Study on the coordination properties of [CeFx]4-x and [BFy]3-y in sulfuric acid medium by complexometric potentiometric titration

Article abstract of DOI:10.1016/j.jfluchem.2013.09.007

Complexometric titration experiments were carried out in three systems of F-Ce, F-B and F-Ce-B to investigate the coordination properties of CeF _x^4-x and BF_y^3-y with ion-selective electrode as indicator. The str

Full text of DOI:10.1016/j.jfluchem.2013.09.007

Journal of Fluorine Chemistry 156 (2013) 106–111
Contents lists available at ScienceDirect
Journal of Fluorine Chemistry
journal homepage: www.elsevier.com/locate/fluor
Study on the coordination properties of [CeFx]4ꢀx and [BFy]3ꢀy in
sulfuric acid medium by complexometric potentiometric titration
a
b
a
Jingui He a, Yong Li a, Xiangxin Xue a, , Hongqiang Ru , Xiaowei Huang , He Yang
*
a School of Material and Metallurgy, Northeastern University, Shenyang, Liaoning 110819, China
b National Engineering Research Center for Rare Earth Materials, General Research Institute for Nonferrous Metals, and Grirem Advanced Materials Co., Ltd.,
Beijing 100088, China
A R T I C L E I N F O
A B S T R A C T
Article history:
Complexometric titration experiments were carried out in three systems of F–Ce, F–B and F–Ce–B to
4ꢀx
3ꢀy
Received 29 November 2012
Received in revised form 18 August 2013
Accepted 12 September 2013
Available online 20 September 2013
investigate the coordination properties of CeFx
and BFy
with ion-selective electrode as indicator.
The structure of the complex was characterized by Raman spectrometer. The results show that in sulfuric
acid medium, Fꢀ coordinates with Ce4+ to form CeF22+ complex, and B(OH)3 can take the Fꢀ from CeF22+ to
form BF4ꢀ. The BFy3ꢀy is more stable than CeFx4ꢀx. The KBF4 product was prepared and confirmed by XRD.
The results indicate that it is feasible to separate fluorine from cerium with B(OH)3 and recover fluorine
as KBF4 before extraction from fluorine-bearing rare earth sulfate solution.
Keywords:
2+
CeFꢀ2
ß 2013 Elsevier B.V. All rights reserved.
BF4
Coordination properties
Complexometric potentiometric titration
Ion-selective electrode
1. Introduction
solid. The discharge of fluorine-containing wastes will cause
serious fluoride pollution of surface and ground water and fluorine
Rare earth has been widely used in metallurgy, nuclear
industry, electronic industry, agriculture, medicine and other
fields due to its unique optical, electrical, magnetic properties.
There are about 200 rare earths (RE) minerals worldwide, including
halides, carbonates, oxides, phosphates, silicates, etc. [1]. Among
them, bastnaesite (ReCO3F, Re = Rare earths) is one of the most
important mineral resources containing about 75% combined rare
earth oxides (REO) especially La and Ce [2]. The main deposits of
bastnaesite resource in China distribute in Sichuan, Shandong and
Inner Mongolia provinces, while the Inner Mongolia Baotou region
is the most abundant one. There are several hydrometallurgical
methods for extracting rare earths from bastnaesite, such as
‘‘concentrated sulfuric acid roasting-water leaching’’, ‘‘oxidation
roasting-hydrochloric acid leaching’’, ‘‘oxidation roasting-sulfuric
acid leaching’’ [3–5]. The ‘‘oxidation roasting-hydrochloric acid
leaching’’ and ‘‘concentrated sulfuric acid roasting-water leaching’’
technologies have the disadvantages of equipments corrosion, long
process flow and discharge of fluorine via waste liquid or waste
resource waste. Recently, more attention has been paid to develop
a more environment-friendly technology of bastnaesite [4,6,7]. The
‘‘oxidation roasting-sulfuric acid leaching’’ process can dissolute
almost all of Ce4+, RE3+, Fꢀ and Th4+, so it is considered to be more
advanced technology for bastnaesite utilization [8–10]. Bastnae-
site theoretically contains about 7–9% fluorine. Previous investiga-
4ꢀx
tions [11–13] show that Fꢀ coordinates with Ce4+ to form CeFx
complex in sulfuric acid system because of the high electronega-
tivity and small ionic size, which makes it difficult to separate Fꢀ,
Ce4+ and RE3+, and brings about the formation of the third phase
during the subsequent extractive separation of rare earths.
Therefore, the key point of sulfuric acid leaching process is to
eliminate the influence of fluorine.
Presently, electrochemical method is widely used in the study
of coordination chemistry in solution due to the convenient
measurement operation, the simple equipment, the accurate result
and especially the continuous appearance of new ion-selective
electrodes [14]. Complexometric potentiometric titration is
frequently applied to the determination of ion complexation.
Ion selective membrane electrodes and mercury electrodes are
generally used as indicator electrodes in titration [15]. It is known
that B3+ can complex with Fꢀ strongly [16]. The ability of Fꢀ
extraction by B(OH)3 is well known and it has been utilized to
prepare various oxides [17–19]. The fluorinated compound KBF4 is
an important industrial chemical which is widely used in industry,
*
Corresponding author at: No. 11, Lane 3, Wenhua Road, Heping District,
Shenyang City, Liaoning Province 110819, China. Tel.: +86 24 83687371;
fax: +86 24 23906316.
E-mail addresses: hejingui_neu@126.com (J. He), Xue_Xiangxin@126.com
(X. Xue).
0022-1139/$ – see front matter ß 2013 Elsevier B.V. All rights reserved.
http://dx.doi.org/10.1016/j.jfluchem.2013.09.007

J. He et al. / Journal of Fluorine Chemistry 156 (2013) 106–111
107
and is oneof the waysof recovering and utilizing of fluorineresource
from bastnasite [20,21]. In this study, based on environmental
protection and fluorine resource utilization, the coordination
4ꢀx
3ꢀy
properties of CeFx
and BFy
in sulfuric acid medium were
investigated by complexometric potentiometric titration with ion-
selective electrode as indicator. The fluorine can be separated from
cerium and recovered as KBF4 before extractive separation of rare
earths. It is expected to eliminate the influence of fluorine on the
extractive separation of rare earths, and carry out exploration
research on industrial application of ‘‘oxidation roasting-sulfuric
acid leaching’’ technology. Recovery of fluorine in the form of KBF4
can also increase the economical value of bastnasite.
2. Results and discussion
2.1. Theory
In order to investigate the coordination property of CeFx4ꢀx and
BFy3ꢀy in sulfuric acid medium, a PF-1 fluoride ion electrode and a
PBF4-1 fluoroborate ion electrode were selected as the indicator
electrodes to record the potential values. Linear relationship exists
between the potential and logarithm of ion activity according to
Nernst equation:
ꢀ
ꢁ
2:303RT
E ¼ K ꢀ
lga
(1)
nF
When the total ion strength of the solution is constant and
sufficient, the activity coefficient is a fixed value, so the potential
values of ions satisfy the following formula:
ꢀ
ꢁ
2:303RT
E ¼ K0 ꢀ
lgC
(2)
nF
where, E is the potential value, C is the ion concentration, K and K0 are
constants, isthe activityF is the faraday constant, R isthe molar gas
a
constant, n is the electron transfer number, T is the temperature.
The potential and logarithm of ion concentration present good
linear relationship, so the variations of Fꢀ and BF4ꢀ concentrations
can be characterized by measuring the potential values with a
fluoride ion electrode and a fluoroborate ion electrode.
Fig. 1. Complexometric titration curves of F–Ce system.
2.3. F–B complexometric titration
2.2. F–Ce complexometric titration
F solution was titrated with B solution to investigate the
coordination property of fluoride and boron under the same acidity
as above. The results are shown in Fig. 2. These curves show the
To imitate and investigate the coordination property of fluoride
and cerium in fluorine-bearing rare earth sulfate solution, F
solution was titrated with Ce solution under the acidity of 1.0 mol/
l.
The complexometric titration curves are shown in Fig. 1. As
shown in Fig. 1(a), the potential of Fꢀ increases with the addition of
the titrant, showing that the Fꢀ concentration decreases according
to Eq. (2). A mutation is observed at the titrant volume of 1 ml, and
the increase of potential approaches equilibrium with further
titration. This is because that Fꢀ can complex with Ce4+ in sulfuric
acid medium, which can be expressed as:
increase in potential of Fꢀ and decrease in potential of BF4 with
ꢀ
the addition of B(OH)3 in solution, indicating the decrease of Fꢀ and
formation of BF4ꢀ ions. Besides, the change of potential of Fꢀ is fast
in the beginning, and becomes slow when the ꢀtitrant volume is
over 1 ml, but the variation of potential of BF4 is nearly stable
during the whole titration process. It is evident fromꢀFig. 2(b) that
the variation velocities of potentials of Fꢀ and BF4 are similar
when the nF/nB ratio is over 4, however, the variation of potꢀential of
Fꢀ slows down while the variation of potential of BF4 keeps
almost stable when the nF/nB ratio is below 4. To evaluate the
reaction rate, the kinetics experiment was carried out and the
Ce4 þ xFꢀ Ð CeF4xꢀx
(3)
results were obtained by constructing a plot of
DE/Dt against t as
Fig. 1(b) represents E against nF/nCe for the complexometric
titration. It can be seen that in the initial stage of titration, the
increase of potential value is unobvious at high nF/nCe ratio. With
further increase in Ce4+ concentration, the nF/nCe ratio changes
more and more slowly. A dramatically increase of potential value is
detected when the nF/nCe ratio is close to 2, and approaches
equilibrium when the nF/nCe ratio is over 2. It can be deduced that
the complex-ratio of Fꢀ and Ce4+ is close to 2. This result is
shown in Fig. 3. From the plots the decrease of Fꢀ and formation of
BF4ꢀ is fast in the initial 20 min, and keeps almost invariable after
40 min of reaction. Fig. 4 depicts the Raman spectra of F–B solution.
The pꢀeaks at 980 cmꢀ1, 1050 cmꢀ1 [23] are attributed to SO42ꢀ and
HSO4 species, respectively. The peak at about 876 cmꢀ1 [24] is
assigned to a total symmetric stretching vibration of BO3 unit in
B(OH)3. The peak at around 771 cmꢀ1 [25] is due to the BF4ꢀ bonds.
As seen in Fig. 4, the band at 876 cmꢀ1 is not observed in the Raman
spectrum of F–B solution at the early stage of reaction, while the
771 cmꢀ1 band begins to appear, indicating that B(OH)3 is quickly
2+
consistent with that Fꢀ coordinates with Ce4+ to form CeF2
complex at high acidity reported in literature [11–13,22].

108
J. He et al. / Journal of Fluorine Chemistry 156 (2013) 106–111
Fig. 4. Raman spectra of F–B solution.
the results mentioned above, possible reactions are suggested in
Eqs. (4)–(8) [29]:
ꢀ
BðOHÞ3 þ 4Fꢀ ! BF4
(4)
(5)
(6)
(7)
(8)
fast
BðOHÞ3 þ Fꢀꢀ!BFðOHÞꢀ3
fast
BFðOHÞꢀ3 þ Fꢀ þ Hþꢀ!BF2ðOHÞ2ꢀ þ H2O
fast
BF2ðOHÞꢀ2 þ Fꢀ þ Hþꢀ!BF3OHꢀ þ H2O
Fig. 2. Complexometric titration curves of F–B system.
slow
BF3OHꢀ þ Fꢀ þ Hþꢀ!BFꢀ4 þ H2O
complex with Fꢀ to form BF4 ions in solution. After 40 min of
reaction, the 876 cmꢀ1 band is observed and the relative intensity
of the band at 771 cmꢀ1 has no evident change. The bands at
771 cmꢀ1 and 876 cmꢀ1 both increase with the proceed of
reaction. These results firmly support the results of complexo-
metric titration.
In the present study, it has been reported that fluoroborate ions
(BFn(OH)4ꢀnꢀ, 1 ꢁ n ꢁ 4) are formed in H3BO3–Fꢀ–H+ system [26–
28]. Wamser [26,27] has shown by conductance measurements
that the main boron-containing species in H3BO3–Fꢀ–H+ solution
are: HBF4, HBF3OH, HBF2(OH)2, HBFOH3, and H3BO3. According to
ꢀ
Grassino and Hume [28] have reported that the formation of
BF3OHꢀ is very rapid, but it is slow to form BF4ꢀ from BF3OHꢀ. In
the initial stage of reaction, the Fꢀ concentration is high enough
that Eq. (4) is the dominant reaction, resulting in a large number of
BF4 ions in solution. However, with the decrease of the Fꢀ
ꢀ
concentration in solution, 3Fꢀ reacts quickly with B(OH)3 to form
BF3OHꢀ, which leads to large amounts of BF3OHꢀ ions in solution,
thereafter the BF3OHꢀ ions will complex with residual Fꢀ to form
ꢀ
BF4
.
2.4. F–Ce–B complexometric titration
In order to separate fluorine from cerium and eliminate the
influence of fluorine on extractive separation of rare earths. The F–
Ce mixed solution was titrated with B(OH)3 solution to investigate
4ꢀx
3ꢀy
the coordination property of CeFx
and BFy
under the same
acidity as above. The complexometric titration curves are shown in
Fig. 5.
It is found from Fig. 5 that the same tendencies of Fꢀ and BF4ꢀ as
ꢀ
Fig. 2 are observed, showing that BF4 ions are generated in F–Ce
solution with the increase of H3BO3 concentration in solution. The
variation range of potential of Fꢀ is much smaller than that of Fig. 2.
The reason is the formation of CeF22+ complex ions in sulfuric acid
medium and few free Fꢀ is leꢀft in solution. Meanwhile, the decrease
velocity of potential of BF4 in F–Ce–B solution is nearly stable
during the titration. Fig. 6 represents the relation between the
reaction rate and response time for F–Ce–B solution. It is found that
t values of Fꢀ and BF4 are much smaller than that of
ꢀ
the
D
E/
D
Fig. 3, showing that the formation rate of BF4ꢀ in F–Ce–B solution is
much slower than F–B solution. The formation of BF4ꢀ is fast in the
Fig. 3. The relation between the reaction rate and response time.

J. He et al. / Journal of Fluorine Chemistry 156 (2013) 106–111
109
Fig. 7. Raman spectra of F–Ce–B solution.
reaction of B(OH)3 and Fꢀ in F–Ce–B solution can be proposed
(Eqs. (9)–(11)):
BðOHÞ3 þ 2CeF22þ ! BF4ꢀ þ 2Ce4þ
(9)
(10)
(11)
fast
2BðOHÞ3 þ 3CeF22þꢀ!2BF3OHꢀ þ 3Ce4þ
slow
2BF3OHꢀ þ CeF22þꢀ!2BFꢀ4 þ Ce4þ
Because Fꢀ and Ce4+ exist in the form of CeF22+ in solution, when
F–Ce mixed solution was titrated with B(OH)3 solution, B(OH)3 can
take the Fꢀ from CeF22+ to form BF4ꢀ. In the initial stage of reaction,
Fig. 5. Complexometric titration curves of F–Ce–B system.
BF4 and BF3OHꢀ ions are produced simultaneously in solution.
ꢀ
With the increase of the reaction time, BF3OHꢀ continue to react
first 100 min, and slows down with the proceed of reaction. Fig. 7
gives the Raman spectra of F–Ce–B solution. Because of the high
Ce(SO4)2 concentration in the solution, the very strong and clear
Raman spectra peaks at 980 cmꢀ1 and 1050 cmꢀ1 are observed,
and the peak at about 680 cmꢀ1 [30] is assigned to in-plane
with Fꢀ from CeF22+ to form BF4ꢀ. These results suggest that BFy
3ꢀy
4ꢀx
is more stable than CeFx
in sulfuric acid medium.
2.5. Preparation of KBF4 from fluorine-bearing rare earth sulfate
solution
2ꢀ
bending vibration of SO4 species. In the early stage of reaction,
the bands at 771 cmꢀ1 and 876 cmꢀ1 are observed and the band
intensity of the 771 cmꢀ1 is quite low in comparison with F–B
solution, indicating the low formation velocity of BF4ꢀ. Thereafter,
the bands at 771 cmꢀ1 and 876 cmꢀ1 both increase with further
reaction. Thus, considering the results obtained from complexo-
metric titration and Raman spectra, a mechanism for the complex
To verify the results obtained by complexometric potentiomet-
ric titration and Raman spectra, Ce(SO4)2 and KF were used as raw
Fig. 6. The relation between the reaction rate and response time.
Fig. 8. The XRD pattern of prepared product.

110
J. He et al. / Journal of Fluorine Chemistry 156 (2013) 106–111
materials to simulate the fluorine-bearing rare earth sulfate
solution with the nF/nCe = 2, and stoichiometric ratio of B(OH)3
and KCl were added into the solution. Because of the micro-
solubility of KBF4 in water, it is easy to precipitate from fluorine-
bearing rare earth sulfate solution. Fig. 8 gives the XRD pattern of
prepared product, and the product is confirmed as KBF4. The result
suggests that it is feasible to recover fluorine in the form of KBF4
before the extractive separation of rare earths from fluorine-
bearing rare earth sulfate solution.
Titration of F–B: 2 ml of F solution was transferred into a 250 ml
vessel and diluted to 100 ml. The acidity was adjusted to 1.0 mol/l
with sulfuric acid. The solution was then stirred and titrated with
0.5 mol/l B solution, and the titration rate of B solution was 0.1 ml/
min.
Titration of F–Ce–B: In order to obtain F–Ce mixed solution with
nF/nCe ratio of 2, 2 ml of F solution and 1 ml of Ce solution were
transferred into a 250 ml vessel and diluted to 100 ml. The acidity
was adjusted to 1.0 mol/l with sulfuric acid. The mixture was then
stirred and titrated with 0.5 mol/l B solution, and the titration rate
of B solution was 0.2 ml/10 min.
3. Conclusion
A saturated calomel electrode was used as reference electrode,
and a PF-1 fluoride ion electrode and a PBF4-1 fluoroborate ion
electrode were used as indicator electrodes to record potential
values (E) with the increase of titrant volume (V). Then, the
titration curves were drawn to analyze the concentration changes
4ꢀx
3ꢀy
were
The coordination properties of CeFx
and BFy
investigated by complexometric potentiometric titration with
ion-selective electrode as indicator. The KBF4 was prepared from
fluorine-bearing rare earth sulfate solution. The conclusions drawn
from the studies can be summarized as:
of Fꢀ and BF4
.
ꢀ
(1) Under the acidity of 1.0 mol/l, Fꢀ coordinates with Ce4+ to form
4.3.2. Kinetics experiment
2+
CeF2 complex, and combines with B(OH)3 to form stable
F–B system. 2 ml of F solution was transferred into a 250 ml
vessel and diluted to 99.5 ml with the acidity of 1.0 mol/l. The
solution was mixed on a magnetic stirrer, and the electrodes were
placed into solution. 0.5 ml of B solution was added into solution
quickly, and began timing immediately. The potential values (E)
were recorded with the increase of time (t).
ꢀ
complex ions of BF4
.
(2) In F–Ce–B mixed solution, B(OH)3 can take the Fꢀ from CeF2
2+
to form BF4ꢀ, showing that BFy3ꢀy is more stable than CeFx4ꢀx in
sulfuric acid medium. The prepared product is confirmed to be
KBF4 by XRD analysis.
(3) The results indicate that it is feasible to separate fluorine from
cerium with B(OH)3 and recover fluorine as KBF4 before the
extractive from fluorine-bearing rare earth sulfate solution.
F–Ce–B system. 2 ml of F solution and 1 ml of Ce solution were
transferred into a 250 ml vessel and diluted to 99.5 ml with the
acidity of 1.0 mol/l. The solution was mixed on a magnetic stirrer,
and the electrodes were placed into solution. 0.5 ml of B solution
was added into solution quickly, and began timing immediately.
The potential values (E) were recorded with the increase of time (t).
4. Experimental
4.1. Apparatus
4.3.3. Preparation of KBF4
Ce(SO4)2 and KF were used as raw materials to simulate
fluorine-bearing rare earth sulfate solution, and the acidity was
adjusted to 1.0 mol/l. The simulated solution was mixed with
stoichiometric ratio of B(OH)3 and KCl, then stirred magnetically
for 5 h at room temperature. After reaction, the product was
filtered in vaccum, washed three times with deionized water and
dried 2 h at 80 8C–100 8C, then ground for detection.
Titrations were carried out by a Leici ZD-2 automatic potentio-
metric titrator (Shanghai Precision Scientific Instrument Co., Ltd.).
Thestirrer, theburet and the glass vessels werestandard equipment.
A PF-1 fluoride ion electrode and PBF4-1 fluoroborate ion electrode
in conjunction with saturated calomel electrodes were used to
determine the electric potential of Fꢀ and BF4ꢀ. HR800UV laser
Raman spectrometer (Horibajobin Yvon) was used to record the
Raman spectra of aqueous solution in the range of 600–1200 cmꢀ1
using the excitation wavelength of 633 nm. Ultima IV X-ray
diffraction (Rigaku) was used to examine the phase structure in
Acknowledgements
The financial aids from the following programs are gratefully
acknowledged. They are the key program of National Natural
Science Foundation of China (NSFC: 50934004), National Natural
Science Foundation of China (51274061), Major State Basic
Research Development Program of China (973 Program:
2012CBA01205), Fundamental Research Supporting Project of
Northeastern University (N110602006).
the 2
u a radiation.
of 10–908 at a rate of 88 minꢀ1 with Cu K
4.2. Reagents
All reagents used of analytical grade were purchased from
Shenyang Guoyao Group Chemical Reagent Co., Ltd. All of the
chemicals were diluted with deionized water.
F solution, 0.5 mol/l, was prepared by dissolving 2.0995 g of
sodium fluoride in distilled water and diluted to 100 ml. B solution,
0.5 mol/l, was prepared by dissolving 3.0915 g of boric acid in
distilled water and diluted to 100 ml. Ce solution, 0.5 mol/l, was
prepared by dissolving 20.215 g of cerium sulfate tetrahydrate in
dilute sulfuric acid and diluted to 100 ml.
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