization of Synthetic Bastnaesite

Chemical Analysis of Synthetic Bastnaesite

Four different batches were selected for analysis and identification purposes. The only differences among these is the amount of cerium carbonate per 700 cc used in the synthesis, batches A and B starting with 0.5753 g, and batch C with 0.4602 g of cerium carbonate. Batch D had several hours of ultrasonic dispersion of the cerium carbonate before the HF was added. The chemical analyses, which were carried out at the Mountain Pass Research Laboratory of Molycorp (Molybdenum Corporation of America) are summarized in Table 1. A comparison of the chemical analysis of the synthetic products with the theoretical composition of bastnaesite indicates that the synthetic product is close to pure bastnaesite.

Table 1 Analysis of synthetic cerium bastnaesite and natural mineral samples

Sample	% Re2O3	% Ce or RE	% F
Theoretical CeFCO3	74.9	63.95	8.67
Synthetic CeFCO3
Batch A	70.2	59.9	9.3
Batch B	71.3	60.9	9.0
Batch C	69.96	59.7	8.6
Batch D		63.15	8.83
Natural (Molycorp) bastnaesite crystal (Ce,La)FCO3	69.17	59.04	6.88
Hand picked* (under UV light) Birthday Claims, Mountain Pass sample (Ce,La)FCO3	60.9	51.98	-

*  Bastnaesite fluoresces (green color) in UV light (without filter) and so it can easily be hand picked.

Specific Surface of the Synthetic Ce-Bastnaesite Powder

Using the multi-point BET method, the specific surface area of the powder was measured by the adsorption of nitrogen at liquid nitrogen temperature. It was found to be 9.27 m²/g, giving an average diameter of the particles of 0.14 μm.

X-Ray Diffraction Studies

A Philips diffractometer was used to obtain the x-ray diffraction patterns for the various synthetic samples of bastnaesite. The beam was Cu Ka with a Ni filter having a wavelength of 1.5418 Å. Table 2 compares the values of the d-spacings reported in the literature with those obtained for the synthetic bastnaesite from this work. The values are quite consistent with those of the bastnaesite (CeFCO₃) sample synthesized by Jansen et al.⁸⁾ and those of the natural bastnaesite from Mountain Pass.

Table 2 X-ray diffraction pattern data for Mountain Pass bastnaesite and for synthesize samples of cerium bastnaesite compared to the synthetic bastnaesite synthesized by Jansen et al.⁸⁾

Jansen’s Synthetic		Natural Bastnaesite		Synthetic CeFCO3 (this work)
CeFCO3		(Ce, La)FCO3 (Molycorp)		Batch A	Batch B	Batch C	Batch D
d, Å	I	d, Å	I	d, Å	d, Å	d, Å	d, Å
4.96	70	4.88		4.874	4.874	4.874	4.901
3.52	100	3.564	70	3.548	3.562	3.562	3.568
2.92	100	2.879	100	2.867	2.876	2.876	2.881
2.46	9	2.445	10	2.442	2.442	2.442	2.446
2.29	2	2.273	3	-	-	-	-
		2.238	3	-	-	-	-
2.09	50	2.057	40	2.049	2.053	2.953	2.050
2.03	50	2.016	40	2.006	2.014	2.014	2.019
1.92	35	1.898	40	1.892	1.895	1.895	1.849
1.81	6	1.783	9	1.778	1.778	1.778	1.784
1.70	18	1.674	21	1.670	1.670	1.670	1.676
1.64	2	1.629	1	-	-	-	-
1.59	13	1.573	15	1.568	1.570	1.573	1.573
1.49	9	1.481	9	1.478	1.482	1.481	1.482
1.46	9	1.439	11	1.435	1.437	1.439	1.439
1.37	4	1.347	7	1.344	1.344	1.344	1.345
1.32	15	1.298	15	1.283	1.296	1.296	1.297
1.29	4	1.277	7	1.274	1.276	1.277	-

It is also apparent from the data that the methods employed for synthesis did indeed yield bastnaesite mineral particles. The intensities of the peaks (not shown here) obtained with the synthetic bastnaesite samples also compared quite well with the reported literature values.

Thermogravimetric Studies on Bastnaesite

A Perkin-EImer TGS-2 Thermogravimetric System was used to obtain the weight loss as a function of the temperature of heating. The corresponding first derivative of this curve was also recorded. The heating rate was programmed to be 50 °C per minute through a temperature range of 30 to 900 °C. Since nitrogen was being flushed through the system during heating, the oxygen concentrations must have been minimal. The heating curve as well as its first derivative is plotted in Fig. 2a and b, respectively. When heated in a nitrogen atmosphere, cerium carbonate exhibits two plateaus with the final weight loss at 900 °C asymptotically approaching 30 %, whereas cerium fluocarbonates show an asymptotic weight loss of 20 % at 900 °C. This total weight loss can be understood in terms of the following reactions in a nitrogen atmosphere.

Fig. 2

Thermogravimetric analysis results for cerium carbonate, synthetic bastnaesite and natural bastnaesite (top figure) and the respective differential curves (bottom figure).

1. i) Cerium carbonate Ce₂(CO₃)₃ · anhydrous or xH₂O in a nitrogen environment :   
   $$\begin{array}{lll}{\text{Ce}}_2{\left({\text{CO}}_3\right)}_3\cdot {\text{xH}}_2\text{O}-{\text{N}}_2\text{envir}-\hfill & \to \hfill & {\text{Ce}}_2{\text{O}}_3+3{\text{CO}}_2+{\text{xH}}_2\text{O}\hfill \\ 460.24g\left(\text{anhydrous}\right)\hfill & \hfill & 328.24\text{g}\hfill \end{array}$$(3)
   ∴ Therefore anhydrous cerium carbonate weight loss = 28.7 %
2. ii) Cerium Fluocarbonate CeFCO₃ (synthetic Ce-bastnaesite) in nitrogen atm:   
   $$\begin{array}{lll}3{\text{CeFCO}}_3-{\text{N}}_2\text{envir}-\hfill & \to \hfill & {\text{Ce}}_2{\text{O}}_3+{\text{CeF}}_3+3{\text{CO}}_2\hfill \\ 3\times 210.2\text{g}\hfill & \hfill & 328.24\text{g}+197.12\text{g}\hfill \end{array}$$(4)
   ∴ Synthetic Ce-bastnaesite weight loss = 20.08 %
3. iii) Lanthanum Fluocarbonate LaFCO₃ (synthetic La-bastnaesite) in nitrogen atm:   
   $$\begin{array}{lll}3{\text{LaFCO}}_3-{\text{N}}_2\text{envir}-\hfill & \to \hfill & {\text{La}}_2{\text{O}}_3+{\text{LaF}}_3+3{\text{CO}}_2\hfill \\ 653.73\text{g}\hfill & \hfill & 521.73\text{g}\hfill \end{array}$$(5)
   ∴ Synthetic La-bastnaesite weight loss = 20.2 %
4. iv) Natural Bastnaesite (Ce,La)FCO₃ which contains cerium and lanthanum as the majority fluocarbonates, then would be expected to lose about 20 % of its weight upon decomposition, as was observed in this work.

Since the weight loss obtained for synthetic cerium fluocarbonate was very close to this value, the above reaction is very probable.

Note that here the final products of decomposition are the corresponding oxides and fluorides. In an oxygen environment, cerium can also be oxidized to its tetravalent state since CeO₂ is the most stable oxide product. Then the reaction could be   

$$2{\text{CeFCO}}_3-{\text{O}}_2\hspace{0.17em}\text{envir}-\to 2{\text{CeO}}_2+{\text{F}}_2+{\text{CO}}_2$$(6)

No studies on oxidative heating have been reported in the literature. For rare-earth carbonates precipitated from aqueous solutions, Head and Holley, Jr.⁹⁾ conducted thermal decomposition studies under several different conditions in vacuum, carbon dioxide and water vapor environments. The decomposition reaction has been postulated as follows:   

$${\text{M}}_2{\left({\text{CO}}_3\right)}_3\cdot {\text{xH}}_2\text{O}\to {\text{M}}_2{\left({\text{CO}}_3\right)}_3\to {\text{M}}_2{\text{O}}_2\cdot {\text{CO}}_3\to {\text{M}}_2{\text{O}}_3\cdot {\text{yCO}}_2\to {\text{M}}_2{\text{O}}_3$$(7)

where x varies from compound to compound and the composition M₂O₃ · yCO₂ ( y < 1) is not always observed. The general shape of their curve is consistent with our results except that their cerium carbonate was hydrated (x = 8). The plateau observed around 195 – 200 °C in our case is perhaps due to some intermediate in the decomposition of cerium carbonates as postulated by Head and Holley⁹⁾.

The decomposition peak for bastnaesite from Mountain Pass was found to occur between 500 to 675°C. Vlasov⁷⁾ reported a decomposition peak between 420 and 600 °C for a bastnaesite from Mongolia. X-ray analysis of its calcination product also showed a cubic phase with a = 5.555 Å which was in fact a solid solution of lanthanide sesquioxides CeLaO_3.5.

This is due to the dissociation of the mineral involving the loss of CO₂ and possibly F in an oxygenated environment. Our results also show a shift in the decomposition peak from 600 °C for natural bastnaesite to about 500 °C in the case of synthetic bastnaesite. With the limited work conducted on this system, it is difficult to explain this shift except to note that these two bastnaesites have different rare-earth compositions and that in the case of natural bastnaesite, small amounts of calcite and barite are also present. The presence of impurities apparently stabilizes the fluocarbonate such that decomposition is delayed by 100 °C. It is interesting to note that the final products are the same since the final weight losses are identical. Perhaps the enthalpies of the decomposition reactions in the two cases are different.

Solution Chemistry of Synthetic Bastnaesite

Similar to such minerals as apatite, barite and calcite, bastnaesite is a sparingly soluble salt mineral. The soluble species resulting from dissolution of the mineral form ion complexes in the solution as well as at the water-solid interface. Ionic complexes may participate in interfacial reactions and play a significant role in subsequent processing. The synthetic Ce-bastnaesite was therefore subjected to detailed investigation of its solution chemistry.

Thermodynamic Calculations

For all minerals containing carbonate as lattice anions, the solubility and the corresponding solution behavior is primarily controlled by the partial pressure of CO₂. For a system open to the atmosphere, the activity of carbonate species in water will be determined by the aqueous solubility of CO₂ and the mineral. In addition, Ce-bastnaesite (CeFCO₃) will contribute Ce³⁺ ions, which will undergo hydrolysis and form a series of hydroxo complexes in solution^10–13). Similarly, fluoride ions can form such complexes as HF, HF₂, CeF²⁺. Since fluoride and cerium ions in solution result from the dissolution of bastnaesite, their total concentration will be identical.

Based on reported data in the literature, the following charged species are expected to exist in aqueous solution in equilibrium with bastnaesite through the reactions listed in Table 3: Ce³⁺, HCO₃⁻, CO₃²⁻, F⁻, CeOH²⁺, Ce(OH)₂⁺, Ce(OH)₄⁻, Ce₃(OH)₅⁻, HF₂⁻, CeF²⁺, CeHCO₃²⁺. Except for the solubility product of CeFCO₃, K_SO, and the solubility constant of CeHCO₃²⁺, K_CeHC, (which were determined experimentally in this work) all the values of equilibrium constants were taken from the published literature.

Table 3 Chemical Reactions in the Synthetic Cerium Bastnaesite-Water System Open to the Atmosphere

			Reference
[1]	CeFCO3 = Ce3+ + F− + CO32−	KSO = 10−16.1	(this work)
[2]	HCO3− + H2O = H2CO3 + OH−	K1 = 10−7.65	15)
[3]	CO32− + H2O = HCO3− + OH−	K2 = 10−3.67	15)
[4]	CO2 + H2O = H2CO3	KH = 10−1.47	15)
[5]	Ce3+ + OH− = CeOH2+	K3 = 105.9	10,16)
[6]	Ce3+ + 2OH− = Ce(OH)2+	K4 = 10−2.3	10,16)
[7]	Ce3+ + 3OH− = Ce(OH)3o[aq]	K5 = 10−12.0	10,16)
[8]	Ce3+ + 3OH− = Ce(OH)3[s]	K6 = 1019.7	17,18)
[9]	Ce(OH)3o[aq] = Ce(OH)3[s]	K7 = 1031.8	*
[10]	Ce3+ + 4OH− = Ce(OH)4−	K8 = 10−24.0	10,16)
[11]	3Ce3+ + 4OH− = Ce3(OH)54+	K9 = 10−18.8	10,16)
[12]	Ce3+ + HCO3− = CeHCO32+	K10 = 103.4	(this work)
[13]	Ce3+ + F− = CeF2+	K11 = 103.99	13)
[14]	F− +H+ = HF	K12 = 102.91	13)
[15]	2F− +H+ = HF2−	K13 = 103.48	13)

*  K₇ is calculated from Reactions [7] and [9].

Determination of K_so for Synthetic Ce-Bastnaesite

The solubility product of Ce-bastnaesite is defined as follows:   

$${\text{CeFCO}}_3={\text{Ce}}^{3+}+{\text{F}}^{-}+{{\text{CO}}_3}^{2-}$$(8)

  

$${\text{K}}_{\text{SO}}=\left({\text{Ce}}^{3+}\right)\left({\text{F}}^{-}\right)\left({{\text{CO}}_3}^{2-}\right)$$(9)

The total carbonate and fluoride concentrations in solution were measured with the corresponding ion selective electrodes and the cerium concentration was measured by a spectrophotometric technique¹⁴⁾. One ml of concentrated H₂SO₄ was added to a 10 ml sample solution and heated to fume off SO₃. Water was added and then heated again to thoroughly remove fluoride. The solution was placed into a 25 ml volumetric flask to which 2 ml of 0.1M citric acid and 1 ml of 7% H₂O₂ were added. Next, 6 M NaOH solution was added till the solution reached a slightly basic pH (when a yellow color appeared), The resulting solution was diluted to 25 ml volume and the resultant absorbance was measured with a Bausch and Lomb Spectrophotometer at 390 mm.

Fig. 3 shows the measured total equilibrium concentration of Ce³⁺, F⁻ and CO₃²⁻ as a function of pH at a constant partial pressure of 10^−3.5 atm. CO₂ in an open system. Based on these solubility data and the thermodynamic data available in the literature for reactions given in Table 3, the molar concentrations of various ionic species in solution were computed as a function of pH¹⁹⁾. The ionic strength of solutions under these conditions was calculated using these results and the corresponding activity coefficients for Ce³⁺, F⁻ and CO₃²⁻ ions were computed with the extended Debye-Huckel Law^17–18). Knowing the activity coefficients and the free ion concentrations, the solubility product of bastnaesite was estimated as a function of ionic strength as plotted in Fig. 4. The value of pK_SO for Ce-bastnaesite is thus 16.1, on extrapolation to zero ionic strength of the solution.

Fig. 3

The pH dependence of total cerium, total fluorine and total carbonate concentrations in aqueous solutions in equilibrium with air.

Fig. 4

The solubility product of synthetic cerium bastnaesite in aqueous solutions as a function of ionic strength.

Determination of the Stability Constant of CeHCO₃²⁺

Titration of H₂CO₃ in aqueous solution in the presence of Ce(NO₃)₃ was carried out in order to determine the stability constant of the complex cation CeHCO₃²⁺. For a pure aqueous, air-saturated solution, the pH is defined by the charge balance:   

$$\left[{\text{H}}^{+}\right]=\left[{{\text{HCO}}_3}^{-}\right]+2\left[{{\text{CO}}_3}^{2-}\right]+\left[{\text{OH}}^{-}\right].$$

If a metal cation which will form hydroxyl complexes with OH⁻ or ion-pair complexes with carbonates is added to the solution, the pH value will become lower since the complexing reactions will consume OH⁻. Thus by measuring the equilibrium pH values at different additions of Ce(NO₃)₃, the concentration of species containing cerium can be calculated from the measured pH and the cerium total concentration based on charge balance equations (14). From this, the stability constant of CeHCO₃²⁺ is simply   

$${\text{K}}_{\text{CeHC}}=\left[{{\text{CeHCO}}_3}^{2+}\right]/\left[{\text{Ce}}^{3+}\right]\left[{{\text{HCO}}_3}^{-}\right]$$(10)

For the titration experiment, a weighed amount of Ce(NO₃)₃ was introduced into a 100 ml solution of NaOH or HNO₃ of known concentration and the solution was thoroughly stirred under an open-system atmosphere until a stable pH reading was achieved. The pH was determined by a Fisher digital pH/mV meter. The value of K_CeHC obtained was 3.40 ± 0.21.

Using the stability constant thus determined with Equation 10 and the thermodynamic data given in Table 3, the relative concentrations of various solution spcies in equilibrium with synthetic cerium fluocarbonate in a system open to atmosphere at 25 °C, were computed. The results ace plotted in Fig. 5. The main charged species in a solution in equilibrium with synthetic cerium bastnaesite in a system open to atmosphere are Ce³⁺, CeF²⁺, HCO₃⁻, CO₃²⁻, F⁻, CeOH²⁺, Ce(OH)₂⁺ and CeHCO₃⁺ in addition to H⁺ and OH⁻. As shown in Fig. 3, the total species concentration is minimum in the pH range of 7.2 - 7.6.

Fig. 5

Equilibria in the synthetic cerium bastnaesite/water system open to the atmosphere at 25 °C.

Summary

A sample of cerium f1uocarbonate (CeFCO₃) was synthesized in the laboratory and characterized through X-ray diffraction and chemical analysis. Thermogravimetric studies on the synthetic bastnaesite exhibited a decomposition peak at 500 °C as compared to 600 °C in the case of natural bastnaesite obtained from Mountain Pass, California. The total weight loss when heated to 900 °C was 20 percent, consistent with the decomposition reaction for the mineral.

The solubility product K_SO as well as the stability constant K_CeHC were experimentally determined through solubility and titration measurements. Using the thermodynamic data available in the literature as well as the constants, a speciation diagram for cerium fluocarbonate has been presented.

Author’s short biography

Pradip

After graduating from the Indian Institute of Technology in Kanpur, Pradip entered the University of California at Berkeley to continue graduate studies in mineral processing. He obtained his M.S. degree in 1977 and Ph.D. degree in 1981. After serving as Scientific Officer, Ore Dressing Section, at the Bhabha Atomic Research Centre, he joined the Tata Research Design and Development Centre where he currently is Chief Scientist and Head of the Process Innovation Laboratory.

Charles C.H. Li

After graduating from Central South Institute of Mining and Metallurgy in Changsha, Charles Li entered the University of California at Berkeley to pursue graduate studies in mineral processing, and obtained his M.S. degree in 1982 and his Ph.D. degree in 1986. He then joined CRA (now Rio Tinto) in Australia, where he worked as a research engineer on mineral processing problems until his retirement.

Douglas W. Fuerstenau

After receiving his Sc.D. degree at MIT, Dr. Fuerstenau spent a six-year period teaching at MIT and working in industry, after which time he joined the faculty of the University of California. At Berkeley he established an extensive program of teaching and research in mineral processing, applied surface chemistry and particle technology. With his graduate students and post doctoral researchers, he has published a wide range of seminal papers in these fields. He currently is P. Malozemoff Professor Emeritus of Mineral Engineering in the University of California.

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