Prediction of Solubilities of Aluminium Hydroxyfluoride
Hydrate Recovered from Spent Pot Lining
Ubong Ntuk*, Edward White, Stephan Tait and Karen Steel
School of Chemical Engineering, The University of Queensland, St Lucia, 4072, Queensland, Australia.
*Email: [email protected]
Abstract. Spent pot lining (SPL) is a hazardous waste
containing fluoride, generated in the production of
primary aluminium. About 7 to 30 kg of SPL is generated
per tonne of primary aluminium. Over the years, much
research has been carried out to offer an economic method
of treatng this waste, but to date no such method is widely
accepted. Chemical leaching and recovery of valuables
have been found to be a promising option. A previous
study has successfully recovered the fluoride values by
crystallizing aluminium hydroxyfluoride hydrate (AHF).
However the crystal size obtained was too small for
potential industrial interest. This work is focused on means
of producing larger crystals. A vital property for
crystallization is solubility, and there have not been
reliable solubility measures for AHF. Knowing the
solubility product (Ksp), the solubility at various conditions
may be predicted by a speciation program with an
appropriate activity model. In this work, experimental
measurements determined pKsp as 23.5 ± 0.3 (95%
probable error). Using this pKsp value in the speciation
program PhreeqC, solubilities at various conditions were
predicted. The solubility in water was found to be very
low, but much higher in acid (with a negative effect of
increasing temperature). The effect of an impurity on the
solubility is also presented. The predicted phase diagrams
suggest potentially viable crystallization pathways for
further testing through follow-up experimental work.
Keywords Spent pot lining, aluminium production, chemical
leaching, speciation program, crystallization, aluminium
hydroxyfluoride.
I.
INTRODUCTION
Aluminium metal is smelted in carbon lined electrolytic
cells (pots). After time (typically between 3-8 years) the
carbon lining degrades and is replaced. The disposal of the
spent pot linings (SPL) which contain considerable fluoride
(and also amounts of cyanide) is a problem for the aluminium
smelting industry. About 7 to 30 kg of SPL is generated per
tonne of primary aluminium produced.
Over the years, various attempts have been made to treat
SPL. But to date, there is no widely accepted commercial
method of treatment. Available technologies can be grouped
into three categories; inertisation, co-processing and material
recycling [1]. Inertisation processes makes SPL a benign
waste suitable for disposal to the environment while coprocessing utilises components of SPL in industries other than
the aluminium industry. Neither of these two methods fully
recovers the potentially valuable materials in SPL. Coprocessing (predominantly in cement, steel and brick
manufacture) uses mainly the carbon content of SPL, mostly
wasting the other valuable components. Material recycling
attempts to recover valuable products of SPL while generating
by-products that are useful or environmentally friendly.
Processes can be sub-classified into physical separation,
chemical leaching or thermal treatment [2].
Physical
separation properties include density, water affinity and
particle size. Chemical approaches selectively dissolve the
components of SPL using discriminating solvents, with
subsequent reaction and precipitation. The thermal approach
mostly utilises the fuel value of the graphite component.
Processes that have had considerable development [1] are
the pyro-metallurgical Ausmelt process [3], co-processing in
cement plants [4], [5] and Alcan’s leaching process to produce
CaF2 [6].
Previous research [1], [7] suggested that an acid leaching
process could be developed to produce crystalline aluminium
hydroxyfluoride (AHF, AlF2OH) as the hydrate (hAHF,
AlF2OH.H2O) [8]. AHF can readily be processed to produce
AlF3 which can be recycled to the smelting pots and the
leached carbon residues can be recycled. Lisbona [1]
undertook a considerable number of laboratory trials which
produced AHF, but the crystal size was too small for industrial
use.
The present work is to determine crystallization
conditions which will allow the production of larger crystals.
II. ALUMINIUM HYDROXYFLUORIDE HYDRATE
Where necessary for calculation and in the absence of a
definitive formula, AHF has been taken as AlF2(OH)·H2O.
The properties of AHF have not been significantly
investigated. There is uncertainty over its molecular formula
and its water of hydration. The major property reported is its
x-ray diffraction XRD pattern [9].
Aluminium hydroxyfluoride (AHF) has been described as a
ralstonite-like mineral [general formula NaxMgxAl2-xF6The Na and Mg free end member,
y(OH)y.nH2O].
Al16(F,OH)48.12-15H2O, occurs in nature and has been
observed in furamolic deposits, in
alkali granites, in
carbonatite veins as well as in Ivigtut cryolite deposit in
Greenland [10, 11].
With regard to its preparation and chemical formula,
compounds have been synthesized by several methods. From
an ionic balance the non-hydrated formula must be AlF3x(OH)x. Cowley and Scott [12] precipitated AHF from mixed
solutions of aluminium fluoride and sulphate treated with
ammonia at 100oC. They suggested a formula ranging from
x=1 to x=2 [AlF(OH)2 to AlF2(OH)]. Roberson and Hem [13]
obtained a compound from solutions of aluminium
perchlorate, NaF and NaOH that had a similar XRD pattern to
natural ralstonite. As there was no Mg in the composing
materials, it was thought to be the Mg free ralstonite end
member, Al (F,OH)3. Grobelny [14] found AHF to be one of
the phases present while recrystallizing AlF3 in water solution,
at a temperature range of 100-150oC. Its formula was
measured as AlF2.5(OH)0.5. A later study of the phase relations
in the system AlF3–Al2O3–H2O–HF, investigated between
400° and 700°C, suggested a range of composition for AHF
and
which included, AlF1.33(OH)1.67, AlF1.1(OH)1.9,
AlF1.7(OH)1.3. The AHF in aluminium hydroxyfluoride
hydrate (hAHF) was found to possess the empirical formula,
AlF3-x(OH)x [11]. According to this author, AHF and hAHF
are chemically similar, but structurally different.
The preparation method [15], said to yield reproducible
phases with the composition AlF2.3(OH)0.7.H20 was followed
by Konig et al. [16] but a different composition of
AlF1.9(OH)1.1.H20 was reported. Their work led to crystalline
aluminium hydroxyfluorides with tunable Al:F ratios as
AlFx(OH)3-x.nH2O. Lisbona and Steel [8] reported the
composition of their crystals as AlF2(OH)·1.4H2O. The XRD
reference contained in the International Centre for Diffraction
Data library, reports AlF3-x(OH)x.nH20 and Al (OH)1.38F1.62.
Though the varying compositions by different authors may
be due to the different synthesis methods, because of the
closeness in size of the F and OH units, it is possible that this
compound crystallises with varying F:OH combinations.
However the ratio of Al: (F,OH) is fixed at 1:3. What changes
is the ratio of F:OH and this work will use the more general
formula Al(F,OH)3 for aluminium hydroxyfluoride. The water
of crystallization will be taken as 1 mole per mole of Al.
III. SOLUBILITY MEASUREMENTS ON AHF
Very little has been reported on the solubility of AHF. This
is perhaps due to its limited application in commercial
chemical processes. Solubility data is essential for the design
of an operating crystallizer system. The solubility of a
substance determines the potential extent of recovery.
Chemical leaching and crystallization techniques can be
used to separate AHF from the SPL mass. While the ultimate
goal would be the production of AlF3, unfortunately AlF3 has
a relatively high solubility in aqueous solution [8]. It appears
that AHF can easily be converted to AlF3, which is used
constantly throughout the life of a typical aluminium smelter.
AHF was prepared by saturating 0.6M sulphuric acid with
AR Al(OH)3 at 25oC. After filtering, the Al2(SO4)3 solution
was saturated with NaF powder and filtered. The pH was
between 2-3. AHF was crystallized by the addition of 2M
NaOH to a pH of 4 - 5.5. Other species such as cryolite
(NaAl3F6) [13] and gibbsite (Al(OH)3) could also precipitate if
the pH further increases. The XRD was measured for the
washed precipitate and was in excellent agreement with the
published results [9], used as the reference AHF in the ICDD
library. The aluminium content of the solid was determined by
x-ray fluorescence XRF and in solution by inductively
coupled plasma atomic emission spectroscopy ICP-AES.
Fluoride contents were determined with an ion selective
electrode following a separation method proposed by [17].
Solubilities were measured by adding excess of the dried
AHF prepared above to water or diluted sulphuric acid in
stirred Schott bottles in a constant temperature bath. The test
solutions were allowed to equilibrate for 36 h, after which a
sample of the supernatant was taken for analysis. The
resulting values were used with the speciation program
PhreeqC to evaluate the solubility product Ksp.
Ksp was defined for the dissociation of AHF into its basic
ions according to the following assumed molecular formula,
with
AlF2OH ⇔ Al3+ + 2 F- + OHKsp = { Al3+}{ F-}2 { OH-}
= [ Al3+][ F-]2 [ OH- ] γAl3+ γF-2 γOH-
where curly brackets { } indicates the molal activity of the
ionic species in solution, square brackets [ ] the concentration
and γ the activity coefficient (to account for the effect of ionic
strength). PhreeqC was used with a modified Davies equation
to estimate activity coefficients. Figure 1 shows the results
obtained for AHF dissolved in water at 25 and 30oC. The
solubility of AHF is extremely low in water (~ 100 ppm).
From the available data an overall value of pKsp of 23.5 has
been chosen for the predictive calculations presented below.
IV. SOLUBILITY SIMULATION RESULTS
Figure 1. pKsp results for AHF in water. [p = - log10].
At 25oC in water, the pH should have a single value. The
variation in pH in Figure 1 was thought to result from minor
amounts of residual acid in the dried AHF. pKsp could be
higher at 30oC than for 25oC, but there are insufficient
observations at 25oC to fully support this, so the average of all
the results (pKsp = 23.5 ± 0.3) is shown.
Libona [1] in his development of the AHF recovery method
carried out a number of solubility measurements of AHF in
impure industrial liquors from 30 to 90oC. The pKsp value
derived using PhreeqC from his results are shown in Figure 2.
His solubilities (as anhydrous AHF) range up to 20 g/l ( ~ 2%
w/w) which is in a likely industrial range. His results indicate
no variation with temperature and an average value of pKsp =
23.5 ± 0.2. The values with solubilities below 3 g/l (dotted
vertical line) resulted from analyses with very poor accuracy
and thus have been ignored in finding the average pKsp. The
ionic strength of these solutions ranged up to 6 M, so they are
quite concentrated. That these values using PhreeqC were
similar to those in water suggests the activity coefficient
prediction equation is relatively sound.
The solubility results for AHF in water, predicted using
PhreeqC are shown in Figure 3 as a function of temperature.
As already indicated the solubilities are very low (AHF is
practically insoluble in water). The estimated pH’s are also
shown. The predicted pH at 25oC is larger than those shown
in Figure 1.
A minimum in solubility is predicted but the effect really is
small (change less than 10%) so it is of little interest.
The predicted solubilities in suphuric acid are substantial at
lower pH (Figure 4). Note that the solubility is shown on a log
scale. The corresponding plot on a linear scale is shown on
Figure 5. The uppermost row of points corresponds to a
sulphuric acid molality of 1.0. The subsequent lower rows
correspond to molalities of 0.5, 0.2, 0.1, 0.05, 0.02, 0.01,
0.005, 0.002, 0.001, 0.0005, 0.0002, 0.0001 and 0. Thus the
lower most row of points is the solubility in water (as in
Figure 3).
Figure 3. Predicted AHF solubility in water.
Figure 2. pKsp results for AHF in industrial liquors (Lisbona,
[1].
reduce the concentration of F-, but it did not because the
addition of aluminium sulphate decreases the pH
consuiderably (by 1.2 pH units between the end points shown
on Figure 8) and the decrease in OH- cause extra F- to be
diussolved.
The addition of a non-common ion impurity (e.g. SO42-) on
the other hand has little effefct on the solubility as shown in
Figure 9 for the addition of sodium sulphate at the 1 molal
level.
Figure 4. Predicted AHF solubility in aqueous sulphuric acid
vs. pH. Note log scale for solubility.
Predictions using hydrochloric acid (HCl) instead of
sulphuric acid (H2SO4) gave the same results as Figures 4 and
5 except that twice the number of moles of HCl than for
H2SO4 were required for the same effect. This suggests that
the acid is primarily a source of protons (H+ ions).
Figure 6. Predicted AHF solubility in aqueous sulphuric acid
vs. temperature. Note log scale for solubility.
Figure 5. Predicted AHF solubility in aqueous sulphuric acid
vs. pH with solubility on a linear scale.
Solubility shows an inverse temperature behavior. The
highest of the curves on Figures 4 and 5 are for 25oC.
Increasing temperature decreases the solubility as shown in
Figure 6.
Of importance, at higher pH, gibbsite rather than AHF is
the stable phase. Figure 7 shows the predicted transition
conditions. This transition is also shown on Figure 4 together
with the corresponding curve for amorphous Al(OH)3, which
from PhreeqC, has a different solubility from crystalline
gibbsite Al(OH)3.
3+
Adding a common ion (e.g. Al ) would be expected to
considerably alter the solubility. As shown in Figure 8 the
addition of aluminium sulphate increases the solubility in
water (no acid). Adding extra Al3+ ion would be expected to
Figure 7. Predicted transition conditions for AHF – gibbsite.
the crystallization kinetics and these will be determined in
later studies.
V.
VALIDATION
As with all predicted results, experimental validation is
required. The predicted results for water using the selected
Ksp are in reasonable agreement with the measured results
given in Figure 1.
Experimental studies are under way to measure the
solubility of AHF in aqueous sulphuric acid. Until then the
predicted results need to be used with caution.
VI. CONCLUSIONS
Figure 8. Effect of the addition of aluminium sulphate to
water saturated with AHF.
1. From the solubility in water experiments and the prior
studies on the crystallization of AHF from industrial
strength liquors the solubility product of AHF has been
estimated as pKsp = 23.5 ± 0.3 (95% probable error).
2. This estimated Ksp was used with the speciation program
PhreeqC to predict the solubilities of AHF. These
solubility values still remain to be validated and can only
be considered as indicative at present.
3. AHF is practically insoluble in water.The solubility is of
the order of 0.01 wt %
4. AHF is very soluble in aqueous sulphuric acids and the
change of solubility with pH and temperature would be the
basis of a crystallization process .
5. Gibbsite could also crystallize if the pH becomes too high
(> 4 – 5).
Figure 9. Addition of sodium sulphate (1 molal) on solubility.
The solubility plots may be used to suggest possible
industrial crystallization pathways. From Figure 5 at 25oC a
10% sulphuric acid solution (~ 1m) could be saturated with
AHF (pH ~ 3) to give an AHF solubility of ~ 20 kg AHF per
100 g water (~ 2 moles AHF/kg water). When neutralised to a
pH of ~ 4.3 (keeping above the gibbsite transition condition of
Figure 7) the solubility is ~ 2 kg AHF per 100 kg of water (~
0.2 moles AHF/kg water) giving a good AHF recovery. If
operated to always be at equilibrium, the crystallization path
will follow the 25oC isotherm in Figure 5. In practice with a
finite supersaturation, the curve would be to the right of the
isotherm (at positions governed by the neutralisation routine)..
Alternatively the 25oC solution could be heated to 90oC to
obtain the reduced solubility. Crystallization reduces the pH so
the path on Figure 5 would need to be calculated once cooling
cconditions are selected. The process choice will depend on
6. Further predictions are required to determine the effect of
excess (above the stoichiometric ratio) fluoride and
aluminium. Aluminium ions have been shown to have a
marked effect on the solubility of AHF in water.
7. The solubility appears to be relatively insensitive to neutral
non-common ion impurities.
8. Hydrochloric acid gives the same solubility for the same
concentration of supplied hydrogen ions.
9. The solubility charts allow tentative industrial processes to
be considered.
VII. ACKNOWLEDGEMENTS
The authors would like to thank the Australian Research
Council and Fluorsid SpA for providing an Australian Postgraduate Award - Industry (APAI) for Ubong Ntuk through
the ARC-Linkage funding scheme.
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