Magnesium removal in the
electrolytic zinc industry
Proefschrift
ter verkrijging van de graad van doctor
aan de Technische Universiteit Delft
op gezag van de Rector Magnificus prof.dr.ir. J.T. Fokkema
voorzitter van het College voor Promoties
in het openbaar te verdedigen op maandag 12 mei 2003 om 13.30 uur
door Jacco Leendert BOOSTER
doctorandus in de geochemie en in de milieukunde
geboren te Maassluis
Dit proefschrift is goedgekeurd door de promotor:
Prof. dr. M.A. Reuter
en de toegevoegd promotor:
Dr. A. van Sandwijk
Samenstelling promotiecommissie:
Rector Magnificus, voorzitter
Prof. dr. M.A. Reuter
Dr. A. van Sandwijk
Prof. dr. G.J. Witkamp
Prof. dr.ir. G. van Weert
Prof. dr. A.J. Monhemius
Prof. dr. J.S.J. van Deventer
Ir. C.F.M. Copini
Technische Universiteit Delft, promotor
Technische Universiteit Delft, toegevoegd promotor
Technische Universiteit Delft
Oretome Ltd., Canada
Imperial College, London, United Kingdom
University of Melbourne, Australia
Pasminco Budel Zink, adviseur
The project was supported by a Dutch E.E.T. grant (Economy, Ecology and
Technology) from the Ministry of Education, Culture and Sciences and the Ministry
of Housing, Spatial Planning and the Environment.
ISBN 90-9016857-5
Copyright © 2003 by J.L. Booster
All rights reserved
Printed in The Netherlands
CONTENTS
CONTENTS....................................................................................................I
LIST OF TABLES ....................................................................................... V
LIST OF FIGURES .................................................................................. VII
LIST OF SYMBOLS ..................................................................................XI
1
OVERVIEW OF ZINC METALLURGY .......................................... 1
1.1
1.2
1.3
1.4
1.5
Introduction.................................................................................................................1
Zinc metallurgy...........................................................................................................6
Hydrometallurgical zinc production ...........................................................................9
1.3.1 Roasting ..........................................................................................................9
1.3.2 Neutral leach .................................................................................................10
1.3.3 Purification....................................................................................................11
1.3.4 Electrolysis....................................................................................................11
1.3.5 Hot acid leach ...............................................................................................12
1.3.6 Jarosite precipitation .....................................................................................12
1.3.7 Goethite precipitation....................................................................................14
1.3.8 Hematite precipitation...................................................................................16
Magnesium control in electrolytic zinc plants ..........................................................17
Pasminco Budel Zink – residue free zinc production ...............................................20
2
AN ALTERNATIVE MAGNESIUM BLEED PROCESS.............. 25
2.1
2.2
Introduction...............................................................................................................25
Experimental .............................................................................................................27
2.2.1 Precipitation ..................................................................................................27
2.2.2 Conversion ....................................................................................................28
2.2.3 Electrolysis....................................................................................................28
Results29
2.3.1 Precipitation ..................................................................................................29
2.3.2 Conversion ....................................................................................................31
2.3.3 Electrolysis....................................................................................................32
Discussion .................................................................................................................33
2.4.1 Precipitation ..................................................................................................33
2.4.2 Conversion ....................................................................................................34
2.4.3 Electrolysis....................................................................................................34
2.4.4 Proposed flowsheet .......................................................................................35
Conclusions...............................................................................................................35
2.3
2.4
2.5
3
SELECTIVE PRECIPITATION OF A MAGNESIUM
FLUORIDE PRODUCT ..................................................................... 37
3.1
3.2
Introduction to precipitation processes ....................................................................37
Thermodynamic modelling of magnesium fluoride precipitation in concentrated
zinc sulphate environment ........................................................................................40
3.2.1 Formulation of the thermodynamic model ...................................................40
3.2.2 Determination of mass stability constants ....................................................43
3.2.3 Results and discussion of the thermodynamic model ...................................47
3.2.4 Conclusions drawn from the thermodynamic model ....................................52
Treatment of fluoride containing industrial zinc sulphate solutions.........................53
3.3.1 Literature review...........................................................................................53
3.3.2 Experimental verification..............................................................................55
Characterisation of hydroxyl-bearing magnesium fluoride containing physically
bound water...............................................................................................................58
3.4.1 Introduction...................................................................................................58
3.4.2 Experimental .................................................................................................60
3.4.3 Results and Discussion .................................................................................61
3.4.4 Conclusions...................................................................................................69
3.3
3.4
4
CONVERSION OF MAGNESIUM FLUORIDE TO
MAGNESIUM HYDROXIDE ........................................................... 71
4.1
4.2
4.3
Introduction...............................................................................................................71
Experimental .............................................................................................................73
Results.......................................................................................................................76
4.3.1 Reaction temperature ....................................................................................76
4.3.2 Stirring velocity ............................................................................................77
4.3.3 Washing intensity..........................................................................................77
4.3.4 Holding time .................................................................................................78
4.3.5 Leach concentration ......................................................................................78
4.3.6 Phase characterisation...................................................................................80
4.3.7 Other impurities ............................................................................................81
Thermodynamics of the conversion reaction............................................................82
Thermal decomposition of brucite to periclase.........................................................90
4.4
4.5
5
REGENERATION OF CHEMICALS IN A MEMBRANE
CELL REACTOR ............................................................................... 97
5.1
5.2
Theory.......................................................................................................................97
Regeneration of zinc fluoride and sodium hydroxide in an electrodialysis
process.....................................................................................................................101
Cell design ..............................................................................................................103
Experimental ...........................................................................................................104
5.4.1 Experimental set up of the first series.........................................................104
5.4.2 Experimental set up of the second series ....................................................104
5.4.3 Experimental set up of the third series........................................................105
5.4.4 Experimental set-up of the long-term experiment ......................................106
Results107
Discussion of the experimental results ...................................................................111
5.6.1 Transference numbers .................................................................................111
5.6.2 Water transport............................................................................................114
5.3
5.4
5.5
5.6
II
5.7
5.8
5.6.3 Active zinc anode........................................................................................115
5.6.4 Scaling.........................................................................................................116
5.6.5 Maximum current density ...........................................................................117
5.6.6 Long-term experiment ................................................................................117
Membrane performance in a larger scale experiment.............................................124
Conclusions.............................................................................................................131
6
CONCLUSIONS AND PROCESS INTEGRATION .................... 133
6.1
6.2
6.3
6.4
6.5
6.6
Chapter 1.................................................................................................................133
Chapter 2 and 3 .......................................................................................................133
Chapter 2 and 4 .......................................................................................................134
Chapter 2 and 5 .......................................................................................................135
Process Integration..................................................................................................136
Recommendations...................................................................................................139
APPENDIX A ............................................................................................ 143
REFERENCES.......................................................................................... 147
III
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LIST OF SYMBOLS
Chapter 3
Symbol
Dimension
Description
a
a±,m
a1
a2
ci
k
F
m
Mi
Nav
Q
r
R
T
zi
γ
γ±,m
γs
ν
ν’
ρf
mol l-1
1.38066*10 -23 J K-1
96 485 C mol-1
mol kg-1
g mol-1
6.02214*10 23 mol-1
m
8.3145 J K-1mol-1
K
J m-2
kg l-1
activity
mean activity
activity in the supersaturated solution
activity at equilibrium
concentration of the ith ion
Boltzmann’s constant
Faraday constant
molality
molecular weight of species i
Avogadro’s constant
(ν’+ν’+ν’-ν’-)1/ν’
radius of the nucleus
gas constant
temperature
valence of the ith ion
activity coefficient
mean activity coefficient
surface energy per unit area
stoichiometric coefficient
ν’+ + ν’solution density
Symbol
Dimension
Description
Cn
F
i
Jn
Q
Un
V
Δx
zn
Δ[F-]
Δф
[Amp]t
[Amp]0
[Ano]t
[Zn2+]
mol m-3
96 485 C mol-1
A m-2
mol m-2 s-1
C
m2 V-1 s-1
l
m
mol l-1
V
mol l-1
mol l-1
mol l-1
mol l-1
concentration of the nth ion
Faraday constant
current density
transmembrane flux of the nth ion
supplied charge
mobility of ion the nth ion
volume of the compartment solution
membrane thickness
valence of the nth ion
change in fluoride concentration
potential difference
concentration in the ampholyte at t
concentration in the ampholyte at t=0
concentration in the anolyte at t=0
molar increase of the zinc concentration
in the anolyte (mol/l)
Symbol
Description
superscript m
membrane phase
Chapter 5
-
1
OVERVIEW OF ZINC METALLURGY
Chapter 1 provides a detailed overview of the hydrometallurgical zinc production
process in which zinc is electrowon from zinc sulphate solution. The latter is obtained
from roasting zinc sulphide concentrates in air and subsequent leaching of the calcine
in sulphuric acid. It will be made clear that for this process, except for magnesium, the
zinc sulphate has to be very pure. As zinc sulphide concentrates typically contain a
whole range of elements, purification of the zinc sulphate solution is imperative. In
fact, the three most important process variants are named after the way the iron
impurity is removed. Another impurity that has to be removed is magnesium. The
most common process to force magnesium bleed is described and it is shown that a
large amount of waste gypsum is produced. Waste gypsum is usually dumped on the
zinc plant site, but in view of the growing environmental awareness, it is argued that
dumping of gypsum might become prohibited in the near future. Pasminco Budel
Zink is the first plant that already faces such a ban on the dumping of gypsum and
therefore magnesium removal had to be integrated with an existing biological
wastewater treatment facility. From the detailed description of this integration, it will
become clear that this process is expensive. Therefore, it will be concluded that the
development of the principles of an alternative magnesium bleed method is useful
both for Pasminco Budel Zinc as well as for other zinc plants.
1.1
Introduction
Zinc is a useful element for society with a variety of applications, which is
exemplified by Figure 1.1.
Semi-Maufactures
6%
Compounds
8%
Zinc Base Alloys
13%
Brass
18%
Others
5%
Galvanising
50%
Figure 1.1 Applications of zinc in 1999. [1]
The most widely applied function of zinc is as a coating to protect iron and steel
against corrosion. Due to its position in the electrochemical series, zinc being less
noble than iron, iron oxidation will not start as long as there is a sufficient amount of
zinc present in contact with iron or steel. Zinc forms an impervious coating of its
oxide on exposure to the atmosphere, whereas iron forms a loose, porous oxide.
Therefore, zinc corrodes at a much lower rate than iron. The application of the zinc
coating (galvanising) is accomplished electrolytically or by hot-dip methods. The
automobile industry is one of the largest consumers of galvanised steel, but it is also
extensively used in the construction industry.[2]
Prolonging the life of steel contributes to the saving of resources such as iron ore and
energy. This implies that the material cycles of zinc and iron are interconnected and
actually should be considered simultaneously when for example environmental and
economic issues regarding the zinc and steel industry are discussed.
Another major application of zinc is brass, a copper-zinc alloy with the proportion of
zinc ranging from 5 to 40%. Thus, zinc is also linked to the copper cycle. Brass is
used in plumbing, heat exchange equipment and a wide range of decorative hardware.
Sometimes other alloys are preferred as brass is rather expensive due to the copper
content, and these alloys usually contain zinc.[2]
The interconnectedness of elemental cycles is not restricted to the use phase; also
during the production phase it is meaningless to consider elements separately. Apart
from the element of interest, resources typically contain a whole range of other
elements. Clearly, deportment of these impurities to phases from which they can be
extracted in order to yield valuable products should be the fundamentally sound
objective both from environmental as well as economic perspective. This implies that
the residues mentioned in Figure 1.2, which presents a comprehensive view on the
industrial zinc cycle, should be processed and recycled to appropriate material cycles.
Figure 1.2 The industrial zinc cycle.[3]
2
If residues are not suitable for further processing, unfortunately dumping is often
opted for them. In 1979, the Second Chamber of the Netherlands adopted the Lansink
resolution, which defines the preferred order of waste treatment options. In decreasing
order of preference these options are:
•
•
•
•
•
prevention
reuse
recycling
incineration
dumping
Clearly, dumping is the least preferred measure to deal with residues. Therefore,
ideally from an environmental point of view, but with increasing penalties on
dumping of residues also from an economic point of view, alternative process routes
should be investigated in order to produce different types of residues, which are
suitable for further processing. Although quite extensive, this Ph.D. thesis can be
considered as a case study concerning the deportment of one impurity, magnesium,
from the electrolytic production process of zinc, which at present contributes to a
residue, gypsum. Gypsum cannot be processed further, because it collects various
harmful minor elements.
In the early eighties, attention of the Dutch government was focussed mainly on the
clean-up of existing dump sites. In 1989, the Dutch National Environmental Policy
Plan shifted attention to the prevention of the formation of waste. From that
perspective, it is not surprising that Pasminco Budel Zink, the only zinc refinery of the
Netherlands, was not allowed to permanently store jarosite and gypsum (both residues
from the production process) in ponds at the plant site anymore. Jarosite precipitation
is a common method for electrolytic zinc plants to reject iron, which is usually the
major impurity in zinc sulphide ores. Gypsum is produced when sulphate and
magnesium are bled from the electrolytic production process simultaneously
deporting various minor elements.
The decision of the Dutch government to prohibit further storage of jarosite and
gypsum posed for Pasminco Budel Zink the serious threat of closure. The
development of processing technology for further processing jarosite was still in its
infancy and judged not to be economically feasible to apply. It is evident that the
financial means to treat residues will be limited when competing electrolytic zinc
plants abroad are still allowed to dump their residues particularly since the profit
margins on zinc production are small.
The discovery of a large ore body (Century) in Australia with a low iron content was a
new lease on life for Pasminco Budel Zink as the production process could be altered
to treat these concentrates and jarosite storage was no longer required. For some time,
there was a discussion about the question whether the jarosite already stored in the
ponds (approximately 2 million tons) should be cleaned up; especially when it became
clear that one of the ponds leaked heavy metals and sulphates into the groundwater. In
addition, there appeared to be another source of pollution. Before Pasminco Budel
Zink (at that time Budelco) started their activities in 1973, the site was already in use
for almost a century by a zinc company employing a pyrometallurgical production
process. The residues from this process, so-called zinc ashes, were used as landfill
3
around the site. Percolating ground- and rainwater leach heavy metals from these
ashes.
In order to deal with both sources of pollution simultaneously, it was decided to
install a geohydrological management system, which pumps up the contaminated
groundwater. The groundwater is subsequently treated with sulphate reducing bacteria
(SRB). Anaerobic bacteria reduce sulphate to sulphide resulting in precipitation of
metal sulphides. Excess sulphide is oxidised to elemental sulphur. Both, metal
sulphides as well as the elemental sulphur are fed to the roaster of the zinc plant to
recover zinc and sulphur dioxide again.[4]
The same principles can be used to treat process liquor in order to bleed magnesium
and so called wash tower acids. However, the sulphate load is too high to treat such
liquors directly in the SRB plant. Therefore, Pasminco Budel Zink built a separate
biological treatment plant using PAQUES-THIOPAQ® technology in which the
moderately concentrated zinc sulphate solution can be converted to ZnS. Although the
operation costs are rather high, the production of gypsum is thus avoided. In Section
1.4, the process will be reviewed more extensively after the treatment of the standard
electrolytic process of zinc in Section 1.3
Another alternative for magnesium bleeding is simply selling part of the process
liquor to other (sometimes smaller-scale) recycling/metallurgical companies. In fact,
this is the major magnesium bleed method applied by Pasminco Budel Zink.
Summarising, Pasminco Budel Zink used to bleed magnesium from its circuit by the
combination of selling purified zinc sulphate solution and the gypsum producing
process. However, due to the ban on dumping, the gypsum producing process had to
be replaced by a biological process. For selling purified solution, Pasminco is
dependent on other companies. If the demand for purified zinc sulphate solution
decreases, all magnesium bleed has to be accomplished via the biological process,
which is rather costly. Therefore, the present Ph.D. research has been initiated with
the objective to develop the principles of an alternative magnesium bleed method,
which is environmentally as well as economically acceptable. Of course, such a
method would be applicable in other electrolytic zinc plants too.
In the remainder of Chapter 1, the principles of zinc production will be treated in
terms of composition of the raw materials being used. Emphasis will be laid on the
roast-leach-electrowinning (RLE) production process. Pyrometallurgical process
routes, although essential for the closure of the zinc cycle, will not be touched. The
flowsheet of Pasminco Budel Zink will be addressed in more detail and the present
magnesium removal technologies will be reviewed.
Chapter 2 presents an alternative closed cycle process for magnesium bleeding from
the electrolytic zinc production process. The process consists of three unit processes:
Mg-precipitation, conversion of the precipitate and electrodialysis to convert
intermediate products.
Chapter 3 elaborates on the precipitation of magnesium fluoride by mixing purified
zinc sulphate solution with zinc fluoride solution. Special attention will be paid to the
4
residual fluoride content of the equilibrium solution, as a (relatively) high fluoride
content is unacceptable in electrolytic zinc production.
Chapter 4 deals with the conversion of magnesium fluoride to magnesium hydroxide
by contacting magnesium fluoride with sodium hydroxide solution. Furthermore,
some thermodynamic aspects and the subsequent pyrohydrolysis are presented.
The subject of Chapter 5 is the electrodialysis of the resulting sodium fluoride
solution in a membrane cell reactor, yielding zinc fluoride and sodium hydroxide
solutions, which both can be recycled to the magnesium removal process.
Finally, Chapter 6 is dedicated to conclusions and recommendations. It aims at
integration of the three unit processes and discusses alternative process routes and
economic aspects.
5
1.2
Zinc metallurgy
Raw materials for zinc production can be sub-divided in the following categories:
•
•
•
zinc recycled from end-of-life products
zinc from mines
zinc metal from stock
In Figure 1.3 (which is reproduced from reference [1]), the raw materials cycle for
zinc production is depicted, demonstrating that in 1996, the amount of zinc produced
from recycling contributes to about 30% of total zinc production. Primary zinc
production has increased to 9.2 Mt in 2001[5] mainly due to increased production in
Asia. However, as other figures have not been updated yet, it was not possible to
update Figure 1.3.
Zinc from
mines 6.6 Mt
2.1 Mt
Total zinc
recycled
2.9 Mt
Zinc
Recycling
Circuit
Ne
pr w s
oc c ra
ess p
re and
sid
ue
s
Zinc metal 0.1
from stock Mt
Primary
zinc
production
7.4 Mt
Se
pr con
od da
uc r y
t io z
n inc
0.8 Mt
scrap feed
Zinc in
products
8.1 Mt
1.5 Mt
Gross consumption
9.6 Mt
Zinc recycled from
end-of-life products
1.4 Mt
Figure 1.3 Origin of raw materials for zinc production in 1996 (millions of metric tons). [1]
Although the proportion of recycled zinc is gradually increasing, ores from mines are
still the major source for zinc production. A distinction can be made between primary
and secondary zinc production. Primary zinc plants are those based almost entirely on
6
zinc concentrates, which are derived from ores (although in some cases secondary
materials are also processed). Secondary zinc plants produce zinc metal, zinc alloys
and zinc oxide through largely pyrometallurgical recycling of scrap zinc materials.
Therefore, these processes play an essential role in the closure of the zinc cycle.
Although it was already mentioned that pyrometallurgical process routes will not be
reviewed in detail, some processes are depicted in Figure 1.4 in order to illustrate the
zinc cycle. Primary as well as secondary zinc (from Zn-Refining in Figure 1.4) can be
used for the galvanising of steel, which is the major application of zinc. After a
product life, e.g. as a car, the steel scrap is recycled via a steel plant. The flue dusts
produced by these plants contain a significant amount of zinc. These dusts can be
further enriched, for example in a Waelz kiln. The zinc oxide product is mixed with
other secondary materials and lead-zinc concentrates and sintered in order to make it
suitable for treatment in the Imperial Smelting Process (ISP), which is the most
common pyrometallurgical process route for zinc. Finally, the zinc product of the ISP
has to be refined by a distillation process.
Zn/Pb Concentrate
Secondary Materials
Sintering
(Dwight-Lloyd)
Coke/Briquettes
Sinter
Hot Blast
Galvanising Scrap
Imperial Smelting
Furnace
Product
Product
e.g. car )
Zn
Slag /
Bullion/
Flue dust
Zn-Refining
Sheet steel Galvanising
Recycled Steel
Primary
Zinc
Steel Plant
Waelz Kiln / Oxide Preparation
Figure 1.4 Overview of pyrometallurgical processing (closure of the zinc cycle). [3]
More than 80% of primary zinc production takes place by means of the electrolytic
production process.[1] As the objective of this work is the development of a
magnesium bleed method for the electrolytic production process, only zinc ores from
mines will be considered. The principal mineral of concern in zinc ores is sphalerite
(ZnS) usually accompanied by galena (PbS) with impurities of pyrite (FeS2) and
chalcopyrite (CuFeS2). Although zinc ore bodies can be formed in different geological
settings, these sulphide minerals are usually precipitated from hydrothermal solutions.
The crystal lattice of sphalerite can accommodate several other elements in
substitution for zinc, e.g. iron. The formation of an ore-body can be described as a
successful attempt of nature to mobilise a substance from a large, low-grade body,
transport it, and precipitate it in a small, high-grade volume[6]. While the average zinc
content in crustal rocks is about 79 ppm, ore bodies have zinc contents ranging from
5% to more than 15% [1] [2]. Concentration of these ores at the mine site is necessary
because the metal content of the ore is usually so low that transport and direct
smelting would be too costly. Ore concentration is accomplished by means of
liberation (crushing and grinding) and subsequent flotation[3]. Grant[7] collected
averages for many concentrates, which are summarised in Table 1.1.
7
Table 1.1 World zinc concentrate data (adopted from Grant, 1993).
%
Zn
Al2O3
As
Ba
Bi
C
CaO
Cd
Cu
Fe
K2O
MgO
Mn
Na2O
P2O5
Pb
S
S as SO4
Sb
SiO2
Sn
Ti
ppm
Ag
Au
Ce
Cl
Co
Cr
F
Ga
Ge
Hg
In
Mo
Ni
Rb
Se
Sr
Ta
Te
Tl
U
Average
53.040
0.342
0.138
0.199
0.029
1.086
0.583
0.243
0.589
7.340
0.138
0.311
0.359
0.074
0.060
1.558
31.956
0.255
0.043
1.738
0.033
1.206
Average
164.1
0.7
2.9
178.5
159.0
39.9
176.1
40.0
46.7
105.7
154.4
38.9
52.0
5.0
62.3
14.0
1.1
27.1
20.2
1.0
Max.
65.200
3.200
2.000
1.791
1.000
2.950
5.000
1.000
4.250
15.000
0.480
2.000
3.000
0.340
0.070
20.000
38.400
0.400
1.000
9.550
0.400
19.000
Max.
2200.00
6.31
2.93
2100.00
3000.00
125.00
1100.00
40.00
1600.00
3000.00
2000.00
100.00
900.00
5.00
460.00
15.00
1.07
240.00
100.00
0.97
Min.
29.000
0.032
0.000
0.002
0.000
0.082
0.005
0.010
0.025
0.370
0.024
0.005
0.001
0.007
0.050
0.001
25.000
0.110
0.000
0.053
0.000
0.001
Min.
0.3
0.01
2.93
5.00
1.00
1.00
4.00
40.00
1.00
0.30
1.00
4.00
1.00
5.00
0.50
12.97
1.07
0.20
1.00
0.97
Std. Dev.
4.542
0.368
0.223
0.414
0.106
1.319
0.747
0.173
0.570
3.304
0.152
0.402
0.490
0.113
0.010
1.927
1.710
0.145
0.117
1.444
0.072
4.594
Std. Dev.
241.80
1.17
0.00
313.98
382.94
37.82
215.83
0.00
180.45
329.52
365.79
35.75
113.04
0.00
82.38
1.02
0.00
42.66
25.90
0.00
count
182
108
168
36
119
3
161
178
176
179
7
167
114
7
2
179
172
2
157
158
137
16
count
172
87
1
135
119
16
130
1
85
128
63
13
127
1
77
2
1
56
20
1
Table 1.1 clearly shows that to produce zinc, one has to deal with many other
elements. It should be realised that the standard deviations of some elements are
8
rather high and therefore, the numbers in Table 1.1 must be considered indicative. The
last column in Table 1.1 represents the number of times an element is encountered in
a concentrate sample.
1.3
Hydrometallurgical zinc production
As stated before, more than 80% of the world’s primary zinc production is
accomplished by means of an electrolytic process. Zinc sulphide concentrates are
roasted in a fluid bed roaster and the calcine is in multiple stages leached with
sulphuric acid. After deportment of iron and further purification, the resulting zinc
sulphate solution is electrolysed to produce very pure zinc (>99.995%) and sulphuric
acid, which is recycled to the leaching. The basis flowsheet of electrolytic zinc
production is depicted in Figure 1.5. In the sections below, the respective stages of the
general electrolytic process are treated in more detail.
Spent acid
to acid plant
SO2
concentrate
Roast
calcine
Leach
Solution
purification
Electrowin
Zn
Figure 1.5 Basis flowsheet of electrolytic zinc production.
1.3.1 Roasting
Zinc concentrates are fed to a fluid bed roaster with air (oxygen) at a temperature of
about 910°C as depicted by Figure 1.6.[8]
Figure 1.6 Fluid bed roaster. [3]
9
The most important roasting reactions are the following[3]:
(ZnxFe1-x)S + (31/2-1/2x)O2
2FeS + 31/2O2
2FeS2 + 51/2O2
2ZnO + SiO2
ZnO + Fe2O3
→ (11/2x-1/2)ZnO + (1/2-1/2x)ZnFe2O4 + SO2
→ Fe2O3 + 2SO2
→ Fe2O3 + 4SO2
→ Zn2SiO4
→ ZnFe2O4
The process is strongly exothermic and the heat of the reaction may be used to
produce high-pressure steam (1 kg of concentrate yields about 1 kg of high pressure
steam). Steam can be used to supply heat necessary in the leaching stage.
The calcined coarse particles leave the fluid bed reactor at the bottom, but the fine
particles are transported upward by the gas. These fine particles are caught by waste
heat boilers, cyclones and electrofilters and then combined with the coarser fraction
for further processing. In a separate plant, the sulphur dioxide gas is converted to
sulphuric acid, an important by-product, according to the following reactions[3]:
SO2 + 1/2O2
SO3 + H2O
→ SO3
→ H2SO4
1.3.2 Neutral leach
The calcine, mainly consisting of zinc oxide, zinc ferrites and zinc silicates, is usually
leached in two steps; a relatively cold, neutral leach in which zinc oxide dissolves and
a (super) hot acid leach in which the zinc content of the zinc ferrites and zinc silicates
is recovered. In the neutral leach, the calcine is contacted with dilute sulphuric acid at
a temperature of 65°C; zinc oxide dissolves:
ZnO + 2H+
ZnO + H2SO4
→ Zn2+ + H2O
→ ZnSO4 + H2O
or
Under these mild conditions, zinc ferrites will not dissolve and any iron present will
precipitate as ferric hydroxide as the compound has been traditionally described.
Dutrizac[3], however, argues that the precipitated compound does not have the
composition of Fe(OH)3. Instead, ferrihydrite, for which he suggested the chemical
notation FeO(OH, H2O, vacancy), is precipitated. Ferrihydrite has a scavenging role
and adsorbs for example SiO32-, GeO42-, SO42-, AsO43-, TeO42-, SeO32-, SeO42- and
PO43-. Cations are less extensively adsorbed, but apparently ferrihydrite can adsorb
Cu2+, Zn2+, Cd2+, Mn2+, Co2+, Ni2+ and Hg2+. Thus ferrihydrite precipitation, apart
from being a useful technology for removing the last traces of iron, contributes to
purification of the electrolyte. However, due to the poor filtration properties of
ferrihydrite, only small amounts of iron can be precipitated, not the large amounts
resulting from the dissolution of zinc ferrites. Therefore, another dissolution and
precipitation step is required to deal with the bulk amount of iron.[3] [9]
10
1.3.3 Purification
After the solid/liquid separation following the neutral leach, the zinc sulphate solution
(pH≈5) still contains some impurities such as copper, nickel, cadmium, cobalt and
magnesium. As very pure zinc sulphate solutions are required for electrolytic zinc
production, a purification step is imperative. Purification is accomplished by the
addition of zinc dust, which results in the following cementation reaction:
Me2+ + Zn → Zn2+ + Me
in which Me = Cu, Co, Ni, or Cd
Two two-step purification processes are the most common options, so called “hotcold purification” and “reverse purification”.
In hot-cold purification, zinc dust and the activator arsenic oxide (As2O3) are added at
a relatively high temperature of 90ºC, which results in the precipitation of copper,
cobalt and nickel. Subsequently, cadmium precipitation is brought about at the lower
temperature of 75-80ºC by the addition of zinc dust and copper sulphate. Usually, the
cadmium precipitate is processed further, eventually resulting in electrolytically
produced pure cadmium.[3] Until recently, cadmium could be considered a valuable
by-product. However, as the use of cadmium is restricted more and more, zinc plants
experience increasing difficulties to market their cadmium.
The reverse purification process involves addition of zinc dust at a relatively cold
temperature of 65ºC, which results in the precipitation of copper and cadmium.
Subsequently, the temperature is raised to about 80ºC; zinc dust, copper sulphate as
well as antimony tartrate are added in order to precipitate cobalt. After solid/liquid
separation, a (dilute) acid leach recovers undissolved zinc dust and redissolves the
cadmium precipitate, which can be processed further. Both copper cement as well as
cobalt cement can be sold as by-products.[3]
1.3.4 Electrolysis
The purified solution, typically containing 125-175 g/l zinc, 10-20 g/l magnesium and
2-5 g/l manganese, is cooled to a temperature of about 35ºC and subsequently
electrolysed.[3] This can be represented by the following reaction:
Zn2+ + 2eH2O
Zn2+ + H2O
→
→
→
Zn
2H+ + 1/2O2 + 2eZn + 2H+ + 1/2O2
(cathode)
(anode)
From thermodynamics, the evolution of hydrogen would be expected according to the
following half-reaction:
2H+ + 2e- → H2
It is the high overpotential for the formation of hydrogen at a zinc surface that makes
the electrolytic production of zinc possible. Impurities, sometimes even at
concentrations as low as the ppb level, could decrease this overpotential and thus
11
partly or entirely stop the electrodeposition of zinc. Other impurities (e.g. cadmium
and copper) tend to coprecipitate with zinc, thereby lowering its commercial value.
Too high a fluoride or chloride level pits the aluminium starter sheets, causing zinc to
“stick” very tightly to the starter sheets. This hampers the automatic stripping of the
zinc plates from the starter sheets, which is usually performed after 35 hours when the
thickness of the zinc plates is a few millimetres. Often additives are added to the
electrolyte, some to counteract adverse effects of remaining impurities, others to
prevent acid mist formation.[3] [9] [10] [11] [12]
A typical tankhouse may contain up to about 20,000 lead-silver anodes and a
matching number of aluminium cathodes depending on plant capacity and the area of
the electrodes (1.1 to 3.2 m2). Current densities of 400-600 A/m2 and power supplies
of 3-3.5 MWh/t may be required. After stripping, the zinc cathodes are melted and
recast into commercial ingots, which vary in size (25-4000 kg), form and
composition.[3]
1.3.5 Hot acid leach
The spent electrolyte is returned to the leaching section of the zinc plant where it is
used to leach calcine in a counter-current process. First, spent electrolyte together
with sulphuric acid (produced at the plant site) is contacted with the solid residues of
the neutral leach. Apart from the small amount of ferrihydrite, the solid residues are
zinc ferrites, zinc silicates and Pb/Ag-residues. In order not to lose the zinc from the
zinc ferrites and silicates, hot acid leach conditions (50-150 g/l H2SO4; 90°C) are
created, resulting in the following reactions:
ZnO⋅Fe2O3 + 8H+
Zn2SiO4 + 4H+
→ Zn2+ + 2Fe3+ + 4H2O
→ 2Zn2+ + SiO2·2H2O
Subsequently, another solid/liquid separation is necessary. The solid residue of this
hot acid leach contains lead and silver and can be sold to a lead smelter. However, as
a result of the hot acid leach, iron is dissolved as well. Because iron gives rise to
various problems in the production process, it has to be removed. Until the late
sixties, there was no appropriate iron removal process. Therefore, hot acid leach was
not a serious option and zinc ferrites were only pyrometallurgically treated or
dumped. At present, most electrolytic zinc plants remove iron by means of a
precipitation process; the precipitated iron compound is usually jarosite, goethite or
hematite, although the latter is only applied in the Iijima Refinery (Akita Zinc Co.) in
Japan.[3] [9] [13]
1.3.6 Jarosite precipitation
Jarosite precipitation involves the introduction of a monovalent cation at pH≈1.5 and
at a temperature of about 98ºC, thus inducing the following reaction:
M+ + 3Fe3+ + 2SO42- + 6H2O → MFe3(SO4)2(OH)6 + 6H+
in which M+ may be NH4+ ,Na+or K+
12
Apart from the introduced monovalent cation, cations already present in solution can
be incorporated, for example H+, Ag+, Tl+, Rb+, 1/2Pb2+ or 1/2Hg2+. In fact, jarosite
precipitation for most plants is the principal way to bleed thallium. Not only the
monovalent cation can be substituted by other elements. Also iron (Al3+, In3+, Ga3+
and Cr3+) and sulphate (SeO42-, CrO42- and to a small extent AsO43- and PO43-) can be
replaced. Even OH-groups may be replaced by fluoride, although it is not entirely
clear whether fluoride is substitutional or adsorbed. Jarosite filters well, although
inevitable some process liquor leaves the circuit with the jarosite, thus providing an
additional sink for impurities.[3] [9] [14]
At this point, a basic flowsheet (Figure 1.7) of an electrolytic zinc plant, which uses
jarosite precipitation, can be presented. As can be seen in the reaction equation above,
the reaction is acid producing and therefore, the addition of a neutralising agent such
as calcine or zinc oxide is an additional requirement.
Concentrates:
ZnS, FeS, FeS2
roasting
SO2
+ Cu, Cd, Mg, Pb, Ag, Co,
Ni, As, Mn, Hg, Ge etc.
sulphuric
acid
plant
calcine
electrolysis
purification
neutral
leach
ZnSO4
Calcine/ZnO +
ammonia
solid residue
Zn
H2SO4
Cd, Cu, Co (zinc ferrites, etc.)
hot acid
leach
Pb/Ag
Figure 1.7
sulphuric
acid
jarosite
precipitation
jarosite
Simplified flowsheet of the electrolytic producing process of zinc using a jarosite
precipitation process to deport iron.
If calcine would be used to neutralise the acid formed during jarosite precipitation, it
should be realised that calcine also contains zinc ferrites. These ferrites do not
dissolve under the jarosite precipitation conditions and are lost with the jarosite. In
order not to loose the zinc content of these ferrites, some zinc plants redirect the
jarosite slurry to the hot acid leach (“Dor variant”). Zinc ferrites dissolve, while
jarosite virtually remains unattacked due to the high ferric levels and ends up with the
lead/silver residue thereby making it valueless. This is the major drawback of this
alternative and plant management usually has to choose between recovering precious
metals or recovering the zinc content of zinc ferrites.[3] [9] It will be clear that the use
of zinc oxide would be advantageous in order to avoid this problem. PCT World
Patent WO 01/25497 A1[15] describes the production of zinc oxide from complex
sulphide material. Firstly, the complex sulphide material is leached with hydrochloric
acid and oxygen. Subsequently, iron is precipitated (as akageneite) from the leach
solution using magnesium oxide and lead, copper, silver, cadmium and cobalt are
13
recovered by cementation with zinc dust. In a second precipitation step, zinc oxide is
precipitated using magnesium oxide. The residual magnesium chloride solution may
be spray roasted to regenerate hydrochloric acid and magnesium oxide. The process
allows for the on-site treatment of low-grade material, resulting in the production of
an iron-free zinc oxide that can be shipped to an electrolytic zinc refinery. In spite of
the apparent advantages, this development has not been implemented (yet).
Jarosite is not marketable and therefore has to be dumped in storage ponds with all the
impurities it contains. This implies both the desired deportment of harmful impurities
(e.g. AsO4- and F-) as well as the loss of valuable by-products (e.g. In, Ga, Ag).[3]
Jarosite precipitation is the most widely employed method to bleed iron, because of its
low operating costs. However, in view of increasing environmental awareness, this
situation is likely to change as zinc plants might be held responsible for their wastes
and possibly could be charged in proportion to the mass or volume of waste they
produce. In this respect, jarosite will be the most costly option as the iron content per
ton of jarosite is lower than for goethite and hematite.
Anticipating future liabilities, large effort has been made to increase the stability of
jarosite as attempts to convert jarosite to a marketable product all have failed.
Stabilisation by mixing with cement seems to be the most successful stabilisation
method, although it should be realised that the amount of waste to be dumped actually
will be increased. Furthermore, the long-term stability of the products is unknown and
the method cannot be used for ammonium jarosite due to the likely release of
ammonia in alkaline media.[3]
As an alternative, a waste-to-waste technology involving the mixing of jarosite and
sewage sludge in a pressure vessel has been proposed resulting in the oxidation of the
sewage sludge according to the following reaction (for sodium jarosite):
NaFe3(SO4)2(OH)6 + (C6H10O5)n + 11/2MgO →
Fe3O4 + n[TOC + C +CO2] + 2SO42- + Na+ + H2O + 11/2Mg2+
Since the organic products were not fully identified, they are represented by TOC
(total organic carbon). Magnesium is added as a neutralising agent in order to prevent
metal dissolution. After solid/liquid separation, the effluent can be treated in a
biological process, while the solid residue is more stable than jarosite. These residues
could be melted at 1500ºC in order to evaporate zinc, which can be used as secondary
raw material upon condensation. The remaining melt containing iron, sand and clay
minerals can be cast into building blocks.[16]
1.3.7 Goethite precipitation
Goethite precipitation is also a widely employed iron removal process. Goethite is the
orthorombic or α-form of the four different FeO.OH polymorphs. The amount of iron
per tonne of goethite is higher than for jarosite, but lower than for hematite.[3]
Goethite can only be precipitated from dilute Fe2(SO4)3 solutions (<2-3 g/l Fe3+),
which seems to make the method not suitable for the zinc industry as usually tens of
grams of iron are dissolved per litre hot acid leach filtrate. However, two different
14
processes are in use solving this problem, the so-called Vieille-Montagne process and
the paragoethite process.[3]
In the Vieille-Montagne variant, the ferric iron concentration is controlled by
reduction with concentrate at a temperature of about 90°C:
2Fe3+ + ZnS → 2Fe2+ + Zn2+ + S0
After solid/liquid separation, the resulting residue (elemental sulphur and unreacted
concentrate) is recycled to the roaster. Subsequently, the reduced solution is
neutralised with calcine to 3-5 g/l H2SO4 and another solid/liquid separation yields a
zinc ferrite residue, which can be recycled to the hot acid leach. Finally, oxidation is
accomplished by injection of air and simultaneous calcine addition in order to keep
the pH between 2.8 and 4, resulting in the immediate precipitation of goethite:
4Fe2+ + O2 + 6H2O → 4FeOOH + 8H+
As the rate of precipitation equals/exceeds the rate of oxidation, the ferric iron
concentration remains below the critical 2-3 g/l. Calcine or zinc oxide has to be added
in order to neutralise the produced acid.[9] A flowsheet of the Vieille-Montagne
goethite process is depicted in Figure 1.8.
S0
Concentrates:
ZnS, FeS, FeS2
+ Cu, Cd, Mg, Pb, Ag, Co,
Ni, As, Mn, Hg, Ge etc.
electrolysis
H2SO4
purification
Cd, Cu, Co
roasting
SO2
sulphuric
acid
plant
sulphuric
acid
calcine
neutral
leach
ZnSO4
concentrate
solid residue
(zinc ferrites, etc.)
S0
hot acid
leach
reduction
calcine
Pb/Ag
zinc ferrite neutralisation
air
goethite
calcine
goethite
precipitation
Figure 1.8 Simplified flowsheet of an electrolytic zinc plant employing the Vieille Montagne goethite
process to deport iron.
15
As it is not possible to subject the goethite to an acid wash (which would result in
dissolution of the goethite), the ferritic part of the calcine added during goethite
precipitation is lost. Therefore, also in the goethite process the use of zinc oxide
instead of calcine is advantageous. Numerous active sites on the goethite surface give
the process also a valuable purifying function. Another advantage in environmental
sense is that goethite, unlike jarosite, can directly be pyrometallurgically treated, for
example in an Ausmelt reactor or a Waelz kiln in order to produce an environmental
acceptable slag to discard. As smelting operations are rather costly, efforts have been
undertaken to stabilise the goethite residue in a low temperature process. An example
is the Graveliet® process developed by Umicore (Union Minière), which implies the
mixing of washed goethite residues with various steel plant slags, lime and possibly
cement in order to produce a hard and inert gravel that can be used in the building
industry.[3][9][17]
The paragoethite process accomplishes the required reduction of the ferric iron
concentration by diluting the hot acid leach solution in a large reservoir of ferrous
sulphate solution at a pH 3.0-3.5 and a temperature of 85-95ºC. Simultaneous addition
of a neutralising agent results in the immediate precipitation of goethite. [3]
1.3.8 Hematite precipitation
The flowsheet of a hematite producing zinc plants is provided in Figure 1.9.
Concentrates:
ZnS, FeS, FeS2
+ Cu, Cd, Mg, Pb, Ag, Co,
Ni, As, Mn, Hg, Ge etc.
electrolysis
purification
Cu, Cd, Co
Zn
H2SO4
SO2
roasting
SO2
sulphuric
acid
plant
H2SO4
calcine
neutral
leach
residue
zinc ferrites etc.
leaching/
reduction
CaCO3
neutralisation
pH =2
CaCO3
neutralisation
pH = 4.5
H2S
gypsum
copper
precipitation
hematite
precipitation
gypsum
O2, steam
p = 20 bar
CuS
hematite
Figure 1.9 Simplified flowsheet of an electrolytic zinc plant employing the hematite process to deport
iron.
16
The neutral leach residue (mainly zinc ferrite) is reductively leached with sulphur
dioxide at a temperature of 105ºC in order to introduce iron to the solution in the
divalent state. For precipitating hematite, the following reaction has to be brought
about:
Fe2(SO4)3 + 3H2O → Fe2O3 + 3H2SO4
In order to achieve a reasonable reaction rate, temperatures of about 200ºC are
required, which implies that capital intensive pressure vessels (autoclaves) are needed.
The precipitation reaction is carried out at 200ºC resulting in a pressure of about 20
bar. In order to support the reaction, all produced acid as well as all acid already
present in solution has to be neutralised with limestone or calcine. If limestone is
used, the produced gypsum of the first neutralisation is found to be saleable, the
gypsum of the second neutralisation is contaminated and must be considered chemical
waste.[3] [9] [14]
The iron content of hematite (Fe2O3) is the highest of the three precipitation methods.
Besides, hematite can, in principle, be used as a raw material for steelmaking.
However, in practice, the content of elements such as zinc, cadmium and residual
sulphate in the hematite residue has been too high for further processing in the steelmaking industry and the only application seems to be in the cement industry.
Although from an environmental perspective hematite precipitation seems to be the
“best” method to deport iron, the high costs of the equipment involved have limited
the widespread application of this process variant strongly.[3]
1.4
Magnesium control in electrolytic zinc plants
In the preceding sections, it was made clear that electrolytic zinc production also
implies processing all kind of other elements. One of these elements is magnesium,
which usually originates from dolomite or serpentine in concentrates. During roasting,
dolomite is transformed into calcium sulphate and magnesium compounds. These
magnesium compounds occur as solid solutions with zinc and iron compounds in the
calcine; on the one hand as easily soluble oxides, orthosilicates and sulphates and on
the other hand as sparingly soluble ferrites. Thermodynamic calculations of this
system showed that it is possible to predict the distribution of magnesium over the
respective phases based on the chemical composition of the concentrate and the
reaction temperature in the roaster.[18]
The average magnesium content of zinc concentrates is not particularly high
(typically 0.2wt%[7]), but as there is hardly any natural magnesium bleed,
accumulation can occur in process liquors (up to the saturation level if no measures
are taken). When magnesium is allowed to accumulate to 8 g/l at a current density of
600 A/m2, the voltage drop per cell increases by 0.2 V due to reduced conductivity of
the solution. Thus power consumption is increased by about 6% when we recall the
usual voltage drop of 3 V and the same current efficiency. In addition, cooling
requirements are increased as well[12]. Furthermore, high magnesium levels result in
increased density and viscosity of the circulating solutions, which increases the
required pumping capacity and may cause problems with filtering and settling in the
17
leaching and purification section of the zinc plant. For these reasons, the magnesium
level must be controlled between 10 and 15 g/l.[18], [19], [20]
The only ‘natural’ magnesium bleed in the process is solution losses (e.g. entrapped
liquor in iron precipitates). It will be clear that this bleed decreases when zinc
recovery is increased by, for example, more thoroughly washing of iron precipitates.
Usually, additional measures have to be taken to control the magnesium level in the
zinc plant. Various methods to bleed magnesium are listed below:
•
•
•
•
•
•
basic zinc sulphate process
selling process solution
treatment of neutral leach residue pyrometallurgically instead of by hot acid
leaching
further electrolysis of spent electrolyte (in combination with a dialysis process)
preleaching of concentrates
integrated into a biological waste water treatment process
(Pasminco Budel Zink, §1.5)
A widely employed method to control the magnesium balance (as well as the
sulphate, chloride and water balance) is the so-called Basics Process, which was
developed by the ‘Electrolytic Zinc Company of Australasia’ in the early seventies. In
this process, zinc is selectively precipitated as basic zinc sulphate by addition of a
neutralising agent such as slaked lime:
3CaO + 4ZnSO4 + 13 H2O → (3Zn(OH)2)·ZnSO4·4H2O + 3CaSO4·2H2O
Other neutralising agents such as limestone can also be used. After solid/liquid
separation, the filtrate, depleted in zinc, but still containing magnesium (and
sulphate), can be used to wash iron precipitates and thus leave the process as
entrapped solution. Alternatively, the filtrate can be treated in a wastewater treatment
device and subsequently discharged. The basic zinc sulphate and unreacted limestone
can be used as neutralising agent in other parts of the zinc plants, which results in the
dissolution of these compounds. Gypsum, however, does not dissolve and has to be
dumped as it contains impurities, which exclude the possibility of selling it.[20],[21]
Solvent mediated conversion of gypsum to hemihydrite or anhydrite could be a
possible process route. Impurities are temporarily released and can captured with an
extractant, ion exchanger or hydrocyclone (insoluble particles). However, it should be
realised that in order to be saleable the product will have to be very pure. Besides, this
process route makes a zinc plant dependent on other companies. If the demand
decreases, the converted product still has to be disposed of.
It should be realised that the amount of magnesium that has to be bled from the
process is actually increased when limestone is used, as limestone typically contains
0.4 wt.% magnesium[21]. Calcine or zinc oxide could also be used as a neutralising
agent to precipitate basic zinc sulphate, thus no additional magnesium would be
introduced and more important, the precipitation of gypsum would be avoided:
ZnSO4 + 3ZnO + 7H2O → ZnSO4·3Zn(OH)2·4H2O
18
If a zinc plant is not allowed to dump gypsum anymore, the Basics process using
calcine as a neutralising agent would be an attractive alternative. Nevertheless, some
considerations should be taken into account. In the Basics process, the filtrate is
usually used to wash other precipitates and/or leach residues. However, when a plant
is not allowed to dump gypsum anymore, it is likely that the plant is not allowed to
dump jarosite or goethite either. Some zinc plants may be able to obtain concentrates
with a very low iron content (like Pasminco Budel Zink) in which case it is not
possible to bleed magnesium as entrapped solution. Other zinc plants may decide to
treat their iron residues pyrometallurgically (e.g. with Ausmelt technology) in which
case it is still possible to discard some entrapped solution. In that case magnesium
will end up in a slag. Finally, new iron removal methods could be developed when a
ban on dumping iron residues becomes operative in more countries than the
Netherlands and it is uncertain if such new processes offer the possibility to bleed
filtrate.
The use of calcine instead of limestone also involves some process adjustments. In
order to remove sufficient sulphate, the reaction has to be carried out in autoclaves or
calcine has to be milled, which are both rather costly process modifications [22]. Even
if the reaction approaches completion, there will be still some sulphate left. It is
probably not allowed to discharge sulphate streams to surface water. For example,
Pasminco Budel Zink faces a standard of 450 ppm (Maximum Allowable
Concentration)[22]. Pasminco Budel Zink would be able to send this low sulphate
stream to the sulphate reducing bacteria (SRB) plant, which will be described in
Section 1.5. Other zinc plants, however, usually do not possess such an installation.
Furthermore, there is still some zinc left in the solution, typically 1 g/l and a separate
effluent treatment plant is required to recover this zinc.
Another magnesium bleed method that can be applied by zinc plants is selling
purified solution to other companies. However, selling purified solution makes a zinc
plant dependent on other companies and the agreed price is usually not very
favourable. Besides, if the demand for purified zinc sulphate solution decreases,
alternative methods should be available.
Instead of subjecting it to a hot acid leach, the solid residue of the neutral leach could
also be pyrometallurgically treated in order to recover zinc and lead, etc. In that case,
the prediction of the distribution of magnesium over the respective phases based on
the chemical composition of the concentrate and the reaction temperature in the
roaster is particularly useful[18]. When a substantial part of magnesium is present in
zinc ferrites, Ausmelt submerged lance technology could be applied. In a paper by
Altepeter and James[23], it is shown that virtually all magnesium will end up in a slag,
which subsequently passed the EPA TCLP (Toxicity Characteristic Leaching
Procedure) test satisfactorily. This way of operation implies major alterations of the
production process and consequential high investment costs. The choice for such a
process will not depend solely on the need for a magnesium bleed method.
Another proposal is to subject spent electrolyte to further electrolysis until most of the
zinc is removed from solution and then pass the strongly acidic magnesium sulphate
solution to a waste water treatment plant.[21] However, the kind of wastewater
treatment facility was not specified, although the costs of operating this process were
considered unattractively high.
19
1.5
Pasminco Budel Zink – residue free zinc production
As mentioned in Section 1.1, Pasminco Budel Zink is not allowed to dump any
residues at the plant site anymore. Consequently, the process had to be altered in order
to prevent the formation of jarosite and gypsum. The use of Century concentrate with
its very low iron content required major process modifications. To prevent the
dissolution of zinc silicates, the neutral leach is kept at higher pH (pH≈4). Sufficient
zinc oxide dissolution is accomplished by increasing the residence time and crushing
and grinding the calcine more intensively (80%<150μm). The neutral leach flow is
aerated resulting in oxidation of the small amount of iron present, which results in the
immediate precipitation of iron hydroxides containing impurities such as arsenic and
antimony. After solid-liquid separation, the major part of the neutral leach solution is
sent to purification. Spent acid is added to the neutral leach residue and the mixture is
heated to 95ºC. Simultaneously, zinc silicates dissolve and SiO2 precipitates. Apart
from silica, the solid residue contains lead and silver. After solid-liquid separation,
this so-called Budel Leach Product (BLP) can be sold to a lead smelter.[24]
In 1999, the SRB-plant was expanded in order to treat low sulphate (<1g/l)
wastewater streams from the zinc plant in addition to the contaminated groundwater.
However, with the ban on the dumping of gypsum two more bleed streams must be
treated:
•
•
wash tower acid (scrubber discharge from the roaster acid plant)
magnesium bleed
These streams have a high sulphate content and cannot be treated directly in the SRBplant. Therefore, a Biological DeSulphurisation (BDS) plant has been developed and
installed, which is capable of dealing with more concentrated sulphate streams. The
production of gypsum could be avoided by integrating the BDS plant with the existing
SRB plant. Figure 1.10 schematically depicts the combination of these two biological
processes and clearly shows that magnesium finally leaves the process as a dilute
MgCl2 solution.
calcine CaCl2 Mg-bleed
typically
0,5 m3/h
Contaminated groundwater
biogas
typically
285 m3/h
scrubber
acid
Effluent
typically
25 m3/h
BDS
CaF2
ZnS
SRB
dilute
MgCl2
solution
Sludge
Figure 1.10 Simplified combined flowsheet of BDS and SRB.
20
More detail concerning these processes can be found in the remainder text and figures
of this section. As described in Section 1.1, a geohydrological containment system (12
shallow and deep wells) was installed in the early nineties as the groundwater beneath
the plant site had become contaminated with heavy metals (notably zinc and
cadmium) and sulphates. Approximately 2.5 million m3 groundwater is pumped up
annually and treated in the sulphate reducing bacteria (SRB) plant. In Figure 1.11, the
Upflow Anareobic Sludge Blanket (UASB) reactor is depicted.
Biogas
CH4, CO2, H2S
effluent
influent
ethanol
sludge
metal sulphides
biomass
Figure 1.11 Schematic representation of the UASB reactor (reproduced from Scheeren et al., 1993 [4]).
The sulphate and heavy metals containing influent enters the reactor and is distributed
evenly over the active biological bed. Anaerobic bacteria reduce sulphate to sulphide
using ethanol as a carbon source and electron donor, which immediately results in the
precipitation of insoluble metal sulphides. As a result acetate is produced, which is
converted to carbon dioxide and methane by methanogenic bacteria. The biogas is
trapped by the ‘hoods’, leaving the top layer of the reactor relatively undisturbed,
which facilitates settling of solids. These solids will settle on the tilted plates of the
gas traps and eventually sink into the sludge bed.[4]
The biogas is periodically burned in a flare after scrubbing the hydrogen sulphide gas.
The liquid effluent still contains sulphide, which is oxidised to elemental sulphur by
aerobic bacteria in a subsequent reactor. Controlled air addition must prevent further
21
oxidation of elemental sulphur to sulphate. A tilted plate settler and ‘Dynasand’ filter
recover more solids, which are flushed out of the reactor. Together with the sludge of
the UASB reactor and the elemental sulphur, these solids are fed to the roaster. In
Figure 1.12, a flowsheet of the SRB process is depicted.[4]
NaOH-solution
CH4
flare
scrubber
CO2 H2S CH4
compost filter
sulphide oxidation
reactor
liq
(H2S)
INFLUENT
EFFLUENT
O2
UASB
CO2
SLUDGE TREATMENT
(ROASTER)
(S)
SANDFILTER
liq
TILTED PLATE
SETTLER
Figure 1.12 Flowsheet of the SRB process [3].
A flowsheet of the BDS plant is depicted in Figure 1.13. The influent flow (wash
tower acid) is typically 25 m3/h and has the following composition[25]:
•
•
•
•
10 g/l H2SO4
0.5 g/l HF
1 g/l HCl
0.5 g/l Zn
Firstly, the wash tower acid is neutralised in two steps with calcine (to pH=3) and
sodium hydroxide (to pH=5.1), which results in a zinc sulphate solution. The slurry is
aerated, thereby oxidising and precipitating the iron and other minor trace elements.
The neutralisation residue and iron are extracted in a thickener underflow. In Figure
1.13, this neutralisation step is represented as a reaction in a single vessel.
Subsequently, the fluoride is removed as CaF2 in a Crystalactor®. The Crystalactor®
is a cylindrical vessel, which is operated based on the concept of supersaturation
through dilution. The reactor is partly filled with seed material to which CaCl2 and
recirculating solution are fed in an upward stream at such a rate as to fluidise the
particle bed. The effluent overflows the top of the reactor, while fluorite particles
become progressively heavier. As a result, they move towards the bottom of the bed
of which the lower portion is periodically discharged.[26]
Only then, the magnesium bleed (typically 0.5 m3/h) can be added, because the
sulphate content of this bleed is too high to add it prior to fluoride removal. Gypsum
would precipitate rather than fluorite. After this addition, the zinc sulphate solution is
reduced by bacteria, which use H2 as an electron donor, resulting in the precipitation
22
of ZnS. The hydrogen gas is produced on site from natural gas according to the
following reaction:
CH4 + 2H2O → 4H2 + CO2
After solid-liquid separation and dewatering, the zinc sulphide is fed to the roaster,
while the effluent goes to the SRB-plant. This way, magnesium finally leaves the
process as a very dilute MgCl2 solution.[25]
Magnesium bleed
Calcine (ZnO)
Overflow to SRB
Scrubber acid
CaF2 pellets
crystalactor
H2/CO2
ZnS to roasting
Figure 1.13 Schematic representation of the BDS-plant
[25]
.
23
2
AN ALTERNATIVE MAGNESIUM BLEED
PROCESS
In this chapter, a new approach to bleed magnesium is presented and discussed, which
involves magnesium fluoride precipitation from purified zinc sulphate solutions. The
magnesium fluoride is contacted with sodium hydroxide solution and thus converted
to magnesium hydroxide, which could be a saleable product. It is shown that the
resulting sodium fluoride solution can be electrolysed in a reactor equipped with ion
exchange membranes, an active zinc anode and a stainless steel cathode. This yields
zinc fluoride and sodium hydroxide solution, which both can be recycled to the
proposed magnesium removal process. Chapter 2 has been published in Minerals
Engineering[27].
2.1
Introduction
From the description of presently applied magnesium bleed methods in Sections 1.4
and 1.5, it will be clear that a governmental ban on the production of gypsum or high
charges on dumping would pose serious problems for zinc plants. Alternatives such as
biological wastewater treatment or a pyrometallurgical process route imply high
investment costs, which can only be justified when, apart from providing a
magnesium bleed method, other objectives are served. Therefore, the development of
an alternative method to exclusively remove magnesium from zinc plant solution in a
closed cycle process is imperative.
The composition of a typical purified solution of a zinc plant is summarised in Table
2.1 indicating that in addition to zinc sulphate, the process liquor contains magnesium
sulphate and manganese sulphate. As Inductively Coupled Plasma (ICP) Spectroscopy
was not available, X-ray fluorescence (XRF) analysis was selected to characterise the
solution composition. Although this analysing method is not as accurate as atomic
absorption spectrometry (AAS), it has the advantage of providing information on a
wide range of elements instead of just a single one as is the case with AAS. It was
considered necessary to get information of more elements than could be provided by
the available hollow cathode lamps. Note that the measured sulphate concentration
reports too low a value; it should be approximately 275 g/l.
Table 2.1 Composition of purified electrolyte (determined with XRF).
contents
Zn
SO4
Mg
Mn
Traces of Ca, Si. K, Al, Cl
g/l
146.6
256.5
13.8
3.2
< 0.5
Instead of precipitating a zinc compound (basic zinc sulphate in the Basics process),
which leaves behind a magnesium/manganese sulphate solution with residual zinc to
process, it was chosen to selectively precipitate a magnesium compound out of
solution. Therefore, a (near) insoluble magnesium salt had to be found of which the
corresponding zinc salt had to be rather soluble. Such a compound could readily be
identified in literature sources such as the Handbook of Chemistry & Physics[28],
which lists the solubility products of magnesium fluoride (7.42*10-11) and zinc
fluoride (3.04*10-2). Based on these solubility products, it was inferred that it would
be possible to precipitate magnesium fluoride out of purified zinc sulphate solution,
while zinc remains dissolved. The solubility products of magnesium fluoride and
calcium fluoride are comparable (Ksp,MgF2 = 7.42*10-11 and Ksp,CaF2 = 1.46*10-10),
while the magnesium concentration significantly exceeds the calcium concentration in
the electrolyte (13.6 g/l Mg and 0.3 g/l Ca). Therefore, it is not expected that calcium
fluoride will precipitate first, acting as nucleus for magnesium fluoride precipitation.
Literature sources[29],[30] describe magnesium fluoride precipitation by a constant
composition method. This method involves the slow mixing of magnesium
nitrate/chloride and potassium fluoride solutions. In our situation, slow addition of
zinc fluoride solution was opted for in order to prevent the introduction of
environmentally and operationally unfavourable cations to the system. A set of
precipitation experiments was performed with purified zinc sulphate solution.
Pasminco Budel Zink provided actual purified zinc sulphate solution. In industrial
solutions often impurities are present, which could have an effect on the precipitation
mechanism. Thus, it was ensured that if impurities would be involved in the onset of
precipitation that it could be repeated real plant situations. The first goal of the
precipitation experiments was to determine the selectivity of the process (i.e. the ratio
Mg:Zn in the precipitate). Furthermore, filtration behaviour was studied as well as the
residual fluoride content in the filtrate. This is crucial since a high fluoride content in
zinc electrolyte results in pitting of the aluminium starting sheets in the cellhouse of
the zinc plant. In turn, this results in a very firm sticking of zinc to the cathode
starting sheets.[10],[12],[31]
There are a few possibilities how to deal with the precipitated magnesium fluoride:
•
•
•
find an application for magnesium fluoride
further processing of magnesium fluoride (in order to recycle the fluorine)
dumping
In view of the Lansink resolution, which was discussed in Chapter 1, dumping is not
an environmentally sound option. However, the amount of waste is dramatically
reduced when compared to the Basics process. The amount of gypsum produced in
the latter process is proportional to the (high) zinc content of the electrolyte solution,
while the amount of magnesium fluoride obviously is proportional to the (relatively
low) magnesium content. An application for industrially produced bulk amounts of
magnesium fluoride is not easy to find. In Chapter 6, dumping as well as possible
applications of magnesium fluoride will be further discussed. Here, it is proposed to
contact magnesium fluoride with sodium hydroxide solution in order to produce
magnesium hydroxide. This produced magnesium hydroxide might be saleable and
the resulting sodium fluoride solution can be electrolysed in a membrane cell reactor,
similar as described by Liao et al.[32] to produce zinc fluoride and sodium hydroxide
solutions. Other possible process routes for magnesium fluoride are also discussed in
Chapter 6.
26
In view of the above, the scope of this chapter is that:
•
•
2.2
Each of the above mentioned process steps will be discussed in terms of various
laboratory experiments, and that
the process steps will be integrated in a provisional flowsheet
Experimental
In this section, the various experimental methods used will be described in detail.
2.2.1 Precipitation
In the precipitation experiments, the following procedure was followed:
•
Simultaneously, 2.1 litre 0.12M ZnF2 solution and 0.8 litre purified zinc sulphate
solution (electrowinning feed) were slowly added to 0.4 litre distilled water over a
period of 45 minutes. The experiments were performed in a baffled vessel (4
litres) for which the diameter = height of vessel and with baffles (4 total) = 1/10 of
diameter.
•
The stirring rate was 200 rpm and after the addition, the suspensions were stirred
for another 75 minutes. The impeller width was 1/3 of the diameter of the vessel.
•
The pH of the industrial purified zinc sulphate solution was low (pH=2.2) and the
pH of zinc fluoride solution was about 5.1. Generally, the pH dropped in the first
few minutes of the experiments (as the starting liquid was distilled water with
pH=7) below pH=4. During the addition of solutions, the pH steadily decreased to
pH≅3.5 and during the additional stirring time further to pH≅3.
•
Four experiments were performed at room temperature (291–295K), two at 327K,
one at 331K and one at 356K.
•
The suspensions were filtered over a ‘Weißbandfilter’ (pore size diameter 4-7 µm)
at 0.5 bar vacuum.
•
During filtration, the filtrate volume vs. time was recorded in all four experiments.
•
The residual fluoride content of all filtrates was determined with a specific
fluoride ion electrode. The use of the ion specific electrode was justified by a
simple verifying experiment. A small amount of sodium fluoride (yielding 291
ppm F) was added to 0.7 M zinc sulphate solution. A triple measurement with the
ion specific electrode respectively yielded 289 ppm, 289 ppm and 285 ppm.
•
The precipitates were washed with dilute sulphuric acid of pH=4, filtered again
and dried at 303K overnight in 5 experiments and at 333K in two experiments.
The precipitates were characterised with X-ray fluorescence (XRF) spectrometry
(Philips PW2400, Uniquant software) and X-ray Diffraction (XRD) (Philips
27
PW3710). Of two precipitates, duplo XRF measurements were obtained and
treated as separate measurements. In Table 2.2, the test matrix is summarised.
Table 2.2 The matrix of the precipitation experiments.
Treaction/K
291
293
294
295
327
327
331
356
Tdrying/K
303
303
303
303
303
303
333
333
XRD
ü
ü
ü
ü
ü
ü
ü
ü
# XRF
1
1
2
1
2
1
1
1
fluoride
ü
ü
ü
ü
ü
ü
ü
ü
filtration
ü
ü
ü
ü
2.2.2 Conversion
Two conversion experiments were performed, one at 291K and one at 327K. The
method followed is:
•
Three grams of magnesium fluoride, produced according to the method described
above, were added per litre stirred 0.2M sodium hydroxide solution into a beaker.
•
After 8 hours, the suspensions were centrifuged and decanted, the precipitate was
washed with distilled water and centrifuged and decanted again.
•
Subsequently, the precipitate was dried overnight at 303K and analysed with XRF
and XRD. For the experiment performed at 327K, a sample was also prepared
after 4 hours contacting with sodium hydroxide solution.
2.2.3 Electrolysis
The electrolysis experiment was performed in a membrane cell reactor, which is
depicted in Figure 2.1.
•
A 0.2M NaF solution was placed into the middle compartment of the membrane
cell reactor (this solution has been termed ‘ampholyte’), 0.01M H2SO4 solution
was used as anolyte, while 0.01M NaOH solution was used as the catholyte.
•
Upon application of an electrical potential difference, the zinc anode gradually
dissolves and fluoride ions move through the anion exchange membrane (AEM).
As a result of the electrical potential, sodium ions move through the cation
exchange membrane (CEM) towards the cathode, where hydroxide ions and
hydrogen gas are produced.
28
•
During a period of 90 minutes, a current of 0.2 A was realised by applying a
potential difference.
•
Subsequently, the compartments were emptied and the composition of the
solutions was determined by XRF. Furthermore, the fluoride content of the
anolyte was measured with the specific fluoride ion electrode.
ZnF2
AEM
CEM
+
Na+
FZn2+
OH-
H2
NaF
zinc
Zn
Zn2+
+
2eV
2H2O + 2e2OH- + H2
NaOH
Figure 2.1 Membrane cell, which is divided in three compartments by a Neosepta-AHA anion
exchange membrane (AEM) and a Nafion-350 cation exchange membrane (CEM).
2.3
Results
The results of the experiments described in Section 2.2 are presented in the sections
below.
2.3.1 Precipitation
In all cases, XRD results confirmed that magnesium fluoride had been precipitated,
although the peaks were broad and sometimes a little bit shifted with respect to their
theoretical positions. This could indicated that some of the precipitates have formed
as amorphous materials instead of crystalline materials.
The results of XRF analyses are depicted in Figure 2.2. The experiment at room
temperature with a drying temperature of 303K was repeated three times. Averages
were calculated and error bars give minimum and maximum values. The experiment
at 327K with a drying temperature of 303K was also repeated. The other two
experiments were not repeated. Measurements of calcium, potassium and manganese
were all below 0.5 wt.% and are not provided in Figure 2.2.
29
5
40
4
30
3
20
2
10
1
0
0
Mg
F
LOI
Zn
wt.%
wt.%
50
SO4
room T, drying 303K
327K, drying 303K
331K, drying 333K
356K, drying 333K
Figure 2.2 Composition of the precipitates for different reaction and drying temperatures. LOI = ‘loss
on ignition’. Note the different scales of the y-axis (left and right).
The magnesium content is on average about 31 wt.%, the fluoride content varies, from
38 to 47 wt.%. The molar magnesium: fluoride ratio is not 1:2 as would be expected,
but 1:1.8 on average. In Section 3.3, it is proven that deviation of the stoichiometric
ratio is caused by substitution of fluoride by hydroxyl groups.
The residual fluoride content of the filtrates, as measured with the specific fluoride
ion electrode, is given as function of reaction temperature in Figure 2.3. The residual
fluoride content of the filtrates decreases with increasing reaction temperature.
residual fluoride (ppm)
400
350
300
250
200
150
100
50
0
253
273
293
313
333
353
373
temperature (K)
Figure 2.3 Residual fluoride content of the filtrate as a function of reaction temperature.
The suspensions were filterable by filters with pore sizes 4-7 µm, which are similar to
industrial filters[22]. Filtration characteristics are depicted in Figure 2.4, two records of
filtration at room temperature are shown, one for filtration of precipitates produced at
331K and one for precipitates produced at 356K.
30
volume (ml)
1800
1600
1400
1200
1000
800
600
400
200
0
356 K
331 K
room T
room T
0
1000
2000
3000
4000
time (s)
Figure 2.4 Filtrate volume vs. time for precipitates produced at different temperatures; the filter area
19.6 cm2.
The slope of the high temperature filtration line corresponds to a filtration velocity of
about 500 l m-2 h-1 at a particle density of about 5 g/l.
2.3.2 Conversion
The composition of the precipitate after it is contacted with sodium hydroxide
solution is given in Figure 2.6 and Figure 2.6. X-ray diffraction analyses on converted
precipitates indicate that magnesium hydroxide (brucite) indeed has been formed.
From Figure 2.6 and Figure 2.6 it can be inferred that magnesium fluoride is
converted into magnesium hydroxide. The fluoride content has remarkably decreased,
but there is still about 4 wt.% fluorine and small amounts of zinc and calcium.
wt.%
40
35
30
25
20
15
8 hours at 291K
4 hours at 327K
8 hours at 327K
10
5
0
Mg
F
Figure 2.5 Composition of converted precipitate: Mg and F.
31
1
wt.%
0.8
8 hours at 291K
0.6
4 hours at 327K
0.4
8 hours at 327K
0.2
0
Zn
Figure 2.6
Ca
Composition of converted precipitate (other elements were all below 0.1 wt.% and are
not given).
2.3.3 Electrolysis
In Figure 2.7, the composition of the solutions (Zn, Na, F), as determined with XRF
and the specific fluoride ion electrode before and after the electrolysis experiment is
summarised.
Anolyte
Ampholyte
Catholyte
content (g/l)
4
3
2
1
0
Zn
before after
F
before after
Na
before after
Na
before after
Figure 2.7 Composition of different compartments of the membrane cell reactor before and after
sodium fluoride electrolysis.
The zinc concentration of the anolyte increased from zero to about 1.3 g/l and the
fluorine content from zero to 0.3 g/l. The sodium content of the ampholyte decreased
from 3.6 g/l to 2.2 g/l and the sodium content of the catholyte increased from 0.3 g/l
to 1.4 g/l.
32
2.4
Discussion
In the sections below, the results will be discussed of successively the precipitation
experiments, the conversion experiments and the experiments with the membrane cell
reactor.
2.4.1 Precipitation
As can be deduced from Figure 2.2, the composition of the magnesium fluoride
precipitates is more or less constant and apparently independent of reaction
temperature or drying temperature. The zinc content is rather low and varies between
0.3 and 1 wt.%, which results in a Mg:Zn ratio varying from 98:1 to 30:1. Therefore,
the process is sufficient selective, as the equivalent zinc loss in that case varies from
10 to 30 kg per ton magnesium removed (on average 20 kg zinc loss per ton
magnesium). It should be realised that the washing of the precipitate with dilute
sulphuric acid after filtration is very important as control measurements of unwashed
precipitates revealed much higher zinc contents (5–6 wt.%).
The loss on ignition (LOI) is mostly water. Obviously, the differences in LOI of
different groups are caused by the change in drying temperature. It should be realised
that this change slightly influences the absolute values of the rest of the composition.
Acceptable fluoride levels in purified zinc sulphate solution are below 10 ppm.[12],[31]
The presently reached 137 ppm at 356K is obviously not low enough. Therefore, it
will be necessary to insert measures in order to prevent accumulation of fluoride in
process liquors. Two possibilities should be investigated in this respect:
1. Optimisation of precipitation conditions in order to reach a lower fluoride level.
2. The insert of an additional fluoride removal step.
In Chapter 3, both aspects are extensively studied. In Section 3.2, a thermodynamic
model is used to predict the minimum fluoride level that can be reached by optimising
precipitation conditions; in Section 3.3, additional fluoride removal steps are
identified based on both a literature review as well as on verifying experiments.
Figure 2.4 shows that filtration improves when the precipitate is produced at higher
temperature, although there does not seem to be a difference (over this filter) between
precipitates produced at 331K and 356K. The improved filtration behaviour is
probably caused by increased particle diameter.
Similar approaches to remove magnesium from zinc sulphate solution or zinc vitriol
are described in Russian literature. Shmurygina et al.[33] propose to feed zinc fluoride
pulp to magnesium containing zinc sulphate solution, resulting in the precipitation of
magnesium fluoride, which was believed to be applicable in the instrument industry
(after washing). It is not clear, however, what is meant with this remark. Furthermore,
the used zinc electrolyte had an initial fluoride content of 100 ppm and after
precipitation a final fluoride content of 300 ppm. The consequences of such a high
residual fluoride content are not discussed. Kurdyomov et al.[34],[35] also proposed
33
magnesium fluoride precipitation by addition of zinc fluoride slurry. The following
precipitation conditions are recommended, which are in reasonable accordance with
the results of the precipitation experiments described above:
•
•
•
temperature between 40-60ºC
reaction time 2-3 hours
75% magnesium removal
The importance of minimising the residual fluoride content was recognised. However,
figures on the residual fluoride concentration of the solution were not provided. The
magnesium fluoride product was believed to be applicable as smelt component and in
the preparation of enamels, glasses and glazes. Fishman et al.[36],[37] suggest to slowly
add a magnesium containing zinc sulphate solution to a stirred zinc fluoride
suspension at a temperature of 50-60ºC. After 1 hour, a filterable magnesium fluoride
residue is obtained. In the above mentioned treatment methods, zinc fluoride pulp was
obtained by contacting zinc sulphate solution or zinc oxide with hydrofluoric acid.
Bulakhova et al.[38] suggest to replace hydrofluoric acid by ammonium bifluoride for
safety reasons and in order to reduce reagent costs.
2.4.2 Conversion
In order to be saleable, the magnesium hydroxide probably has to be more pure than
given by the results of Section 2.3.2. There are a number of possible explanations for
the remaining fluoride. During the conversion and the accompanying hydroxide
consumption the pH might have dropped below the conversion-pH (of
MgF2/Mg(OH)2 transition). Another possibility is that the particles were too large and
that the cores of the particles were never contacted with sodium hydroxide solution.
Therefore, care must be taken to keep magnesium fluoride particles in suspension.
Finally fluoride ions could have been structurally incorporated into the crystal
structure. Additional research on this subject is described in Chapter 4. The fluoride
content of the converted product in the 8 h. experiment is equal to the fluoride content
of the 4 h. experiment, indicating that slow kinetics did not affect the experimental
results.
2.4.3 Electrolysis
It may be concluded from the experimental results that ion transport has occurred and
that the membrane process basically works. It was shown that the molar Zn:F ratio is
not equal to 1:2 as would be expected. The amount of fluoride found in the anolyte
was much lower than expected. An explanation could be loading of the anion
exchange membrane. It should be realised that in this experiment the sodium
hydroxide concentration in the catholyte is rather low. Defining optimal operating
conditions required additional research on this subject, which is described in Chapter
5.
34
2.4.4 Proposed flowsheet
As experiments concerning the different process steps showed promising results, a
first attempt has been made to draw up a flowsheet. Figure 2.8 shows the integration
of different process steps. At this stage, it is assumed that the water input of the
membrane cell reactor is balanced by evaporation losses in the rest of the zinc plant
(otherwise the process liquor would progressively become diluted).
ZnF2 solution
C
purified
electrolyte
e
F- Na+
ZnF2
MgF2
precipitation
NaOH
NaF
zinc
NaOH
to
plant
L
S
MgF2
L
MgF2/Mg(OH)2
conversion
S
water
Mg(OH)2
Figure 2.8 Suggested flowsheet for the magnesium removal process.
2.5
Conclusions
The results discussed in this chapter make the following clear:
•
Selective precipitation of magnesium fluoride out of purified zinc solution is
possible by simultaneous slow addition of purified zinc sulphate solution and zinc
fluoride solution.
•
The residual fluoride content of the filtrate decreases with increasing precipitation
temperature, however, the residual fluoride concentrations could be too high for
industrial application.
•
Filtration characteristics improve as well with increasing temperature, probably
due to the increase of the particle diameter.
•
Magnesium fluoride can be converted to magnesium hydroxide by contacting the
precipitate with sodium hydroxide solution. However, the residual fluoride
content of the product is rather high.
35
•
The resulting sodium fluoride solution can be regenerated in a membrane cell into
zinc fluoride and a sodium hydroxide solutions, which both can be recycled in the
process.
Summarising, the principles of the proposed process seem to be valid. The desired
chemical reactions in the three individual unit processes have been realised. However,
the purity of (intermediate) products, especially with regard to the residual fluoride
content, is still insufficient.
It is the objective of the subsequent three chapters to discuss the purity of the
products. This also in view of creating a sustainable solution to zinc processing as
discussed in Chapter 1. Therefore, each unit process will be extensively studied in the
subsequent three chapters and possible solutions to deal with residual fluoride will be
presented.
36
3
SELECTIVE PRECIPITATION OF A
MAGNESIUM FLUORIDE PRODUCT
Chapter 3 deals with the selective precipitation of magnesium fluoride. First, a short
general thermodynamic introduction precedes the detailed discussion of the
construction of a thermodynamic model capable of predicting fluoride speciation in
zinc sulphate solutions at high ionic strength. This thermodynamic model has been
published in Minerals Engineering. Model predictions are compared with
experimental results. Thereafter, possible methods to treat fluoride containing zinc
sulphate solutions are discussed. The chapter concludes with a detailed study of the
magnesium fluoride precipitation product, which has been published in Powder
Diffraction.[62]
3.1
Introduction to precipitation processes
Although this thesis by no means intends to provide a complete thermodynamic or
kinetic description of precipitation processes, it was thought useful to give some
theoretical background based on a few standard textbooks.[39],[40],[41],[42],[43]
If we consider a system at constant pressure and temperature and the only work being
done is mass transfer, we can write the following equation for the Gibbs energy
change:
dG =
 ∂G 

dni =
 i  p ,T , n 'i
∑  ∂n
i
∑ µ dn
i
i
(3.1.1)
i
The thermodynamic potential, µi, represents the reversible work necessary to add a
mole of substance i to an (infinite amount of) mixture. Chemical reactions and phase
transitions can be considered as the addition and removal of phases. For a phase
transition, such as a precipitation reaction, at constant pressure and temperature, we
can write:
∆G = (µ2-µ1)∆n
with
µ2-µ1 = ∆µ
(3.1.2)
The negative of this quantity, φ=-∆µ, is actually the driving force for crystallisation
and is called the reaction affinity. For a spontaneous process φ > 0 and at equilibrium
φ = 0. If state 1 corresponds to a supersaturated solution and state 2 to the crystal:
∆µ = µ2-µ1 = (µ0 + RT ln a2) – (µ0 + RT ln a1)
(3.1.3)
which can also be written as:
φ/RT = ln (a1/a2)
(3.1.4)
The activity of a component is in a real solution not equal to its concentration.
Therefore, an activity coefficient is introduced to express the deviation from ideality.
For molal units, the following relation holds:
am = mγm
(3.1.5)
If the dissolved component is an electrolyte that dissociates in solution to give
ν’+ cations and ν’- anions, introduction of the mean activity is necessary as it is
impossible to determine the activity coefficients of the individual ions:
am = aν’±, m = (Qmγ±,m)ν’
(3.1.6)
Activities can also be expressed using other concentration scales such as molar
concentrations or mole fractions.
From electrostatic theory, the Debye-Hückel equation can be deduced, which is an
expression for the activity coefficient. However, this equation is only applicable at
very low ionic strength. The ionic strength of a solution is defined as follows:
I = 1/2∑cizi2
(3.1.7)
The Debye-Hückel equation is the only equation, which can be fundamentally
deduced from electrostatic theory; other available equations are (semi)-empirical. A
few examples, including the Debye-Hückel equation, are given below
•
•
•
Debye-Hückel: log γmean,c = -Azi2I1/2/(1+aiBI1/2)
I<0.1M
2 1/2
1/2
Güntelberg:
log γmean,c = -Azi I /(1+I )
I<0.1M
2 1/2
1/2
Davies:
log γmean,c = -Azi I /(1+aI ) + bI I<0.7M
A, B, a, b and ai are constants (some of which are T-dependent)
For more concentrated solutions, there are other methods, e.g. the Bromley method or
methods according to Guggenheim, Meissner or Pitzer may be used. However, the
complex composition of the mixed metal-sulphate solutions we are dealing with in
electrolytic zinc industry makes the calculation of individual activity coefficients
tedious and there is an increasing risk of introducing errors in these calculations.
Therefore, following many other authors, another method is followed, which will be
described in detail in Section 3.2.
In precipitation systems, the (dimensionless) supersaturation ratio is defined as:
S = amean,1/ amean,2
(3.1.8)
In a supersaturated solution, precipitation starts with the formation of nuclei: the
capture of g new growth units reduces the Gibbs energy according to
∆Gv = -∆n*φ = -g/Nav*RT ln S = -gkT ln S
(3.1.9)
However, the formation of a nucleus involves the creation of a solid-solution interface
with a finite surface energy, resulting in an increase in the Gibbs energy according to
38
∆Gs = 4πr2γs for a spherical nucleus
(3.1.10)
Combination of these two equations yields a curve that can be interpreted as an
energy barrier for nucleation. The critical radius depends on the supersaturation
(higher supersaturation results in lowering of the energy barrier). An expression for
the critical radius as a function of the supersaturation ration can be derived and
substituted in eqs. 3.1.9 and 3.1.10, which yields upon rearranging:
∆Ghom =
16πv 2 (γ s ) 3
3( kT ln S ) 2
(3.1.11)
This expression may be regarded as the activation energy for a homogeneous
nucleation process. It will be clear that the supersaturation ratio in the magnesium
fluoride experiments described in Chapter 2 was high enough to overcome this energy
barrier. In many cases, however, heterogeneous instead of homogeneous nucleation
occurs. When foreign particles are already present in solution, the interfacial energy
between the precipitating compound and these foreign particles may be much smaller
than the interfacial energy between the crystal and the solution. In that case,
precipitation takes place onto the surface of the foreign particles at a much lower
supersaturation. In other words, ∆Ghet, has been dramatically reduced. This can be
advantageous when the filtration characteristics have to be improved. Such a
supersaturation can be chosen that homogenous nucleation does not occur, instead the
only precipitation is taking place on foreign surface. This way, large particles can be
produced which are easy to filter. Instead of foreign particles, seed material of the
precipitating compound itself could be used. In fact, this is common practice in many
industrial precipitation processes and although the filterability in the magnesium
fluoride precipitation experiments described in Chapter 2 appears to be sufficient, it
could also be applied in the proposed magnesium removal process.
The activation energy for the nucleation of magnesium fluoride has been
experimentally determined by Kurdyomov et al.[34] at 32 kJ/mol although this
activation energy was determined in an experiment in which magnesium fluoride
precipitation was brought about by addition of solid zinc fluoride to zinc electrolyte
solution. Therefore, apart from homogeneous nucleation also heterogeneous
nucleation at the zinc fluoride surface was possible and thus the determined activation
energy reflects a combination of homogeneous and heterogeneous nucleation. In their
paper, Kurdyomov et al.[34] recommend the following precipitation conditions, which
are in accordance with the results of the precipitation experiments described in
Chapter 2:
•
•
•
temperature between 40-60ºC
reaction time 2-3 hours
75% magnesium removal
The recommendation of 75% magnesium removal was directly related to the
observation that at higher removal percentages, the residual fluoride content of the
solution markedly increased. In Section 3.2, a thermodynamic model is described,
which predicts an optimal magnesium removal percentage of 72%. This percentage
obviously corresponds quite well with the recommendation from experimental work.
39
However, the results are not entirely comparable as Kurdyomov et al.[34] used zinc
sulphate solutions containing appreciable amounts of calcium resulting in the
precipitation of calcium fluoride.
3.2
Thermodynamic modelling of magnesium fluoride
precipitation in concentrated zinc sulphate environment
(Section 3.2 is largely adopted from a publication in Minerals Engineering [44])
3.2.1 Formulation of the thermodynamic model
In Chapter 2, from experiments (as well as from literature sources) it became clear
that the residual fluoride content of solutions after magnesium fluoride precipitation
was rather high. The question came up whether a thermodynamic equilibrium was
reached or that slow kinetics is responsible for the high residual fluoride content.
Therefore, it is useful to determine first what is thermodynamically possible.
Subsequently, experimental results are compared with the thermodynamic equilibrium
situation in order to determine if kinetics is of significant interest. If it turns out to be
thermodynamically possible to reach a sufficiently low residual fluoride content, the
following step would be to optimise precipitation conditions in order to improve
kinetics. If thermodynamic modelling indicates that it is not possible to reach a low
residual fluoride content, it will be decided not to further study the details of the
kinetics of this reaction. Instead, research effort will be put in the search for an
additional fluoride removal method. A thermodynamic equilibrium model has been
formulated, firstly, as a prediction tool for fluoride behaviour in magnesium fluoride
precipitation experiments with purified zinc sulphate solutions and secondly, as a
basis for fluoride behaviour in other parts of electrolytic zinc plants.
Mass balance equations for the major constituents of purified zinc sulphate solutions:
zinc, fluorine, magnesium, manganese and sulphur are considered. In addition to these
elements, traces of other elements are present. When thermodynamic stability
constants of the fluoride complexes of these elements are of the same magnitude as
thermodynamic stability constants of the zinc fluoride and magnesium fluoride
complex, it is legitimate to neglect these species. However, the fluoride complexes of
iron and aluminium are very strong and should be incorporated into the model.
For iron, this remark has to be refined as only ferric iron ions form very strongly
complexes with fluoride ions. During purification, cementation with zinc dust is
common and any Fe(III) still present in the zinc sulphate solution will be reduced:
Zn + 2Fe3+ → Zn2+ + 2Fe2+
(3.2.1)
At equilibrium, Nernst law can be applied:
E 0 ( Zn 2 + / Zn) +
RT
2F
ln a ( Zn 2 + ) = E 0 ( Fe 3+ / Fe 2+ ) +
RT
F
ln a ( Fe 3+ ) / a (Fe 2 + )
(3.2.2)
40
Substituting
E0 (Zn2+/Zn) = -0.76 V
E0 (Fe3+/Fe2+) = 0.77 V
F = 96 485 C mol-1
R = 8.314 J K-1 mol-1
T = 298K
yields upon rearranging
ln
a (Zn 2 + )∗ a (Fe 2 + ) 2
a (Fe 3+ ) 2
= 119.2
(3.2.3)
Therefore, ferric iron will not be present and it seems legitimate to leave the ferric
iron balance out of the thermodynamic model. However, it should be realised that
aeration of the solution prior to the magnesium removal process results in oxidation of
ferrous iron (if present). Subsequently, a higher residual fluoride content will be
reached during the selective precipitation of magnesium fluoride.
Manganese acts the same way as iron, Mn(III) ions form very strong complexes with
fluoride ions. However, Mn(III) ions which are formed during electrolysis are
effectively reduced by ferrous irons in the leaching part of the plant (common practice
in zinc plants in order to prevent corrosion of the heat exchangers). Therefore, Mn(III)
is not considered in this model calculation.
One might assume that siliceous species should be incorporated as well, since the
interaction of fluoride solutions with glassware is notorious. Unfortunately, the
standard reference handbook for stability constants of Smith and Martell[45] does not
provide information on such fluoro-siliceous species. The well-known
thermodynamic software package HSC Outokumpu® does provide a stability
constant for the reaction:
SiO2(aq) +6F- + 4H+ à SiF62- + 2H2O
K = 2.2·1029 (at 298K)
(3.2.4)
At pH=4 and a free fluoride concentration of 2 mM (38 ppm), SiO2(aq) is by far the
dominant species. It will be shown that the equilibrium free fluoride concentration is
lower than 2 mM. Therefore, we can neglect the silicon mass balance in the solution
model for pH=4.
Various authors have discussed the thermodynamic description of relatively
concentrated mixed metal sulphate systems, among them Wang & Dreisinger[46] and
Filippou et al.[47]. Wang and Dreisinger[46] argued that the dinuclear species
Zn2(OH)3+ influences the pH-temperature behaviour of the solution very strongly.
Furthermore, they noticed that Zn(OH)20 was negligible (<10-6M) in their solution
system of pH<5 and they excluded it from the model (along with other species such as
Zn(OH)3- and Zn(OH)4-, which occur only at higher pH). In this model, the choice of
species of Wang and Dreisinger[46] is largely followed, with the exception that sodium
is excluded from the model as the sodium content of the purified zinc sulphate
solutions of Pasminco Budel Zink, which were used in the magnesium fluoride
41
precipitation experiments, was negligible. Furthermore, Al(OH)2+ was included as
well as, logically, various fluoride species.
mass balance equations
[Zn]T =
[Zn2+] + [ZnF+] + [ZnOH+] + 2[Zn2(OH)3+] + [ZnSO40]
[Mg]T =
[Mg2+] + [MgF+] + [MgOH+] + [MgSO40] + {MgF2}solid
[F]T
=
[MgF+] + [ZnF+] + [F-] + [MnF+] + [HF] + 2{MgF2}solid + 2[HF2-] +
[AlF2+] + 2[AlF2+] + 3[AlF30] + 4[AlF4-]
[S]T
=
[SO42-] + [HSO4-] + [ZnSO40] + [MnSO40] + [MgSO40] + [Al(SO4)+] +
2[Al(SO4)2-]
[Mn]T =
[Mn2+] + [MnOH+] + [MnSO40] + [MnF+]
[Al]T =
[Al3+] + [AlOH2+] + [Al(OH)2+] + [AlSO4+] + [Al(SO4)2-] +
2[Al2(OH)24+] + 3 [Al3(OH)45+] + [AlF2+] + [AlF2+] + [AlF30] + [AlF4-]
An electroneutrality condition is required.
electroneutrality condition
2[Mg2+] + [MgF+] + [MgOH+] + [ZnF+] +2[Zn2+] + [Zn(OH)+] + 3[Zn2(OH)3+] +
2[AlF2+] + 3[Al3+] + [AlF2+] + [H+] + 2[AlOH2+] + [Al(OH)2+] + 4[Al2(OH)24+] +
5[Al3(OH)45+] + [AlSO4+] + 2[Mn2+] + [MnOH+] + [MnF+] = [F-] + [HF2-] + [OH-] +
2[SO42-] + [AlF4-] + [Al(SO4)2-] + [HSO4-]
Finally, a number of thermodynamic equations complete the description of the
precipitation system. Substitution of these thermodynamic equations in the mass
balance equations reduces the total number of equations to be solved.
thermodynamic equilibria
KcT,1 = [Mg2+][F-]2
KcT,3 = [ZnF+]/[Zn2+][F-]
KcT,5 = [Zn2(OH)3+]/[Zn2+]2[OH-]
KcT,7 = [MgSO40]/[Mg2+][SO42-]
KcT,9 = [AlF2+]/[Al3+][F-]2
KcT,11 = [AlF4-]/[Al3+][F-]4
KcT,13 = [Al(OH)2+]/[Al3+][OH-]2
KcT,15 = [Al3(OH)45+]/[Al3+]3[OH-]4
KcT,17 = [Al(SO4)2-]/[AlSO4+][SO42-]
KcT,19 = [MnOH+]/[Mn2+][OH-]
KcT,21 = [MgOH+]/[Mg2+][OH-]
KcT,23 = [HF]/[H+][F-]
KcT,25 = [HSO4-]/[[H+][SO42-]
KcT,2
KcT,4
KcT,6
KcT,8
KcT,10
KcT,12
KcT,14
KcT,16
KcT,18
KcT,20
KcT,22
KcT,24
= [MgF+]/[Mg2+][F-]
= [Zn(OH)+]/[Zn2+][OH-]
= [ZnSO40]/[Zn2+][SO42-]
= [AlF2+]/[Al3+][F-]
= [AlF30]/[Al3+][F-]3
= [AlOH2+]/[Al3+][OH-]
= [Al2(OH)24+]/[Al3+]2[OH-]2
= [AlSO4+]/[Al3+][SO42-]
= [MnSO40]/[Mn2+][SO42-]
= [MnF+]/[Mn2+][F-]
= 1/[H+][OH-]
= [HF2-]/[HF][F-]
Input values for the mass balance equations are based upon analyses of industrial
purified zinc sulphate solution and analytical zinc fluoride solution, which were used
in the precipitation experiments.
42
3.2.2 Determination of mass stability constants
The high ionic strengths prevailing in industrial zinc sulphate solutions make it
necessary to distinguish between molar concentrations and activities. Instead of
estimating individual activity coefficients of species, Fillippou et al.[47] used
Vasil’evs[48] modification of the extended Debye-Hückel equation in order to estimate
the mass stability constants from thermodynamic constants at ionic strength
considerably higher than zero:
log KcT , i = log KcT , i +
A∆z 2 Ic
0
1 + 1.6 Ic
+ bIc
(3.2.5)
where
•
•
•
•
KcT,i is the molar mass stability constant as a product of concentrations
KcT,i0 is the thermodynamic equilibrium constant
A is the temperature dependent Debye-Hückel coefficient
n
Ic is the molar ionic strength, I c = 0.5 * ∑i =1 ci zi2
•
∆z 2 =
•
b is a constant that can be estimated by plotting q against Ic
•
∑
q ≡ log(
n
i =1
ν i zi2
KcT ,i
KcT ,i
0
)−
A∆z 2 I c
1 + 1.6 I c
= bI c
(3.2.6)
Values for b were taken from Fillippou et al.[47] or computed using tables of stability
constants of Smith and Martell (1976)[45] and Martell and Smith (1982)[49]. Plots of q
against Ic generally resulted in equations of straight lines (with the slope being the
parameter b). In Table 3.1, parameters b for different species are listed as well as the
accompanying regression coefficients (r2) when the parameters were computed in this
paper from the tables of Smith and Martell[45]. When the regression coefficient is
unity, the equation is based only on two points, and we do not know anything about
the accuracy of the estimated parameter.
Due to lack of data, the b parameter could not be determined for the solubility product
of magnesium fluoride. Furthermore, the Vasil’ev plots of MgSO40 and ZnSO40
yielded unsatisfactory results. Therefore, the molal mass stability constants of
ZnSO40, MgSO40 at high ionic strength are determined from the thermodynamic
stability constants using a Davies type of equation (following Fillippou et al.[47]):
log KmT = log KmT + A∆z 2 (
Im
0
1+
Im
) − 0.2 I m
(3.2.7)
The molar solubility product of magnesium fluoride at high ionic strength is
determined with a similar type of equation.
43
Table 3.1 Molar stability constants (at zero ionic strength) for complex
equilibria and parameters b estimated by constructing Vasil’evs plots.
b
r2
Species
log Kc2980
-8.18
Mg2+ + 2F-→ MgF2(S)
2+
+
1.8
0.2833
0.9587
Mg + F → MgF
1.15
0.2118
0.9512
Zn2+ + F-→ ZnF+
0
22+
2.38
Zn + SO4 → ZnSO4
0
22+
2.23
Mg + SO4 → MgSO4
7.0
0.2452
0.9611
Al3+ + F-→ AlF2+
12.6
0.3872
0.9079
Al3+ + 2F-→ AlF2+
3+
16.7
0.6532
0.9958
Al + 3F → AlF3
19.1
1.6532
0.8768
Al3+ + 4F-→ AlF43+
2+
9.01
0.4501
0.981
Al + OH → AlOH
+
3+
18.7
0.2719
1
Al + 2OH → Al(OH)2
4+
3+
20.3
1.2912
1
2Al + 2OH → Al2(OH)2
5+
3+
42.1
1.579
1
3Al + 4OH → Al3(OH)4
1
3+
2+
3.89
0.0656
0.8541
Al + SO4 → AlSO4
2+
22.26
0.1147
0.8736
Mn + SO4 → MnSO4
3.4
0.5148
0.9833
Mn2+ + OH-→ MnOH+
2
2+
+
1.3
0.2148
0.9979
Mn + F → MnF
1
2+
+
2.8
0.1129
1
Mg + OH → MgOH
+
13.997
0.219
H + OH → H2O
+
3.17
0.2084
0.9934
H + F → HF
1
0.57
0.2226
0.9174
HF + F → HF2
[45]
All constants taken from Smith and Martell , except the numbers accompanied by a
1
footnote:
Martell and Smith[49]
2
Morel and Hering[50]
The molal mass stability constant (of Eq. 3.2.7) can be converted to molar scale by
Eq. 3.2.8:
KcT =
KmT
n
ρ f − 0.001∑ ci M i
(3.2.8)
i =1
The molal ionic strength Im is defined as:
I m = 0.5 * ∑i =1 mi zi2
n
(3.2.9)
in which mi represents the molality of species i in mol kg-1
mi =
ci
n
ρ f − 0.001∑ ci M i
(3.2.10)
i =1
44
The density of the solution can be estimated using empirical relations that combine
tabulated density data of binary solutions. However, in this case, measurement of the
density of the solution was considered more appropriate. The density of the filtered
solution was determined, i.e. without solid magnesium fluoride particles. Therefore,
in Eqs. 3.2.8 and 3.2.10, magnesium fluoride is excluded from the sum.
The thermodynamic dissociation constant for the reaction
Al(SO4)2- → AlSO4+ + SO42-
log KdmT,17 = -1.89
(3.2.11)
is taken from Wang & Dreisinger[46]. As this constant is reported for the condition of
zero ionic strength, conversion from molal scale to molar scale is not necessary. Due
to lack of further information, this constant is also used as mass stability constant at
ionic strengths much higher than zero.
For the determination of the mass stability constant of the bisulphate ion, the work of
Dickson et al.[51] was used instead of the extended Debye-Hückel equation. Dickson
and his co-workers suggested an equation for the dissociation constant of HSO4depending on temperature and ionic strength on molal basis:
log KdmT25 = p1 + p2/T + p3 ln T + p4T + p5T2 – 4fr/ln 10 + p6(Im/T) +
F(Im)(p7T + p8/T) + p9(Im2/T)
with
KdmT25
T
Im
fr
F(Im)
Aф
p1-p9
(3.2.12)
= dissociation constant of HSO 4- on molal basis
= absolute temperature
= molal ionic strength
= -Aф{√Im/(1+1.2√Im) + 2 ln (1 + 1.2√Im)/1.2}
= 1 – (1 + 2√Im)exp(-2√Im)
= 0.322863 + 3.75692 x 10 -4t + 2.55932 x 10-6t2 –9.96273 x 10-11t3
+ 5.98066 x 10-3exp((t-270)/10) + 1.39987 x 10 -2/(t+20) + 18.4374/(315-t) –
554.596/(315-t) 2 + 7684.77/(315-t)3- 54091/(315-t) 4 + 154381/(315-t)5
with t = temperature in ºC
= experimentally determined parameters (Dickson et al., 1990) [51]
Table 3.2 Dickson’s parameters.
p1 = 562.7097
p2 = -13,273.75
p3 = -102.5154
p4 = 0.2477538
p5 = -1.117033 x 10-4
Experimental parameters
p6 = -57.07583
p7 = -1.144759 x 10-3
p8 = 46.72816
p9 = 2.499849
The molal mass stability constant (mi =1/ KdmT25) can be related to the molar mass
stability constant with Eq. 3.2.10.
In order to extend the model to higher temperatures, use was made of the
thermodynamic database of the software package Geochemical Workbench®. As the
tabulated thermodynamic stability constants for 298K correspond with the stability
45
constants derived from Smith and Martell[45], it was considered legitimate to use the
reported thermodynamic stability constants at 333K and 373K (Table 3.3).
Table 3.3 Thermodynamic stability constants derived from Geochemical
Workbench.
species
Mg2+ + 2F-→ MgF2(S)
Mg2+ + F-→ MgF+
Zn2+ + F-→ ZnF+
Zn2+ + SO42- → ZnSO40
Mg2+ + SO42- → MgSO40
Al3+ + F-→ AlF2+
Al3+ + 2F-→ AlF2+
Al3+ + 3F-→ AlF3
Al3+ + 4F-→ AlF4Al3+ + OH-→ AlOH2+
Al3+ + 2OH-→ Al(OH)2+
2Al3+ + 2OH-→ Al2(OH)24+
3Al3+ + 4OH-→ Al3(OH)45+
Al3+ + SO42-→ AlSO4+
AlSO4+ + SO42- → Al(SO4)2Mn2+ + SO42-→ MnSO4
Mn2+ + OH-→ MnOH+
Mn2+ + F-→ MnF+
Mg2+ + OH-→ MgOH+
H+ + OH-→ H2O
H+ + F-→ HF
HF + F-→ HF2-
logK3730
-9.87
2.52
2.01
3.08
2.59
7.38
13.35
17.85
19.96
8.74
17.18
19.31
3.82
1.30
2.99
3.58
2.04
2.80
12.24
3.86
0.51
logK3330
-8.62
2.11
1.59
2.68
2.39
7.16
12.90
17.06
19.45
8.92
17.66
19.77
3.30
1.79
2.60
3.45
1.66
2.46
13.02
3.47
0.54
log K2980
-8.18
1.80
1.25
2.37
2.23
7.00
12.60
16.70
19.10
9.05
17.87
20.30
42.09
3.01
1.89
2.29
3.40
1.34
2.20
13.987
3.17
0.50
In this section, only the temperatures 333K and 373K are considered. When the
thermodynamic stability constants are required at other temperatures, the Gibbs free
energy of reaction can be calculated when the standard entropy change of reaction is
known as well as heat capacity functions for all reactants and products. When heat
capacity functions are not available, Helgeson’s method can be used to obtain a good
estimate of the thermodynamic stability constant (Helgeson, 1967[52] and Fillippou et
al., 1995[47]). This was done for the stability constants of ZnOH+ and Zn2OH3+ at
333K and 373K:
log K T =
with
∆S T0
0
2.303RT
[T0 −
θ
ω
(1 − exp(exp{b + aT } − c +
∆S T0
= the standard reaction entropy change
∆H
= the standard reaction enthalpy change
0
T
R
θ
0
T0
T − T0
θ
))] −
∆H T0
0
2.303RT
(3.2.13)
= reaction temperature (K)
= 8.314 J K-1 mol-1
= 219K
46
ω
a
b
c
= 1.00322K
= 0.01875K -1
= -12.741
= 7.84x10 -4
In order to be consistent with the thermodynamic constants calculated at higher
temperature, the thermodynamic constants at 298K for both species were also
calculated according to this method. Entropy and enthalpy values, summarised in
Table 3.4, were selected from the papers of Wang and Dreisinger[46] and Fillippou et
al.[47] Dependence on ionic strength was accounted for by using Vasil’ev’s equation
for ZnOH+ (b parameter –0.087 from Fillippou et al.).
Table 3.4 Entropy and enthalpy data.
Species
ZnOH+
Zn2OH3+
OHZn2+
ΔST00 (J K-1 mol-1)
81.6
81.6
-10.88
-112.1
ΔGT00 (kJ mol-1)
-330.1
-488.44
-157.29
-147.06
At this point, the model can be solved in an iterative procedure:
•
•
•
•
•
First, estimate Ic and Im
calculate all molar mass stability constants
solve the model (with a suitable mathematical computer program, in this case
Mathcad®)
calculate new values for Ic and Im
If the values are within an acceptable range (<10-3) of the estimated Ic and Im the
iterations stop
3.2.3 Results and discussion of the thermodynamic model
Industrial zinc sulphate solution is assumed to have the composition as it is
summarised in Table 3.5.
Table 3.5 Electrolyte composition.
model composition
Zn
Mg
Mn
Al
SO4
144 g/l
13.7 g/l
5.1 g/l
1 ppm
274.5 g/l
Input values for the mass balance equations are obtained by simulating industrial
electrolyte:ZnF2 solution mixing in a ratio 1:2.75. The model was solved for different
47
percentages of magnesium removal at 298K. This was accomplished by varying the
input fluorine and zinc mass balances. Variation of pH was realised by solving the
model for increasing input values for the sulphate mass balance. In Figure 3.1, the
total residual fluoride content is depicted as a function of pH at different magnesium
removal percentages. From the figure, it becomes clear that magnesium fluoride
precipitation should be performed between pH 4-4.5. At lower pH values, the total
residual fluoride content increases rapidly; at higher pH values, unwanted zinc
hydroxide precipitation can be expected[53].
total residual F (ppm)
350
300
250
13%
55%
77%
90%
200
150
100
50
0
1
2
3
4
5
pH
Figure 3.1 Residual fluoride at different percentages of Mg removal at 298K.
It will be clear that the residual fluoride content of the solution decreases when the
amount of removed magnesium is decreased. However, when a small fraction is
removed, obviously a larger volume of solution will be produced in order to reach the
same amount of magnesium removal. Therefore, the product of total filtrate volume
and residual fluoride content (ppm) should be considered in order to determine the
optimal fraction (i.e. at minimised fluoride loss) of magnesium to be removed. The
residual fluoride content in the solution has been calculated for different fractions of
magnesium removal from typical industrial zinc sulphate solutions (containing 13.7
g/l Mg). Subsequently, the total amount of filtrate solution was calculated for the
removal of 13.7 grams of magnesium at these different removal percentages. An
optimal magnesium removal percentage could be determined at 72% (Figure 3.2).
48
(1/removed fraction)*residual F
180
160
140
120
100
80
60
40
20
0
0
20
40
60
80
100
% magnesium removed
Figure 3.2 Optimal magnesium removal percentage based on total residual fluoride content.
In Figure 3.3, the residual fluoride content is depicted at 55% magnesium removal.
The figure clearly shows that the increase in residual fluoride with decreasing pH is
solely attributable to increasing HF formation (and a small amount of HF2-). It is
worthwhile noticing that aluminium fluoride complexes make up an amount of 2.5
ppm, while the aluminium input was only 1 ppm. One should realise that any
dissolved aluminium in this precipitating system will be present as aluminium
fluoride complex. Generally, industrial zinc sulphate solutions will not contain
dissolved aluminium, but when it occasionally occurs, the fluoride content of the
solution increases very fast.
residual fluoride (ppm)
350
total F
300
250
HF + HF2-
200
150
100
F-, ZnF+,
MgF+,
MnF+, Alfluorides
50
0
1
2
3
4
5
pH
Figure 3.3 Residual fluoride as a function of pH at 55% Mg removal at 298K.
For this situation, the total free fluoride concentration at pH=4 is 32 ppm. Recalling
the equilibrium constant of the following equilibrium:
SiO2(aq) +6F- + 4H+ → SiF62- + 2H2O
K = 2.2·1029 (at 298K)
(3.2.4)
49
[SiF62-] = 4.2·10-4[SiO2]
results in:
Therefore, operation in the pH range 4-4.5 will not result in fluoride loss to fluorosiliceous species. At lower pH levels, both the fluoride and free acid concentration
increase resulting in possible fluoride complexation when dissolved siliceous species
are present (which is usually not the case).
The model is solved for the higher temperatures 333K and 373K as well. The
equilibrium situation is depicted for three different temperatures in Figure 3.4. A
number of different features can be distinguished. With increasing temperature, the
residual free fluoride content of the magnesium fluoride precipitating system
decreases. Unfortunately, stability constants of fluoride complexes increase with
increasing temperature. Therefore, the reduction in residual free fluoride content
between 298K and 333K is balanced by the increase in fluoride complexes. At 373K,
the free fluoride content is so low that even the increased values of stability constants
can not balance it anymore. Therefore, magnesium fluoride precipitation near the
boiling point of the solution seems optimal for minimising residual fluoride.
Furthermore, the equilibrium pH decreases with increasing temperature as a result of
enhanced zinc hydrolysis at higher temperature. Consequently, the concentration of
HF increases with increasing temperature.
6
70
5
60
4
50
3
40
30
pH
Residual fluoride (ppm)
80
2
20
total F
F
HF
other F
pH
1
10
0
273
0
293
313
333
353
373
Temperature (K)
Figure 3.4 Residual fluoride as a function of reaction temperature.
The model calculations make clear that the required residual fluoride content of less
than 10 ppm, will not be reached under any set of pH-T conditions. If no suitable sink
for fluoride can be identified in the zinc electrolysis process, an additional fluoride
removal step has to be developed.
Experimental results (Booster et al., 2000[44] and this thesis, 2.3.1) are compared with
the predictions of the thermodynamic model. Magnesium fluoride precipitation
experiments were conducted at different temperatures. During the precipitation
experiments, 0.12M ZnF2 solution (pH=5.1) and industrial zinc sulphate solution
(composition as given in Table 5) were mixed over a period of 45 minutes. This
resulted in a magnesium removal efficiency of approximately 56%. After the addition,
50
the suspension was stirred (200 rpm) for another 75 minutes. The residual fluoride
content, depicted as a function of reaction temperature in Figure 3.4, was determined
with the specific fluoride ion electrode.
During the experiments, the pH of the suspensions generally dropped to pH≅3. Model
calculations for pH=3 are also depicted in Figure 3.5. As an example of the sensitivity
of the model to uncertainties, the thermodynamic constants of the most important
equilibria (i.e. MgF2, ZnSO4, MgSO4, MgF+ and ZnF+) are varied by a factor of two.
This approach yields minimum and maximum error bars as indicated.
residual fluoride (ppm)
As the difference between experiments and thermodynamics is largest at 298K, it was
questioned whether thermodynamic equilibrium was reached. Therefore, the timedependence of the magnesium fluoride precipitation reaction was investigated in a
long-term experiment. As described, normally, the precipitation experiments were
carried out in a period of 2 hours. In the long-term experiment, the addition of
solutions was also completed in 45 minutes. Subsequently, the solution was stirred for
48 hours, during which samples were taken. Figure 3.6 depicts the residual fluoride
content of the filtrate as a function of time. Figure 3.6 indicates that an equilibrium
value around 250 ppm is reached within 20 hours. This value is also depicted in
Figure 3.5.
400
350
300
250
200
150
100
50
0
273
experimental
thermodynamic
model
323
373
temperature (K)
Figure 3.5 Residual fluoride as a function of reaction temperature.
For a temperature of 298K, the experimental results suggest that the residual fluoride
content predicted by the thermodynamic model is slightly too low. If this would also
be the case at higher temperature, then experimental results would even be closer to
the thermodynamic equilibrium situation as they are now. Therefore, the precipitation
reaction does not seem to be very much affected by slow kinetics. Besides, as both
thermodynamic modelling as well as experimental results indicate that it is not
possible to reach a low residual fluoride content, it was decided not to study the
details of the kinetics of this reaction. In Section 3.3, research effort was put in the
search for an additional fluoride removal method.
51
residual fluoride content (ppm)
400
350
300
250
200
150
100
50
0
0
10
20
30
40
50
60
time (hours)
Figure 3.6 Residual fluoride content as a function of time.
In addition, the reversed approach was followed. Two grams of (dry) magnesium
fluoride precipitate, produced according to the method described in Chapter 2, was
contacted with 350 ml industrial electrolyte. After a stirring time of 24 hours, the
fluoride content of the solution was measured with the specific fluoride ion electrode:
250 ppm. One could argue that the zinc and magnesium content of industrial
electrolyte is significantly higher than in filtrate solutions of the precipitation
experiments (which results in enhanced zinc and magnesium fluoride complex
formation). However, when the proposed magnesium removal process will be
integrated in the electrolytic zinc production process, the solution composition in the
magnesium fluoride precipitation section will be similar to industrial electrolyte rather
than the dilute filtrate of the experiments (which will become clear in Chapter 5 and
Chapter 6). Therefore, this experiment provides a valuable confirmation that the
return solution will have a (too) high fluoride content.
3.2.4 Conclusions drawn from the thermodynamic model
In this section, a thermodynamic model of magnesium fluoride precipitation in
concentrated zinc sulphate environment was evaluated. The model accounts for ionic
strength effects as it assumes that stability coefficients depend on ionic strength
according to Vasil’evs equation. Solution of the model revealed the following
features:
•
Optimal precipitation conditions in order to minimise residual fluoride are
between pH=4-4.5 near the boiling temperature of the solution.
•
At lower pH, residual fluoride increases as a result of increased HF formation.
•
An optimal magnesium removal efficiency of 72% could be calculated for the
purified electrolyte supplied (see Section 2.1).
52
•
The equilibrium pH of the precipitating system decreases with increasing
temperature due to enhanced zinc hydrolysis.
•
Some trace elements, if present, can dramatically increase the residual fluoride
content. Among them, aluminium, siliceous species at low pH and ferric iron.
•
The required residual fluoride content of less than 10 ppm cannot be reached in
the precipitating system. Therefore, an additional fluoride removal mechanism has
to be found.
•
The experimentally determined residual fluoride content is generally two to three
times higher than predicted by the thermodynamic model. However, when the
most important equilibria are arbitrarily manipulated by a factor two, the
experimental results fall within the calculated range.
3.3
Treatment of fluoride containing industrial zinc sulphate
solutions
Both from experiments as well as from thermodynamic modelling, it became clear
that the residual fluoride content of the zinc sulphate solution in a magnesium fluoride
precipitation process will be well above 10 ppm. It was concluded in Section 3.2 that
an additional fluoride removal step is imperative, which is the subject of this Section.
The solution can only be returned to the zinc plant, if one of the following two
conditions is fulfilled:
•
a natural fluoride bleed must be present somewhere in the zinc plant.
•
an additional fluoride removal step for the return solution has to be incorporated
in the magnesium removal process.
In any other case, return of the fluoride containing solution will result in accumulation
of the fluoride level in the zinc plant and eventually lead to problems in the cellhouse.
Therefore, first a short literature study was carried out in order to identify possible
fluoride removal methods.
3.3.1 Literature review
•
natural bleed
Literature sources support fluoride incorporation in jarosite. Dutrizac[3], for example,
mentions that F seems to replace the OH groups in jarosite, but he did not definitely
know whether the fluoride is substitutional or adsorbed. According to a study by
Getskin et al.[54], combination of infrared spectroscopy and thermal analysis proved
that fluoride ions enter the crystal lattice of jarosite partly replacing the hydroxyl ion.
The fluoride incorporation is directly proportional to the fluoride content of the
solution.
53
The use of goethite as a scavenger for fluoride in electrolytic zinc plants is well
known. Various references report fluoride coprecipitation or adsorption with
goethite.[12], [55]
Both iron removal processes seem to provide a fluoride bleed, although practical
application of such a treatment is not investigated. Furthermore, some considerations
should be taken into account. Firstly, some jarosite producing zinc plants subject the
jarosite product to a (hot) acid leach, which could strongly reduce the fluoride
removing capacity. Secondly, the destiny of jarosite or goethite should be considered.
In case that a zinc plant is not allowed to dump iron residues, often pyrometallurgical
treatment will be opted for. Then, the fluoride contained in jarosite or goethite could
be released as gaseous hydrofluoric acid.
•
additional fluoride removal
If a natural fluoride bleed is not present, an additional removal step is required. As a
high fluoride level results in corrosion of the aluminium cathodes in the cellhouse, it
seems advantageous to search for a treatment method in which reduction of the
fluoride level is accomplished by interaction with some aluminium compound.
Indeed, a number of patents could be identified, which claim the removal of fluoride
from zinc electrolyte by aluminium compounds.
According to Rakhmankulov et al.[56], removal of fluoride can be accomplished by
addition of Al2(SO4)3 to zinc electrolyte in the presence of dissolved SiO2.
Experiments are described in which the reaction temperature was 333K and the
addition of aluminium sulphate was followed by neutralisation with zinc oxide to
pH=5.1. Thus, aluminium hydroxide was formed and the removal of fluoride was
explained by a chemisorption mechanism. The fluoride level could be reduced to 10.6
ppm. At such a fluoride level, solutions can be returned to a zinc plant without
causing problems in the cellhouse. The fluoride removal efficiency reportedly
decreases from 8.5 ppm per g/l Al(OH)3 at high fluoride levels (500-700 ppm) to 1.6
ppm per g/l Al(OH)3 at low fluoride levels (10-50 ppm). The form of aluminium
hydroxide determines the effectiveness of fluoride removal (amorphous aluminium
hydroxide being most effective and crystalline gibbsite being hardly effective).
Another promising method, suggested by Hampton et al.[57], is the removal of fluoride
from zinc sulphate solution by mixing with aluminium anodising waste sludge.[57]
Aluminium anodising waste sludge is obtained from neutralising aluminium
containing acid waste streams. After filtration, the sludge contains aluminium
hydroxide and typically 70% moisture. The sludge is reportedly also effective when
dried and crushed before use. The fluoride-loaded sludge can be treated with an alkali
metal hydroxide solution in order to regenerate fluoride, although it is not specified
how to deal with the resulting fluoride solution. Another possibility is the treatment of
the fluoride-loaded sludge in a lead smelter. Such an approach would fit in the
philosophy of closing material cycles, but one could wonder what the destination of
fluoride is in the lead smelter.
Detournay et al.[58] suggest the addition of aluminium ions (2.5-3.5 g/l) and phosphate
ions (8.75-12.25 g/l) in stoichiometric quantities. Subsequently, the solution must be
54
neutralised with CaO, Ca(OH)2 or CaCO3 to pH=4.3-4.7, thereby producing a
fluorinated precipitate. At higher pH, unwanted zinc hydroxide precipitation might
occur, while at lower pH, fluoride adsorption is insufficient. The reaction temperature
should be between 318K and 363K. At higher temperature, again, defluorination is
insufficient, while at lower temperature, the filterability of the precipitate is not good.
After filtration, however, the fluorinated precipitate has to be discarded. In order to
prevent the production of gypsum, it could be worthwhile to use zinc oxide (calcine).
Oohara[59] proposes a method to remove fluoride from zinc sulphate solutions by
addition of titanyl sulphate dehydrate (TiOSO4) instead of an aluminium compound.
The addition is followed by adjusting the pH with zinc oxide to 2.1-4.2 (but
preferably 2.7). Fluoride ions are adsorbed to the hydrolysis product and separated
from solution by filtration. The fluoride content could be reduced to 10 ppm (at
pH=2.7).
A totally different approach is extraction of fluoride (and chloride) from zinc sulphate
solution with quaternary ammonium sulphate in the presence of high molecular ethyl
alcohol or 3-n-butylphosphate, which is suggested by Frolov et al.[60] and Kuzmin et
al.[61]. Frolov et al.[60] were able to reduce the fluoride content (originally being 150
ppm) of industrial zinc sulphate solution to 10 ppm after shaking with the organic
phase in a 1:1 ratio) for 5 minutes. The zinc content was not reduced by this treatment
method. Re-extraction of fluoride and regeneration of the extractant was proposed by
treatment with aqueous solution of sodium (or ammonium) bisulphate. Although it
was not specified what the destination of the resulting aqueous fluoride solution
should be, it is proposed here to treat the resulting NaF/NaHSO4 in a membrane
reactor similar as described in Chapter 2. Kuzmin et al. argue that re-extraction with
sodium bisulphate requires multiple steps and yield dilute halogenous solutions. They
claim to improve the re-extraction by performing the first extraction in a mixture of
tetra ammonium sulphate and a phenol derivative (with alkyl substituents in orthoposition). The extraction is improved and, more important, re-extraction can be
accomplished in a single step by treatment with (1M) sodium hydroxide solution. The
result is a relatively concentrated halogenous solution, of which it is believed that it
can be further processed to yield a raw material product. The extractant can be
regenerated by treatment with dilute sulphuric acid.
3.3.2 Experimental verification
From the literature review, it will be clear that there is sufficient technology available
to remove fluoride from zinc sulphate solution. Amongst others, jarosite, goethite and
aluminium hydroxide were identified as capable agents for reducing the fluoride
content. In this Section, it was experimentally verified whether addition of these
substances to the fluoride containing solution resulted in a decrease of the residual
fluoride content. Furthermore, two other substances, gypsum and calcium chloride
were also investigated.
55
•
jarosite
residual fluoride (ppm)
Dried jarosite particles were added to a fluoride containing zinc sulphate solution (in
fact, a filtrate solution of the magnesium fluoride precipitation experiments described
in Chapter 2 was used), which resulted in a decrease of the fluoride content of the
solution. The efficiency of the process is shown in Figure 3.7.
200
180
160
140
120
100
80
60
40
20
0
0
50
100
150
200
250
300
jarosite (g/l)
Figure 3.7 Fluoride reduction by addition of jarosite particles.
The slope of the linear fitting line is -0.57, which could be interpreted as a removal
efficiency of 0.57 ppm F per g/l jarosite. It could be doubted whether a linear fitting
line is appropriate here (r2 = 0.979). A more accurate approach is the following: for
each addition of jarosite particles the fluoride reduction (mg/l) per g/l of added
jarosite is calculated. The calculated removal capacities of jarosite are displayed as a
function of the residual fluoride content (by attributing the said efficiency for each
addition to the fluoride level {[F-final] + [F-initital]}/2). The resulting instantaneous
removal efficiency of jarosite in Figure 3.8 first decreases, but starts to increase again
when the residual fluoride content is 100 ppm. Probably, the steep increase after the
initial decrease is the result of slow kinetics, which would imply that the fluoride
reduction is due to previous additions. Therefore, the overall efficiency (total fluoride
reduction/total jarosite addition) is shown as well. After an initial efficiency of 1.5
ppm F per g/l jarosite, the removal efficiency decreases and stabilises at
approximately 0.6 ppm per g/l jarosite. The experimental results, depicted in
Figure 3.8, indicate that the residual fluoride content could be reduced to less than 20
ppm, from which it can be inferred that it is possible to sufficiently reduce the fluoride
content in order to justify return of the solution to the zinc plant. The specific surface
area of the used jarosite particles was determined at 11 m2/g (BET analysis), which
implies a surface concentration of 3.2*10-5 mol F/m2 jarosite in the final stage (below
100 ppm, Figure 3.8) of fluoride removal. The incorporation of fluoride into the
jarosite structure could be limited to the surface area and from that perspective, it may
be expected that the efficiency of the incorporation process will be much higher when
the fluoride containing solution is added during jarosite precipitation.
56
1.4
1.2
1
0.8
0.6
0.4
0.2
0
200
150
100
50
fluoride removal efficiency of
jarosite (mg F/g jar)
1.6
0
fluoride content solution (mg/l)
overall efficiency
instantaneous efficiency
Figure 3.8 Fluoride removal efficiency as a function of residual fluoride content.
•
goethite
In a simple indicative experiment, 5.8 grams of goethite was added to 100 ml fluoride
containing (122 ppm) solution, which resulted in a decrease of the residual fluoride
content to 67 ppm. Although the removal efficiency in this experiment is 0.95 ppm
per g/l goethite, the surface coverage is rather low due to the very high specific
surface area. The specific surface area of the used goethite particles was determined at
193 m2/g (BET analysis), which implies a surface coverage of only 2.6*10-7 mol
F/m2. The high specific surface area is caused by a very porous structure, but the
surface is probably not fully accessible for solution and thus fluoride-goethite
interaction.
•
aluminium hydroxide
In a verifying experiment, aluminium hydroxide (spec. surf. area of 1 m2/gram) was
added to 100 ml fluoride containing solution. Addition of 3.7 gram Al(OH)3 resulted
in a reduction of the fluoride content from 143 ppm to 97 ppm. The removal
efficiency (1.24 ppm per g/l Al(OH)3 is significantly lower than reported by
Rakhmanulov, which confirms that fluoride removal is more effective when
aluminium compounds are freshly precipitated.
•
gypsum
Addition of a few grams of gypsum to 150 ml fluoride containing solution
(266 ppm F) did not result in a reduction of the residual fluoride content.
•
CaCl2
Addition of 20 ml 0.4M CaCl2 solution to 150 ml fluoride containing solution
did not result in a reduction of the residual fluoride content.
57
From the literature review as well as from the experimental verification, it will be
clear that there is sufficient technology available to remove fluoride from zinc
sulphate solution. The choice for one of the technologies will depend on reagent costs
and the destination of the fluorinated product.
3.4
Characterisation of hydroxyl-bearing magnesium fluoride
containing physically bound water
(Section 3.4 is based on a publication in Powder Diffraction [62])
3.4.1 Introduction
Routine analyses of five magnesium fluoride precipitates revealed the following
peculiarities:
•
X-ray fluorescence (XRF) measurements indicated that the ratio Mg:F was not 1:2
as stoichiometrically would be expected. The average ratio Mg:F was 1:1.88
(standard deviation 0.026). Furthermore, a large loss of weight (10-15%) was
observed when the samples were heated to 500˚C.
•
X-ray diffraction (XRD) patterns of the precipitates did not match accurately with
that of sellaite (MgF2), but resembled it closely (Figure 3.9).
XRF and XRD analysis of reference MgF2 Patinal® (E. Merck Gmbh, Darmstadt)
yielded results that are to be expected for MgF2:
•
the observed Mg:F ratio in five tablets is consistent with the formula MgF2 (on
average 1:1.98, standard deviation 0.024)
•
XRD-pattern is consistent with sellaite structure (Figure 3.10)
•
no great weight loss when heated to 500˚C (0.2-0.7%)
58
Figure 3.9 Diffractogram of magnesium fluoride precipitate with silicon standard (N.B.S. 640); the
vertical lines represent the theoretical positions of the peaks for magnesium fluoride (ICDD
41-1443).
Figure 3.10 Diffractogram of reference magnesium fluoride (Merck) with silicon standard (N.B.S.
640); the vertical lines represent the theoretical positions of the peaks for magnesium
fluoride (ICDD 41-1443).
Note that fluorite is not found in Figure 3.9 which depicts a typical diffractogram of
the magnesium fluoride precipitate. Based upon the results of Crane and Ehlers
(1969)[63], who found evidence for limited amounts of substitution of hydroxyl ions in
MgF2 at 1 kbar and T>400ºC, it was hypothesised that fluoride was partly substituted
59
by hydroxyl ions. In view of the proposed magnesium removal process, examination
of this hypothesis is useful. Hydroxyl incorporation in magnesium fluoride would
affect the amount of fluoride solution to be used in the removal process thus affecting
the operating conditions. The available analytical techniques in order to verify this
hypothesis were inventoried:
•
XRF analysis
to compare the composition of experimental material with reference MgF2
(Merck)
•
Infrared spectroscopy
to detect the presence of OH-stretching and(if present) bending vibrations
•
Thermal analysis
to measure weight loss versus temperature and monitor enthalpy changes
•
Electronmicroprobe
to detect the possible presence of oxygen and determine particle size
•
XRD
to determine the unit cell and compare the crystallography with that of sellaite
(MgF2)
3.4.2 Experimental
Selective precipitation of magnesium fluoride is accomplished by the slow mixing of
purified zinc sulphate solutions from zinc plants with zinc fluoride solutions, adapting
the procedure described in Chapter 2. The composition of the magnesium containing
purified zinc sulphate solution is summarised in Table 3.6.
Table 3.6 Zinc electrolyte composition (determined with XRF).
element
Zn
Mg
Mn
SO4
g/l
144
13.7
5.1
274.5
A zinc fluoride solution was prepared by dissolving a stoichiometric amount of zinc
oxide in a 0.24M hydrogen fluoride solution. The zinc electrolyte is slowly mixed
with zinc fluoride solution (0.12M) in a ratio 1:2.75. The experiment was carried out
in a baffled. The suspension was filtered over a ‘whitebandfilter’ (pore size diameter
4-7 μm). The precipitate was washed with very dilute sulphuric acid (pH=4), filtered
again and dried overnight at 378K together with reference magnesium fluoride
(Merck). Subsequently, the above listed analysing methods were carried out:
60
•
XRF
Tablets were pressed using approximately 2 grams of sample and boric acid,
which were subsequently analysed with a Philips X-ray spectrometer (type
PW2400) using Uniquant software to correct for matrix effects
•
Infrared spectroscopy
A paste was made of the powders by mixing it with polychorotrifluorethylene-oil
(Kel-F). The paste was pressed between two windows. The IR-spectrum was
obtained using a PerkinElmer Spectrum 1000 FT-IR Spectrometer
•
Thermal analysis
Thermal analyses (thermogravimetry and differential thermal analysis) were
carried out with approximately 100 mg of sample (type NETZSCH STA 409C)
•
Electronmicroprobe
Polished sections were made of dried precipitate and analysed with a JXA-8800M
Electron Probe Micro-analyser.
•
XRD
X-ray diffraction patterns were obtained with a Philips x-ray diffractometer (type
PW3710) using Cu Kα radiation and Ni filter. For unit-cell determinations,
silicon (NBS 640) was added as an internal standard.
3.4.3 Results and Discussion
The composition of the precipitate and reference magnesium fluoride as determined
with XRF analysis is given in Table 3.7. It became clear that the ratio Mg:F in the
precipitate was inconsistent with the formula MgF2. A large amount (14%) of
elements lighter than fluorine (presumably O and H) was found.
Table 3.7 Composition of precipitate and reference (other elements were all
below 0.1 wt.% and are not given).
element
Mg
F
Zn
SO4
Si
Ca
rest
precipitate (wt.%)
33.44
49.45
1.22
1.50
0.20
0.19
13.93
reference (wt.%)
38.98
60.86
0.01
0.08
0.01
-
Some zinc and sulphate was found in the precipitates as well. However, these
amounts of zinc and sulphate probably resemble evaporate traces of the original
solution, which was not completely removed from the filter cake by washing with
very dilute sulphuric acid, rather than that zinc and/or sulphate ions are incorporated
in the crystal lattice. This idea is supported by the strong correlation of the measured
61
zinc and sulphate amounts in various magnesium fluoride precipitates, reflecting
different washing intensities.
450
y = 0.98x + 78.7
R2 = 0.95
SO4 (mmol/kg)
400
350
300
250
200
150
100
50
0
0
100
200
300
400
Zn + Mg + Mn (mmol/kg)
Figure 3.11 Relation between Zn/Mg/Mn and SO 4 in the precipitate.
In Figure 3.11, the molar contents of zinc, magnesium and manganese in the
precipitate are plotted against the molar sulphate content. As magnesium is largely
present as magnesium fluoride, it has been assumed that the fraction of magnesium
present in the precipitate due to insufficient washing is directly proportional to the
zinc content (given by the electrolyte composition of Table 3.6). A linear relationship
through the origin with a slope of 1 would be expected. However, the linear fitting
line does not go through the origin and the applicability of a linear line could even be
questioned. On the other hand, the slope of the linear fitting line corresponds
remarkably well with the expected molar ratio of 1:1.
The molar Mg:F ratio of the precipitate that was used throughout the study described
in this section was 1:1.89, whereas the reference material showed the normal
stoichiometric ratio of 1:2,00. When our hypothesis is correct, infrared analysis at
least should reveal the presence of hydroxyl groups in the precipitate. The infrared
analysis of the magnesium fluoride precipitate is shown in Figure 3.12.
62
53.6
50
45
40
35
30
%T
OH - stretch
precipitate
OH - bend
25
20
15
10
5
0.0
4000
3000
2000
cm -1
1500
1000
600
Figure 3.12 Results of infrared analysis of magnesium fluoride precipitate. The absorption bands of
the IR spectrum of Kel-F without sample is less intensive as a result of the thinner layer.
Both spectra are displayed on the same vertical scale.
In the precipitate, both OH-stretch and OH-bend vibrations are present, which proves
that at least water is present. Besides, hydroxyl-incorporation still remains a
possibility. The reference material does not significantly contain these absorption
bands, indicating that no significant water or hydroxyl is present. The analysis is not
conclusive, however, with respect to the nature of the present water. The possible
solutions may be:
•
•
•
•
MgF2·xH2O
MgF2 with physically adsorbed water
Mg(OH,F)2·xH2O
Mg(OH,F)2 with physically adsorbed water
However, recalling the diffractogram of the magnesium fluoride precipitate, which
deviated only slightly from the sellaite pattern, the XRD results exclude the above
mentioned first and third option. Lattice water would result in a much larger unit cell
of the precipitate with respect to sellaite, causing a larger deviation in the XRDpattern.
In order to refine the characterisation, thermal analysis of both the precipitate as well
as the reference material has been carried out. The thermograms of both solids are
provided in Figure 3.13. The materials were heated at 10ºC/min in air, using
approximately 100 mg of sample. The reference material did not show any weight
loss.
63
MgF2 Merck
DTA / uV / mg
No weight loss
exo
0.2
[1] 78.4°C
100
0
TG %
-0.2
95
[1] 167.4°C
[1] -13.88%
(1)
-0.4
-0.6
85
Enthalpy effect
-0.8
[1] 346.7°C
[1] -1.12%
85
[1] 962.1°
-1.0
[1] -2.72%
-1.2
MgF2 (this study)
Two weight loss stages !!
200
400
600
Temperature °C
800
(2)
80
1000
Figure 3.13 Thermograms of MgF2 precipitate and reference material.
The precipitate experienced two weight loss stages. The first weight loss is very great,
but has no clear enthalpy effect. Therefore, this weight loss stage is assumed to relate
to the release of physically bound water. The second step does have a small enthalpy
effect and could be dehydroxilation. Then, the following reaction would be expected:
MgF1.89OH0.11 → 0.945MgF2 + 0.055MgO + 0.055H2O
This reaction would result in a weight loss of 1.6%, which corresponds reasonably
well with the observed weight loss of 2.7%. In view of this difference, it should be
noted that the formula MgF1.89OH0.11 is too accurate and should be considered
indicative.
In addition, the electronmicroprobe was used to identify the particle size of both the
precipitate (aggregates) and the reference material from the secondary electron images
(Figure 3.14 and Figure 3.15). The precipitate turns out to form aggregates with
individual particles as small as about 100 nm. It is not difficult to imagine the capture
of very small amounts of water within the tiny pores of these aggregates.
Furthermore, qualitative analyses of these aggregates revealed the presence of
oxygen, in accordance with the presence of water and hydroxyl groups. The
aggregates consisting of very small individual particles provide an indication that the
dominant mechanism was homogeneous nucleation, i.e. the start of precipitation was
not dependent on the presence of impurities. For example, calcium fluoride cores
could not be identified.
64
Figure 3.14 Secondary electron image of precipitate.
Figure 3.15 Secondary electron image of MgF 2 reference.
It appeared from the evaluation of routine diffractograms recorded for the magnesium
fluoride precipitates, that these crystallised magnesium fluorides showed XRD
patterns that deviated slightly but in a consistent way from the pattern of synthetic
sellaite (ICDD 41-1443). It was decided therefore to investigate the crystallography of
the sample and to compare this with the unit cell of synthetic MgF2 “Patinal”®. The
precipitate and the reference MgF2 were finely ground, and mixed with an appropriate
amount of Si (NBS 640) as an internal standard. Diffractograms were recorded with
the instrument settings listed in Table 3.8.
65
Table 3.8 Diffractometer type and instrumental settings.
Scan Parameters
Settings
Start Angle (º 2Theta)
End Angle (º 2Theta)
Step Size (º 2Theta)
Time per step (s)
Measurement Type
Scan Mode
Control Unit
Goniometer
5
120
0.020
4
absolute scan
continuous
PW3710
PW3020
(Theta/2 Theta)
40 kV
50 mA
0.154056
1.00
0.04 rad
1.00
Generator Tension (kV)
Generator Current (mA)
Wavelength Cu Kα (nm)
Divergence Slit angle (º)
Incident Beam Soller Slit
Diffracted Beam
anti-scatter slit (º)
Receiving Slit
Diffracted Beam Soller Slit
Diffracted beam filter
Detector
PHD lower level (%)
PHD upper Level (%)
0.2 mm
0.04 rad
Ni
PW3011
30.0
70.0
As the pattern of the precipitate was only slightly shifted from the sellaite pattern, it
was assumed that the crystal system would be similar. The space group P42/mnm,
given for sellaite was therefore taken for the precipitate. Indexing of reflections was
carried out on basis of ICDD card 41-1443. Refinement was carried out with the
program LCLSQ version 8.4[64]. The results of the unit cell calculations are given in
Table 3.9.
Table 3.9 Comparison of theoretical and experimental determined unit cell
parameters.
Literature
ICDD 41-1443
(Sellaite, syn)
Synthetic MgF2
Merck “Patinal” ®
105038
Reference material
0.4622 ± 0.0007
0.3050 ± 0.0003
0.0652 ± 0.0002
Mg(F,OH)2
‘sample”
a0 (nm)
c0 (nm)
V (nm3)
0.46200 ± 0.00003
0.30509 ± 0.00001
0.06612 (calculated)
dx (g/cm3)
dm (g/cm3)
3.177
3.18
3.13 *
3.150
F30=106(0.0092,31)
F25=26(0.030,32)
F13=6(0.064,33)
*: on basis of analytical data indicating the formula Mg(F1.89OH0.11)
SS/FOM
0.467 ± 0.002
0.303 ± 0.001
0.066 ± 0.001
66
The unit cell of the reference MgF2 is, as expected, almost identical to the literature
value (ICDD 41-1443). However, there is a notable difference in the unit cell of the
precipitate with respect to the reference material and the literature data. The a-axis is
slightly larger, and the c-axis is slightly smaller. The effective ionic radii of F- and
OH- do not differ much: variations are in the order of 0.003 nm[65]. The OH- -ion is a
little larger. Only very slight changes in unit cell parameters are to be expected in case
of replacement of F- by OH-. The unit cell data are therefore in agreement with a
partial replacement of F- by OH-, and exclude incorporation of lattice water, as
deviations will have to be much larger then.
Table 3.10 lists the indexed X-ray diffraction patterns of the precipitate and the values
of ICDD-41-1443 are given for comparison. Notable is, apart from the difference in
d-values for several reflections, that the intensities of several X-ray reflections are
also different, which is consistent with significant OH/F replacement, as OH and F
have different atomic scattering factors. Also it can be noted that the (112) reflection
at 0.1380 nm in the sample has a much higher intensity (26 %) than the (112)
reflection (2%) at 0.13821 nm in sellaite. The (301) reflection at 0.13475 nm in
sellaite in turn has a similar intensity as the (112) in the precipitate. As the positions
are very close to each other, one might think that the 0.1380 nm reflection should
perhaps better be indexed as (301) but refinement gave in such case a lesser fit then
when (112) was used. Therefore, (112) is taken to be the correct index for the 0.1380
nm reflection.
It appears that the (110) reflection (0.3304 in the precipitate versus 0.3267 nm in
sellaite) is one of the most affected by OH-incorporation. Possibly, the position of this
strong line can be used to estimate the OH-content. Crane and Ehlers[63] in their paper
on the system MgF2-MgO-H2O used the (220) reflection of MgF2 for estimation of
the amount of solid solution of Mg(OH)F in MgF2. However, they just stated they did
this, but did not give data from which the amount of displacement relative to the OHcontent can be read. Several syntheses of Mg(F,OH)2 in our laboratory, however,
yielded similar water contents and almost exactly the same position for the (110)
reflection : 0.330 nm. The X-ray lines of the precipitate were significantly broader
than those of synthetic MgF2, but this is attributed to the very small particle size of the
precipitate (around 100 nm, Figure 3.14) in comparison with grains of the reference
material, which were up to 150 micrometer size (Figure 3.15). The pattern is
fundamentally different from the pattern for MgFOH, mentioned by Crane and
Ehlers[63]. They were, however, unable to index this XRD-pattern and could not
determine the crystal system of MgFOH. (In an aside from this study, an effort to
refine these data was also made, but it appears that the Crane and Ehlers[63] paper lists
too few reflections of MgFOH in order to yield a reliable symmetry, indexing or even
unit cell.)
67
Table 3.10 X-ray diffraction patterns.
hkl
110
101
200
111
210
211
220
Mg(F,OH)2
sample
Observed
d-value
(nm)
0.3304
0.2538
0.2336
0.2230
0.2085
0.1716
0.1652*
Mg(F,OH)2
sample
Calculated
d-value
(nm)
0.3301
0.2543
0.2334
0.2233
0.2088
0.1720
0.1651
Mg(F,OH)2
sample
I/I0
Sellaite
41-1443
d-value
Sellaite
41-1443
Intensity
%
77
13
5
100
25
57
21
(nm)
0.3267
0.25474
0.23105
0.22309
0.20672
0.17112
0.16335
%
100
15
1
76
24
56
20
0.1516
0.1476
0.1378
0.1327
0.1227
0.1116
0.1058
0.09096
-
14
2
26
3
3
5
3
3
-
0.15259
0.14607
0.14403
0.13821
0.13745
0.13176
0.12814
0.12272
0.11814
0.11549
0.11205
0.11149
0.10889
0.10550
0.10519
0.10330
0.10255
0.09931
0.09811
0.09786
0.097
0.092
0.09208
0.09124
0.090
0.09031
11
3
2
2
22
3
1
3
<1
2
2
6
3
2
3
1
<1
1
<1
1
1
<1
2
3
1
2
(overlap with
Si-line(311)
reflection
002
310
221
112
301
311
320
212
321
400
410
222
330
312
411
420
331
103
322
421
113
430
402
213
510
412
0.1517
0.1470*
0.1380
0.1327
0.1227
0.1118
0.1057
0.09097
-
reflections marked with * were NOT used in the unit cell determination
68
3.4.4 Conclusions
Summing up the results, we find that IR-spectrometry indicates the presence of
molecular water but does not rule out the incorporation of OH- in magnesium fluoride.
Thermal analysis shows a large weight loss and a small weight loss step
(accompanied by an enthalpy effect), which can be explained as slow release of large
amount of adsorbed water and a dehydroxylation (the latter giving a reasonable match
with the chemical composition from XRF results). Microprobe analysis shows
presence of oxygen, consistent with presence of water and/or hydroxyl, and shows the
precipitate to be consisting of aggregates composed of approximately 100-135 nm
sized particles with equally small pores. These pores could accommodate the
adsorbed water, and regarding their size, would release water only slowly because of
strong capillary forces. The small grain size also explains the broader X-ray
reflections for the precipitate in comparison with the reference material. Unit-cell
determinations yield a unit-cell for the precipitate which is so close to that of sellaite,
that it exludes the presence of lattice water. However, the unit cell determinations still
do show a significant difference in lattice parameters of the precipitate with respect to
sellaite and the reference material. This could be explained by OH-incorporation.
Intensity shifts for some X-ray reflections in the pattern of the precipitate further
support the incorporation of an anion different from fluorine. Regarding the combined
results we consider the hypothesis that there is small replacement of fluorine by
hydroxyl in the precipitate confirmed.
Combination of different instrumental analysing methods confirmed the hypothesis
concerning hydroxyl incorporation in selectively precipitated magnesium fluoride
from zinc electrolyte solutions. Specifically, the following conclusions could be
drawn:
•
magnesium fluoride precipitation (pH≈ 4) experiences hydroxyl incorporation
when selectively precipitated from zinc electrolyte solution, indicating that the
stability of pure MgF2 in water should be questioned;
•
aggregates of very small particles, 100 – 135 nm, are formed;
•
aggregates contain physically bound water trapped in pores;
•
the crystal structure is sellaite type, P42/mnm (tetragonal);
•
the stoichiometric formula becomes MgF1.89OH0.11 for this sample and will be
similar for other precipitates depending on the precipitation conditions;
•
as a result of the different unit cell and composition, the density differs slightly
from ‘pure’ sellaite and becomes 3.13 g/cm3;
•
the amount of fluoride solution required in the magnesium removal process is less
than would be expected on basis of a 1:2 stoichiometry.
69
4
CONVERSION OF MAGNESIUM FLUORIDE
TO MAGNESIUM HYDROXIDE
The conversion of magnesium fluoride to magnesium hydroxide was studied as part
of the process to bleed magnesium from zinc sulphate electrolyte. Previous
experiments (Chapter 2) indicated a relatively high residual fluoride content in the
conversion product, which might limit the application possibilities of the magnesium
hydroxide. In this chapter, reduction of the residual fluoride content of the magnesium
hydroxide to <1wt.% is aimed at by optimising operating conditions such as leach
concentration, holding time, reaction temperature and wash procedures. In addition,
the thermodynamics of the conversion reaction are studied. Finally, the thermal
decomposition of magnesium hydroxide to magnesium oxide is studied. A shortened
version of this chapter has been accepted for publication in Minerals Engineering.
4.1
Introduction
One of the possibilities to deal with magnesium fluoride is to convert the precipitate
into magnesium hydroxide by contacting it with sodium hydroxide solution.
Magnesium hydroxide has a wide range of applications. It could be used to neutralise
acid waste streams and treat flue-gases (sulphur dioxide), thus providing costeffective solutions to environmental problems. Other applications are as fertiliser
additive, flame retardant or in the pulp and paper industry.
Introductory experiments (Chapter 2) showed that the conversion of magnesium
fluoride to magnesium hydroxide is possible. However, the residual fluoride content
was rather high, which could limit the economic application possibilities of the
converted product. As an example, in Table 4.1 the composition of one of the
converted products (of which the preparation conditions are given below) is compared
with product specifications of Nedmag (a magnesium hydroxide producer in the
Netherlands).
The preparation conditions of the converted product were the following: industrial
zinc electrolyte was slowly mixed with zinc fluoride solution (0.12M) in a ratio 1:2.5
over a period of 40 minutes in the vessel with dimensions as described in Chapter 2.
After an additional stirring (200 rpm) time of 75 minutes (dimensions of stirring
device as described in Chapter 2), 1 litre suspension (containing approximately 5
grams magnesium fluoride) was filtered and subsequently washed with 160 ml dilute
sulphuric acid (pH=4). After a second filtration step, the washed precipitate was
added to 0.8 litre 0.5M NaOH and stirred at 300 rpm. An abitrarily long holding time
of 114 hours was chosen; it was assumed that after such a long holding time
thermodynamic equilibrium was reached. After solid/liquid separation, the
magnesium hydroxide was washed with distilled water, dried at 378K and analysed
with XRF (Table 4.1) and X-ray diffraction. XRD confirmed the presence of brucite
(Mg(OH)2) and fluorite (CaF2). Especially the presence of fluorite indicates that the
lower limit of the residual fluoride in the product is determined by the presence of
calcium in the industrial zinc electrolyte. Some calcium in the purified electrolyte will
end up in the magnesium fluoride precipitate. As diffractograms of the magnesium
fluoride precipitates did not show any sign of the presence of fluorite, it is considered
likely that calcium substitutes for magnesium in the sellaite structure. The magnesium
fluoride precipitate converts to magnesium hydroxide and calcium fluoride when it is
contacted with sodium hydroxide solution. Calcium fluoride is found in the XRD
pattern of the converted product.
Table 4.1 Comparison of the compositions (wt.%) of an example of marketable
magnesium hydroxide and one of the converted products.[66],[67]
MgO
ZnO
F
Ca(OH)2/CaF2
SiO2
MnO2
Fe2O3
Al2O3
Na
K
Cl
SO4
LOI
Nedmag-H
product spec.
wt.%
66.3-67.5
<0.5
<0.5
0.66-0.80
0.08-0.15
0.08-0.15
0.35-0.40
0.04-0.10
not given
not given
0.35-0.50
0.15-0.30
30.85-30.5
< means <50 ppm
Converted precipitate
wt.%
60.4
1.1
6.4
1.1
<
0.08
<
<
0.7
0.2
<
<
30.03
Under these precipitation conditions, the converted product does not comply to the
product specifications of Nedmag and it could well be possible that the magnesium
hydroxide is not pure enough for other applications. Therefore, one of the objectives
of the testwork described in this chapter is to investigate the possibilities to improve
the conversion reaction (especially regarding the residual fluoride content) by varying
the following parameters:
•
•
•
•
•
reaction temperature
reaction time
stirring velocity
washing intensity
leach concentration
The effects of reaction temperature, reaction time and stirring velocity relate to the
influence of kinetics on the residual fluoride content in the converted product. The
washing intensity indicates the effort required for removing (soluble) sodium fluoride
from the converted product. Experiments in which the leach concentration was varied
indicate what thermodynamically is possible.
72
4.2
Experimental
To ensure practical relevance magnesium fluoride produced from zinc electrolyte and
zinc fluoride solution was used instead of pure, laboratory grade magnesium fluoride.
The experimental method was similar to the method described in Chapter 2 and
above. The industrial electrolyte was slowly mixed with zinc fluoride solution (0.12
M) in a ratio 1:2.4 to 1:3. The suspensions were filtered and subsequently washed
with approximately 250 ml dilute sulphuric acid (pH=4) and dried overnight at 378K.
The composition of eight magnesium fluoride precipitates, which were used as
precipitate feed in conversion experiments, is given in Table 4.2. From the product
amounts, it will be clear that, although the ratio between electrolyte and zinc fluoride
solution was more or less constant, the volumes used varied: respectively from 490 to
1340 ml for the electrolyte and from 1225 to 3440 ml for zinc fluoride solution.
Table 4.2 Composition of various magnesium fluoride feed precipitates
(determined by XRF).
Test
Product
amount
(g)
1
2
3
4
5
6
7
8
8.6
8.6
8.6
24.7
12.4
11.5
11.7
10.8
T (K)
stir.
time
(min)
Mg
(wt%)
Zn
(wt%)
F
(wt%)
Ca
(wt%)
SO4
(wt%)
293
45
33.7 0.5
45.9 0.6
1.0
353
45
32.4 0.8
48.1 0.1
1.0
353
45
32.9 0.7
48.1 0.1
0.9
363
45
33.1 0.8
49.0 0.1
1.2
323
45
31.1 0.7
49.1 0.2
1.0
324
100
32.9 0.6
50.0 0.2
1.0
324
100
31.7 2.7
50.1 0.2
0.8
324
100
33.4 0.5
49.3 0.3
1.0
(LOI = loss of ignition, rest = sum Na, Si, K, Mn, Y, Ba)
rest
(wt%)
LOI
(wt%)
0.5
0.2
0.5
0.2
0.1
0.1
0.3
0.1
17.8
17.4
16.8
15.6
17.8
15.2
14.2
15.4
Despite the slight variations in the precipitation conditions, the composition of the
precipitates is rather constant, which is also shown in Figure 4.1 and Figure 4.2. The
only remarkable exception is the high zinc content in one of the precipitates. The
cause of this high zinc content is unclear, but as the zinc content of washed
precipitates in all other experiments (not only those described in Table 4.2) was below
1 wt.%, it was not further investigated. As the composition of the magnesium fluoride
precipitates was rather constant, it was considered legitimate to assume a constant
feed input, i.e. in the discussion below conversion products are not compared with
individual XRF analyses of magnesium fluoride feed.
73
60
50
wt.%
40
Mg
30
F
LOI
20
10
0
1
2
3
4
5
6
7
8
Figure 4.1 Composition with respect to major elements in magnesium fluoride feed.
3
wt.%
2.5
2
Ca
1.5
S
Zn
rest
1
0.5
0
1
2
3
4
5
6
7
8
Figure 4.2 Composition with respect to minor elements in magnesium fluoride feed.
The molar fluorine to magnesium ratio is not equal to 2:1, but on average about 1.9:1
(Figure 4.3), which is in excellent agreement with the results of Section 3.4 in which
the characteristics of magnesium fluoride were compared with reference MgF2.
74
2.5
F/Mg (mol/mol)
2
1.5
1
0.5
0
1
2
3
4
5
6
7
8
Figure 4.3 Molar fluorine to magnesium ratio in different precipitates. The average of this dataset is
1.9 with a standard deviation of 0.09.
Mass balances were carried out for tests 6 to 8 with respect to fluorine. The results of
these mass balances are summarised in Table 4.3. Between 88 and 91% of the input
fluorine (zinc fluoride solution) is found in the output products (magnesium fluoride,
filtrate solution and wash liquor).
Table 4.3 Mass balances for fluorine in precipitation experiments.
Input
electrolyte
ZnF2
total input
Output
filtrate
wash liquor
Volume (l)
0.65
1.60
0.65
1.65
F content (g/l)
0.52
1.50
0.003
4.46
1.965
0.300
0.25
0.086
10.80
50.0
Volume (l)
2.125
0.290
2.165
0.285
precipitate
total output
recovery
(%)
11.66
Mass F (g)
0.003
4.38
0.002
7.136
7.138
0.32
0.090
0.531
0.025
0.541
0.025
0.629
0.027
49.3
5.750
6.306
88
5.842
6.408
89
5.324
5.980
91
F content (g/l)
Mass (g)
11.50
0.003
4.38
0.25
0.088
0.002
7.227
7.229
0.002
6.570
6.572
Mass F (g)
F content (%)
50.1
These magnesium fluoride precipitates were used in conversion experiments, which
were carried out as follows. The precipitates were ground in a porcelain mortar and
suspended with distilled water to 100 g/l. The suspensions were added to a
(preheated) sodium hydroxide solution and stirred for a predetermined time. After the
solid/liquid separation (filtration or decantation), the precipitates were washed with
100 ml 0.1M sodium hydroxide solution and dried at 378K. Dried precipitates were
analysed with X-ray fluorescence spectrometry and X-ray diffraction.
75
4.3
Results
Diffractograms confirm the presence of fluorite in magnesium hydroxide products. As
the solubility product of CaF2 is much lower than the solubility product of Ca(OH)2, it
seems legitimate to assume that all calcium in the converted product is present as
fluorite. Therefore, in the discussion below, the total amount of residual fluoride is
corrected for the measured calcium content. The remainder fluoride content can be
attributed to the magnesium compound. The fractional influence of this phenomenon
on the molar ratio Mg:F was accounted for in Figure 4.3.
As it was anticipated and soon confirmed that the leach concentration dominates all
other variables, it was decided to assess the other variables in a limited number of
experiments (by changing a specific variable while keeping the others constant). It
should be realised that investigating the variables separately implies the assumption of
variables being independent of each other.
4.3.1 Reaction temperature
The reaction temperature was already investigated in previous conversion
experiments (described in Chapter 2) which indicated that there was no significant
influence of reaction temperature on the final residual fluoride content. In Table 4.4,
the experimental conditions and results of an additional experiment are summarised.
Table 4.4 Influence of reaction temperature on the conversion reaction.
leach ratio
[OH-]/{Mg2+}1
5.2
5.2
1
reaction
temperature
(K)
305
345
stirring
(rpm)
wash
liquor
holding
time (min)
residual F
(wt.%)
200
200
NaOH
NaOH
60
60
5.9
5.5
{Mg2+}= Molar magnesium concentration present in solid form (calculated from the
initial amount of magnesium fluoride).
Increasing the reaction temperature with 40K only results in a slight decrease of the
residual fluoride content. It will be clear that increasing the reaction temperature is not
an effective measure for the reduction of the residual fluoride content. It could be
argued that the [OH-]/{Mg2+}-ratio is not entirely unambiguous as, theoretically, the
same value will be attributed to 20 g magnesium fluoride in a 2.0M NaOH solution
as, for example, to 1 g magnesium fluoride in a 0.1M NaOH solution. In order to limit
the number of experiments, it was decided to neglect this aspect as in practice the
experimental range was only limited to 2-4 g magnesium fluoride in 0.25 to 0.75M
NaOH.
76
4.3.2 Stirring velocity
The effect of the stirring velocity of the axial impeller used on the conversion reaction
was investigated at a higher leach concentration and at intermediate temperature.
Again, the impeller width was 1/3 of the diameter of the vessel. The experimental
conditions and results are summarised in Table 4.5.
Table 4.5 Influence of stirring velocity on the residual fluoride content of the
conversion product.
leach ratio
[OH-]/{Mg2+}
10.4
10.4
reaction
temperature
(K)
325
325
stirring
(rpm)
wash
liquor
holding time residual F
(min)
(wt.%)
400
800
NaOH
NaOH
60
60
3.6
3.4
Increasing the stirring velocity, although only slightly, contributes to a reduction of
the residual fluoride content.
4.3.3 Washing intensity
The influence of the wash method of the magnesium hydroxide products was also
investigated. Not only the difference between washing the product once or twice with
dilute sodium hydroxide solution and not washing the product was examined, but also
the difference between washing with dilute sodium hydroxide solution and washing
with distilled water was assessed. The experimental conditions and results are
summarised in Table 4.6. Magnesium fluoride of a single batch was used to ensure a
constant composition in the starting material.
Table 4.6 Influence of washing method on the residual fluoride content of the
conversion product.
leach ratio
[OH-]/{Mg2+}
5.2
5.2
5.2
5.2
5.2
reaction
temperature
(K)
323
323
323
323
323
stirring
(rpm)
wash liquor
holding time residual F
(min)
(wt.%)
200
200
200
200
200
water
water 2x
NaOH
NaOH 2x
180
180
180
180
180
6.1
5.8
5.8
5.5
5.3
Washing the conversion product with distilled water results in a slight reduction of the
measured residual fluoride content as a result of the deportment of the entrapped
sodium fluoride solution. It will be clear that washing the product a second time with
water is not effective at all. Washing with dilute sodium hydroxide solution is slightly
77
more effective than washing with water. The second washing step of sodium
hydroxide is hardly effective.
4.3.4 Holding time
Residual fluoride content (wt.%)
The conversion experiments described in Chapter 2 made clear that prolonging the
holding time from 4 to 8 hours did not influence the residual fluoride content
However, when the holding time is shorter, there could be a significant effect. In
Figure 4.4, the residual fluoride content of the conversion products of four
experiments is shown with the holding time varying between 1 and 3 hours (indicated
by data labels) at different reaction temperatures.
7
60
60
6
180
5
120
4
3
2
1
0
300
310
320
330
340
350
Reaction temperature (K)
Figure 4.4 Residual fluoride content as a function of reaction temperature and holding time. The
holding time (minutes) of each experiment is indicated with a data label. [OH]/{Mg}=5.2,
wash liquor = NaOH, stirring rate=200 rpm.
Increasing the holding time from 1 to 2 or 3 hours seems to result in a slightly lower
residual fluoride content. However, the difference between the residual fluoride
content of the 3 h. experiment and the value of the one hour experiment is very small.
The combined effect of a temperature increase of 18K and an increase of the holding
time of 2 hours is only a reduction in the residual fluoride content of 0.4 wt.%. In
Section 4.3.1 it was demonstrated that increasing the reaction temperature with 40K
only resulted in a decrease of the residual fluoride content with 0.4 wt.%. In Section
4.3.2, increasing the stirring velocity from 400 rpm to 800 rpm only resulted in a
reduction of the residual fluoride content of 0.2 wt.%. Therefore, although the results
seem to be marginally affected by kinetics, it can be concluded that improving the
kinetics is not effective in reducing the residual fluoride. Consequently, the details of
the kinetics of this reaction were not studied.
4.3.5 Leach concentration
The influence of the leach concentration on the residual fluoride content was
investigated in two experiments of which the results are summarised in Table 4.7.
78
Table 4.7 Influence of leach concentration on the residual fluoride content of the
conversion product.
reaction
temperature
(K)
335
335
323
323
leach ratio
[OH-]/{Mg2+}
5.2
10.4
7.8
15.6
stirring
(rpm)
wash
liquor
holding time residual F
(min)
(wt.%)
200
200
200
200
NaOH
NaOH
NaOH
NaOH
120
120
60
60
5.1
3.9
4.7
2.9
residual fluoride content (wt.%)
The leach concentration has a much larger effect on the residual fluoride content than
the other variables as can be seen in Table 4.7. The dominant influence of the leach
concentration becomes even more apparent when the results of numerous experiments
(in which, apart from the leach concentration also other parameters were varied) are
plotted in Figure 4.5.
12
10
8
6
4
2
0
0
10
20
30
40
50
2+
[OH-]/{Mg }
Figure 4.5 Residual fluoride content of conversion product as a function of the molar ratio
[OH-]/{Mg2+}.
From Figure 4.5, it becomes clear that in all these experiments, the [OH-]/{Mg2+}ratio
dominates all other parameters such as temperature, holding time, stirring velocity etc.
Each vertical intersection in Figure 4.5 encloses the combined effect of all parameters
except the [OH-]/{Mg2+}-ratio. A significant reduction of the residual fluoride content
only can be accomplished by using a very large excess of sodium hydroxide solution.
Optimising other parameters can only be seen as fine-tuning measures.
Performing the conversion reaction with a stoichiometric excess of sodium hydroxide
solution yields a mixture of sodium fluoride and sodium hydroxide solution, which
has to be treated in the membrane cell reactor, as described in Chapter 2. The detailed
study of the membrane cell reactor, which is presented in Chapter 5, will show that
even a small amount of hydroxide in the solution causes scaling and fouling of the
anion exchange membrane. In order to prevent these problems, addition of zinc
fluoride to the NaF/NaOH solution, before it is fed to the membrane reactor, could be
79
considered. Zinc hydroxide precipitates according to reaction (4.3.1) and after solid
liquid separation it can be used as neutralising agent in other parts of the zinc plant.
ZnF2 + 2NaOH → Zn(OH)2 + 2NaF
(4.3.1)
This reaction was experimentally investigated: 0.5 litre 0.12M ZnF2 and 0.5 litre
0.25M NaOH were mixed and stirred at 400 rpm during 60 minutes at room
temperature. After solid/liquid separation, washing the zinc hydroxide product with
sodium hydroxide solution and drying, XRF analysis did not show any sign of
fluorine anymore. Therefore, the zinc hydroxide can safely be used in other parts of
the zinc plant without introducing any fluoride. Another option is neutralisation of
the NaF/NaOH solution with hydrofluoric acid.
Neutralising excess sodium hydroxide with zinc fluoride prevents scaling of the anion
exchange membrane in the electrodialysis reactor. However, the amount of sodium
and fluoride ions to be transported across the membranes is increased, thus
contributing to higher energy consumption. Unfortunately, Figure 4.5 indicates that a
large excess of sodium hydroxide is required in order to obtain an appreciable
reduction of the fluoride content. As an example, the effect of a molar ratio
[OH-]/{Mg2+} of over 40 is shown in Figure 4.5. Even with this molar ratio, the
residual fluoride content is still about 2 wt.%, which is much higher than the product
specifications given in Table 4.1. Obviously, neutralising such an amount of sodium
hydroxide would lead to an unacceptable increase of the energy costs in the
membrane cell reactor. Therefore, the intermediate range of the [OH-]/{Mg2+}-ratio
between 10 and 40 is not investigated. Although, it is possible to convert magnesium
fluoride (sellaite) to magnesium hydroxide (brucite) by contacting it with sodium
hydroxide solution, it must be concluded that from a practical point of view it is not
possible to produce pure magnesium hydroxide. If it is not possible to sell this
fluoride containing magnesium hydroxide, an additional process step is imperative.
4.3.6 Phase characterisation
In Figure 4.6, a typical diffractogram of the conversion reaction product is depicted.
Figure 4.6 Typical X-ray diffractogram of a conversion reaction product.
80
Apart from brucite (Mg(OH)2), fluorite (CaF2) is found as well, but sellaite (MgF2) is
not found. The amounts of fluoride, calcium and sodium found in this particular
sample are given in Table 4.8.
Table 4.8 Composition of a conversion reaction product determined by XRF.
element
F
Ca
Na
wt.%
10.0
0.6
0.1
Based on the calcium content, it can be calculated that the amount of fluoride present
as CaF2 is 0.5 wt.%. Furthermore, one could argue that a small portion (0.1 wt.%) of
fluoride is present as entrapped sodium fluoride solution. The remainder fluoride
content (9.4 wt.%) can be attributed to the magnesium compound. However, it is not
present as magnesium fluoride anymore, as in that case it would have been detected as
sellaite with X-ray diffraction. Therefore, it can be deduced that brucite experiences
substitution of hydroxyl groups by fluoride, which implies that the conversion product
should be represented as MgOH2-yFy. Figure 4.5 can be adjusted to depict the
stoichiometric amount of fluorine in the conversion product (y in the chemical
formula) as a function of the [OH-]/{Mg2+}-ratio, which is done in Figure 4.7.
0.400
y in MgOH2-yFy
0.350
0.300
0.250
0.200
0.150
0.100
0.050
0.000
0
10
20
30
40
50
2+
[OH-]/{Mg }
Figure 4.7 Amount of fluoride in the formula of the conversion product as a function of the molar ratio
[OH-]/{Mg2+}.
4.3.7 Other impurities
The discussion regarding the purity of the magnesium hydroxide product focused on
the residual fluoride content for two reasons. Firstly, because the residual fluoride
content is by far the largest impurity and secondly, because most other impurities are
not affected by the conversion process. When impurities are present in the magnesium
fluoride feed, there is no mechanism in the conversion process to liberate them. As an
81
example, the impurities of the conversion products of the experiments investigating
the washing method (described in Table 4.6) are provided in Table 4.9.
Table 4.9 Impurities in magnesium hydroxide conversion products.
wash liquor
water
water 2x
NaOH
NaOH 2x
F (wt.%)
6.1
5.8
5.8
5.5
5.3
Zn (wt.%)
0.68
0.74
0.74
0.73
0.73
Ca (wt.%)
0.25
0.24
0.26
0.25
0.25
Mn (wt.%)
0.04
0.04
0.04
0.04
0.04
Na (wt.%)
2.25
0.22
0.07
0.39
0.37
The respective impurities will be briefly discussed below.
•
The zinc content of the conversion products is not affected by the washing
method, which could be expected. Any zinc in magnesium fluoride feed
immediately precipitates in the form of zinc hydroxide when the magnesium
fluoride product is contacted with sodium hydroxide solution. Therefore,
reduction of the zinc content in the magnesium hydroxide product can only be
realised by enhanced washing of the magnesium fluoride precipitation product.
•
As discussed, calcium is present as CaF2, which does not convert to Ca(OH)2
under these conversion concentrations. In order to reduce the calcium content (and
simultaneously also fluoride) in the precipitate, it is necessary to select feed
solution with a lower calcium content for the magnesium fluoride precipitation.
•
A very small amount of manganese is still present. However, this small amount
falls within the product specifications of Nedmag (Table 4.1) and may thus be
acceptable.
•
Sodium is present due to evaporation of entrapped NaF/NaOH solution. Washing
magnesium hydroxide precipitates with dilute sodium hydroxide solution
obviously will leave some sodium in the conversion product. When distilled water
is used, virtually all sodium can be removed.
•
The only impurity that has been removed completely by washing the magnesium
hydroxide product is sulphate.
4.4
Thermodynamics of the conversion reaction
As a first attempt to understand the equilibrium of this conversion, the reaction of
pure magnesium fluoride to pure magnesium hydroxide is considered:
MgF2(s) + 2OH-(aq) → Mg(OH)2(s) + 2F-(aq)
(4.4.1)
82
Based on the dissociation constant of water as well as on the solubility products of
magnesium fluoride and magnesium hydroxide, a stability diagram can be constructed
for a fixed fluoride concentration:
K w = [ H + ][OH − ]
↔ [ H + ]2 =
K w2
[OH − ]2
K sp , Mg ( OH ) 2 = [ Mg 2 + ][OH − ]2 ↔ [OH − ]2 =
K sp , MgF2 = [ Mg 2 + ][ F − ]2 ↔ [ Mg 2 + ] =
(4.4.2)
K sp , Mg ( OH ) 2
(4.4.3)
[ Mg 2 + ]
K sp , MgF2
(4.4.4)
[ F − ]2
Substitution of eq. (4.4.3) and (4.4.4) into (4.4.2) yields an expression for [H+]:
K sp , MgF2 K w2
[ H + ]2 = − 2
(4.4.5)
[ F ] K sp , Mg ( OH ) 2
Eq. (4.4.5) can also be written as:
+ 2
− 2
[H ] [F ] =
K sp , MgF2 K w2
(4.4.6)
K sp , Mg ( OH ) 2
Dividing by 2, taking logarithms and rearranging yields:
K sp , MgF2 K w2
1
log[ F ] = pH + log(
)
2
K sp , Mg ( OH ) 2
−
(4.4.7)
Although substitution of the universally accepted value of the dissociation constant of
water (Kw=10-14) is straightforward, selecting values for the other two constants is
more difficult as literature sources report different values for the solubility products of
magnesium fluoride and magnesium hydroxide. A few examples are given in Table
4.10.
Table 4.10 Solubility products of magnesium fluoride and magnesium hydroxide.
Equilibrium
constant
Ksp,MgF2
Ksp,Mg(OH)2
Ref-1
Ref-2
Ref-3
5.16*10-11
5.61*10-12
6.4*10-9
1.2*10-11
8.5*10-9
1.38*10-11
Ref-1 = 78th Handbook of Chemistry & Physics[68]
Ref-2 = 61th Handbook of Chemistry & Physics[69]
Ref-3 = Söhnel &Garside[40]
83
Substitution of the values provided by the different literature references yield the
following equations:
Ref-1:
Ref-2:
Ref-3:
log[F-] = pH-13.5
log[F-] = pH-12.6
log[F-] = pH-12.6
or
or
or
[OH-] = 0.33[F-]
[OH-] = 0.04[F-]
[OH-] = 0.04[F-]
In order to make a choice between the references, a simple experiment was
performed: 3 grams of magnesium fluoride precipitate (which was chosen in order to
deal with the specific particle morphology) was contacted with 1 litre 0.2M
NaOH/0.8M NaF solution (so that [OH-] = 0.20[F-]). References 2 and 3 predict that
the magnesium fluoride will convert (to some extent) to magnesium hydroxide until
the ratio has changed from [OH-] = 0.20[F-] to [OH-] = 0.04[F-]. This would result in
an increase of the fluoride concentration with about 0.1M. The first reference
however does not predict a reaction as there is no solid hydroxide present to dissolve
and thus shift the [OH-]/[F-]-ratio the opposite way. Therefore, according to the first
reference nothing will happen. Analysis of the solution with the specific fluoride ion
electrode after 60 minutes stirring contact indicated that the fluoride concentration
was not changed. Therefore, it is considered appropriate to use the first reference. It
could be noted that the solubility product of magnesium fluoride used here is different
from the one used in Chapter 3. However, the experiment described above relates to
the ratio of the values of the two solubility products and does not give information on
the values of the individual solubility products. Taking the logarithm of eq. (4.4.5),
followed by dividing by 2 yields eq. (4.4.8):
K sp , MgF2 K w2
1
pH = − log( − 2
)
2
[ F ] K sp , Mg ( OH ) 2
(4.4.8)
Substitution of eq. (4.4.2) in eq. (4.4.3) followed by taking logarithms of eqs. (4.4.3)
and (4.4.4) yields the following equations:
log[ Mg 2 + ] = −2 pH + log(
log[ Mg 2 + ] = log(
K sp , MgF2
[ F − ]2
)
K sp , Mg ( OH ) 2
K w2
)
(4.4.9)
(4.4.10)
If the first reference is used and it is assumed that [F-] = 0.1, eqs. (4.4.8) to (4.4.10)
reduce to:
pH = 12.5
log [Mg2+] = 16.7 – 2pH
log [Mg2+] = -8.3
(4.4.8a)
(4.4.9b)
(4.4.10c)
Three equations of straight lines are obtained, which can be plotted in a log(Mg)-pH
diagram (Figure 4.8). Thus, stability areas of magnesium fluoride, magnesium
hydroxide and the free magnesium ion are obtained.
84
-3
log(Mg)
-5
MgF2
-7
Mg(OH)2
-9
Mg2+
-11
-13
0
2
4
6
8
10
12
14
16
pH
Figure 4.8 Stability diagram of magnesium compounds at fixed [F -].
The stability diagram can be extended by including an additional log(F)-axis. This
way, the three-dimensional stability diagram, which is sketched in Figure 4.9 is
obtained.
log (Mg)
-6
MgF2
Mg(OH)2
-8
Mg2+
-10
log (F) 3
1
6
8
10
12
14
pH
Figure 4.9 Three-dimensional stability diagram of magnesium species.
The three-dimensional stability diagram suggests that there is a sharp conversion
plane where magnesium fluoride converts to magnesium hydroxide. However, the
experiments described in this chapter and Chapter 3 proved that a conversion of pure
magnesium fluoride to pure magnesium hydroxide does not occur. Instead, reaction
(4.4.11) is true:
MgF2-x(s)OHx + (2-x-y)OH-(aq) → MgOH2-yFy(aq) + (2-x-y)F-(aq)
(4.4.11)
85
The reaction equation contains two solid solutions with respectively isomorphic
substitution of hydroxide in sellaite and isomorphic substitution of fluoride in brucite.
Therefore, the stability diagrams in Figure 4.8 and Figure 4.9 in practice are not
completely valid.
From a thermodynamic point of view, it could be argued that the Gibbs energy of
formation of the (hydroxyl-containing) magnesium fluoride precipitate and the
(fluorine-containing) magnesium hydroxide conversion product depends on the
degree of substitution of respectively hydroxide and fluoride in the following way:
G MgF2− xOH x = G MgF2 − K 1 ⋅ x
and
(4.4.12)
G MgOH 2 − y Fy = G Mg ( OH ) 2 − K 2 ⋅ y
(4.4.13)
At equilibrium, the following relation holds:
∆G ≡ 0 ≡ G Mg ( OH ) 2 − K 2 ⋅ y + (2 − x − y )G F − + (2 − x − y ) RT ln( F − ) − G MgF2 + K 1 ⋅ x
− (2 − x − y )GOH − − (2 − x − y ) RT ln(OH − )
(4.4.14)
↔
( F ) /(OH ) = exp(−
G Mg ( OH ) 2 − K 2 ⋅ y + (2 − x − y )G F − − G MgF2 + K 1 ⋅ x − (2 − x − y )GOH −
(2 − x − y ) RT
The Gibbs energies are known (values are given in Table 4.11) and can be substituted
in eq. (4.4.14).
Table 4.11 Gibbs energies of formation[70],[76]
compound
MgF2(c)
Mg(OH)2(c)
FOHMg2+
∆G(298K)
kJ/mol
-1071.1
-833.5
-278.8
-157.3
-454.8
In conversion experiments, x and y can be determined with XRF measurements of the
magnesium fluoride feed and the magnesium hydroxide conversion products.
Furthermore, the equilibrium concentrations of [OH-] and [F-] can be determined and
used as approximation of (OH-) and (F-), leaving only K1 and K2 unknown. Two
conversion experiments yield two sets of data, respectively x1, y1, [OH-]eq,1, [F-]eq,1
and x2, y2, [OH-]eq,2, [F-]eq,2. Respective substitution of these data in eq. (4.4.14)
provides two equations with two unknowns (K1 and K2) which can be solved.
It was assumed that x=0.1 in the magnesium fluoride feed (derived from Figure 4.3)
and the equilibrium hydroxide concentration was derived by subtraction of the
apparent hydroxide consumption from the initial hydroxide concentration. It should
86
)
be noted that different sets of experiments resulted in different values for K1 and K2. It
was decided that the most reliable values were obtained at low leach concentrations as
deviation of activity coefficients from unity increases with concentration. The values
of K1 and K2 were determined at 48.7 and 152.2 respectively. Subsequently, the
measured [OH-]/[F-]-ratio in other experiments was compared with the theoretical
[OH-]/[F-] (i.e. calculated when K1, K2, x and y are substituted in eq. 4.4.14). This
comparison is summarised in Table 4.12 and also depicted in Figure 4.10. As might
be expected, K1 and K2 equate the experimental results well at low [OH-]/[F-], but at
higher leach concentrations deviation increases markedly.
Table 4.12 Comparison of theoretical and experimentally determined [OH-]/[F-].
x
0.10
0.10
0.10
0.10
0.10
0.10
0.10
0.10
0.10
[OH-]/[F-] th.
7.94
5.83
5.00
4.10
2.87
2.13
1.51
1.00
0.64
y
0.034
0.088
0.114
0.146
0.202
0.246
0.293
0.347
0.401
[OH-]/[F-] exp.
24.80
9.00
5.70
4.10
2.50
2.00
1.50
1.00
0.60
30
[OH-]/[F-]
25
20
theoretical
experimental
15
10
5
0
0
0.1
0.2
0.3
0.4
0.5
y in MgOH2-yFy
Figure 4.10 Comparison of [OH -]/[F-]theor and [OH-]/[F-]exp.
Since K1 and K2 are known, eqs. (4.4.12) and (4.4.13) can be plotted in Figure 4.11.
As it is not known to what extent respective substitution of hydroxide in the sellaite
structure and fluoride substitution in the brucite structure is possible, the middle part
of Figure 4.11 is dashed. Crane and Ehlers[63] report the existence of an intermediate
compound Mg(OH)F in the MgF2-MgO-H2O system at 1 kbar and elevated
temperature (600-800ºC). However, the solid solution of Mg(OH)F in the brucite
structure reportedly decreases when the temperature and pressure are lowered. This
information is consistent with our experiments at lower temperature and atmospheric
pressure as Mg(OH)F was never found. Consequently, for the temperature range of
87
the experiments, the Gibbs energy cannot continue to decrease from both sides to a
minimum, which would imply the occurrence of Mg(OH)F. From both sides, the
Gibbs energy will start to increase again at some degree of substitution. However, it is
not known at which degree this will happen. Therefore, a dashed line has been
sketched in the middle of Figure 4.11.
-750
-800
G (kJ/mol)
-850
-900
-950
-1000
-1050
-1100
-1150
0
0.5
Mg(OH)2
1
1.5
y in MgOH2-yFy
2
MgF2
Figure 4.11 Gibbs energy of the magnesium fluoride and magnesium hydroxide product.
With the determination of K1, it becomes possible to predict the Gibbs energy of
formation of magnesium fluoride containing some hydroxyl groups. In principle, this
offers the possibility of refining the thermodynamic model described in Section 3.2.
On average (Figure 4.3), the composition of the magnesium fluoride precipitate is
MgF1.9OH0.1 and the corresponding precipitation reaction is:
Mg2+ + 1.9F- + 0.1OH- → MgF1.9OH0.1
(4.4.15)
Under the assumptions made, the standard Gibbs energy change of this reaction is
defined as follows:
∆G 0
= G MgF1.9 OH0.1 − G Mg 2 + − 19
. G F − − 01
. GOH −
= (G MgF2 − 01
. K1 ) − G Mg 2 + − 19
. G F − − 01
. GOH −
(4.4.16)
With the determination of K1 and the Gibbs energies provided in Table 4.11, the
standard Gibbs energy of the precipitation reaction can be calculated:
∆G 0
= -1071.1-(0.1*48.7)+454.8-(1.9*-278.8)-0.1*-157.3)
= -75.7 kJ
(4.4.17)
Subsequently, the reaction constant of the precipitation reaction can be calculated:
K prec = e
− ∆G 0
RT
75700
= e 8.314*298 = 186
. * 1013
(4.4.18)
Finally, the solubility product is obtained from the precipitation constant:
88
K sp , MgF1.9OH0.1 =
1
K prec
=
1
= 5.3 * 10 −14
13
. * 10
186
(4.4.19)
total residual fluoride content
(ppm)
In Section 3.2, the solubility product of pure magnesium fluoride was used as one of
the thermodynamic equilibria. This equation of the solubility product of pure
magnesium fluoride can be replaced with eq. (4.4.19). Model calculations concerning
the total residual fluoride content as a function of pH (59% Mg-removal) were
performed once again. The results are depicted in Figure 4.12.
3
2.5
2
1.5
1
0.5
0
0
1
2
3
4
5
6
pH
Figure 4.12 Total residual fluoride content as a function of pH (59% Mg removal).
Although the trend is identical to the trend depicted in Figure 3.3, the total residual
fluoride content is very low and does not comply with experimental results. In order
to explain this deviation from experimental work at least two arguments can be put
forward:
1.
Although K1 was determined from experiments performed at low leach
concentration, [F-] and [OH-] are not equal to (F-) and (OH-). Thus an
unknown error is introduced in the calculation of K1.
2.
The solubility product of the MgFxOH2-x complex was calculated from Gibbs
energies. When the solubility product of pure magnesium fluoride is
calculated from Gibbs energies, a value of 5.13*10-11 is obtained, which is
considerably lower than the value (6.6*10-9) used in the thermodynamic
model. If this lower solubility product had been used in the thermodynamic
modelling described in Section 3.2, then the calculated total residual fluoride
content would have been much lower. Consequently, the differences between
Figure 3.3 and Figure 4.12 would also be smaller. The occurrence of different
solubility products of magnesium fluoride in the literature has been mentioned
before (Table 4.10). Experimental work described in Chapter 2 and Chapter 3
indicate that the higher solubility product yields results, which are more in line
with experimental data than the lower solubility product.
89
4.5
Thermal decomposition of brucite to periclase
In Section 4.3, it was concluded that very high leach concentrations are necessary to
accomplish appreciable reduction of the fluoride content in the converted precipitate.
As the membrane cell reactor cannot cope with these high leach concentrations
(energy costs would be unacceptable), a fluoride containing magnesium hydroxide
conversion product is produced, which may not be saleable. Since the objective was
to produce a saleable product, the thermal decomposition of the magnesium hydroxide
precipitate was investigated.
At different temperatures, it was tried to convert magnesium hydroxide products with
varying amounts of substitutional fluorine to magnesium oxide. In each experiment,
approximately 2 grams of magnesium hydroxide product was calcined for 45-90
minutes at the specified temperature. The XRF analyses of the calcined products are
summarised in Table 4.13.
Table 4.13 Residual fluoride content after thermal treatment of magnesium
hydroxide products.
T (K)
573
773
1173
1173
1173
1273
1273
1323
residual fluoride (wt.%)
19.7
16.1
17.9
13.4
13.7
0.8
0.4
0.8
time (min)
60
60
60
60
60
90
45
60
The differences in the measured residual fluoride content stem from the varying
amount of residual fluoride in the magnesium hydroxide feed. The relative amount of
residual fluoride even increases due to the release of water vapour. Unfortunately,
only at the relatively high temperature of 1273K, a conversion to magnesium oxide is
accomplished. The results at lower temperature suggest that the following reaction
occurs:
MgOH y F2 − y →
y
2− y
y
MgO +
MgF2 + H 2 O
2
2
2
(4.5.1)
XRF analyses results of thermal decomposition experiments performed at 900°C
show that nearly all fluoride in the calcined precipitate remains in the solid phase and
that all hydroxides are decomposed to oxides. Diffractograms (of which a typical
example is depicted in Figure 4.13) of these products show the presence of three
phases:
•
•
•
periclase
sellaite
fluorite
90
This means that upon heating, the compound decomposes to form Mg(OH)2 and
MgF2 and at the same time, the Mg(OH)2 decomposes to form MgO and H2O (which
escapes). MgF2 particles seem to remain stable at this temperature.
Figure 4.13 Diffractogram of calcined magnesium hydroxide at 900°C.
Experiments performed at or above 1000°C (without air blown through the furnace)
show a different result. XRF-analyses results from these calcined precipitates show
very low fluoride concentrations (F<wt.1%), indicating that reaction (4.5.2) is
occurring at or above 1000°C. A typical diffractogram of this high temperature
conversion is provided in Figure 4.14.
Figure 4.14 Diffractogram of calcined magnesium hydroxide at 1050°C.
91
Only two phases are present, periclase and a small amount of fluorite. Sellaite has
completely disappeared and it seems reasonable to assume that fluorine leaves the
system as HF(g) according to the following reaction:
MgF2(s) + H2O → MgO(s) + 2HF(g)
(4.5.2.)
The results may seem surprising when the thermodynamic equilibrium constant of
this reaction is calculated from the respective Gibbs energies:
Table 4.14 Reaction constants for MgF2/MgO conversion.
T (K)
1273
1373
1473
reaction constant
3.44*10-3
1.48*10-2
5.19*10-2
Based on the equilibrium constant, conversion is not expected. However, obviously
conversion did happen. A possible explanation could be that pH2O is locally (on a
molecular level) very high due to reaction (4.5.1) and evaporation of cake moisture,
thus inducing reaction (4.5.2).
In practical applications, HF(g) can either directly be contacted with the NaF/NaOH
mixture resulting from the conversion of magnesium fluoride to magnesium
hydroxide or it can be captured by a zinc hydroxide slurry in order to produce zinc
fluoride. Like HF(g), zinc fluoride can be used to neutralise the NaF/NaOH mixture
and it can also be used to precipitate magnesium fluoride.
The results of the experimental investigation of the thermal treatment of magnesium
hydroxide products indicate that it is possible to produce pure magnesium oxide from
fluoride containing magnesium hydroxide at temperature at or above 1000°C. Thus, a
saleable product is obtained. After recovering HF(g), two possibilities could be
identified to integrate the recovered gas with the magnesium removal process. The
recovery and subsequent integration, however, was not experimentally investigated
and thus requires some additional research.
If reaction 4.5.2 is true, the fluoride content could be further reduced by increasing
pH2O and by removing HF(g). It might even be possible to directly convert magnesium
fluoride to magnesium oxide. Although investigating another process route falls
beyond the scope of this thesis, the idea was considered interesting enough to do some
preliminary experiments. Three experiments were carried out in which gaseous water
was introduced to the furnace. The experimental set-up is schematically depicted in
Figure 4.15.
92
Off-gases
containing HF
N2
N2/H2O
Covered with heating coil
in order to prevent condensation
T = 1000 0C
Flow meter
Water
reservoir
Sample of MgF2
Oil bath
Temperature device
(80 0C)
N2 vessel
Figure 4.15 Experimental set-up for the direct conversion of magnesium fluoride to magnesium oxide.
The oil bath temperature was kept constant at 353K; the resulting water vapour was
carried by a nitrogen flow of 40 l/h. Three experiments with magnesium fluoride
precipitate were performed. Experimental conditions like reaction time, furnace
temperature and amount of starting material as well as the resulting residual fluoride
content of the calcined product (XRF analysis) are provided in Table 4.15.
Table 4.15 Experimental conditions and results of direct conversion of
magnesium fluoride to magnesium oxide.
Exp. no.
1
2
3
Starting
material
MgF2
MgF2
MgF2
time (min)
T (K)
mass (g)
120
100
100
1323
1323
1288
1.12
1.71
1.82
Residual F
content (wt.%)
0.3
10.0
8.7
Clearly, the fluoride content of the magnesium fluoride precipitate was significantly
reduced in experiments 1 to 3. In the first experiment, the residual fluoride content
was very low. It could be considered remarkable that a relatively small decrease of
reaction time and increase of the amount of precipitate results in a decreased
reduction of the fluoride content (to 10.0 wt.% in experiment 2). Experiment 3 at a
lower reaction temperature yields a residual fluoride content comparable to
experiment 2. However, in all three experiments, appreciable conversion of
magnesium fluoride to magnesium oxide was realised. Therefore, it can be
recommended to further study the details concerning the thermodynamics and kinetics
of this reaction, which will eventually determine the energy requirements and thus
operating costs.
93
Pyrohydrolysis of magnesium fluoride is possible and has been done at the
Keramchemie fluid bed pilot plant, at the kg/h level, jointly for the US Enrichment
Company (USEC) and CAMECO Coproration, as part of the AVLIS uranium
enrichment programme[71]. This information together with the experimental results
described above indicate that pyrohdrolysis of the magnesium fluoride precipitate
indeed could be an alternative for the processing of magnesium fluoride. Zinc,
sulphate and fluoride will be volatilised and can be absorbed in a slurry of zinc oxide
and added to the process. The thermodynamics could even be more favaourable when
silica is added to the pyrohydrolyser.[71],[72]
Slage[73] patented a process which concentrates uranium from contaminated
magnesium fluoride. For the production of uranium metal, magnesium metal is used.
Greensalt (UF4) reacts with magnesium metal yielding magnesium fluoride:
UF4 + 2Mg(g) = U(s) + MgF2
This magnesium fluoride is contaminated with a variety of uranium chemical forms
and oxidation states. This matrix must be considered as a low-level nuclear waste and
must be disposed of. In the patented process, the magnesium fluoride is treated with
high temperature steam, preferable of around 1000°C. The reaction becomes usually
complete in about five to about ten hours. The residue solids are magnesium and
uranium oxides with minor other non-strippable impurities. These solids can be
dissolved essentially completely in a variety of acids. Once dissolved, the uranium
may be extracted by any of a variety of solvent extraction systems. The raffinate may
be evaporated to produce magnesium salt of the acid anion. The extracted liquor may
also be contacted with the recovered hydrogen fluoride in order to regenerate
magnesium fluoride, which can be disposed of.
In two jointly published papers, Messier[74],[75] describes the kinetics of high
temperature hydrolysis of magnesium fluoride. In Part I, he found that the ratedetermining process was chemical reaction at the MgF2-MgO interface. Furthermore,
he observed that under the experimental conditions used the reaction rate did not
change with increasing product layer thickness, which was explained by the ratio of
specific volume of product to reactant. For magnesium oxide and magnesium fluoride
this ratio is 0.57, indicating that the product layer will be sufficiently porous to
provide negligible resistance to the passage of gases. In the first paper[74], Messier
studied the reaction at 745°C, 775°C, 805°C and 845°C and found that the reaction
rate between these temperatures increases from approximately 1015 to 1016 molecules
cm-2 s-1.
This reaction rate does not correspond to the observed residual fluoride content (Table
4.14) in the converted product. The specific surface area of the magnesium fluoride
precipitate was determined at 46 m2/g (BET analysis). With a reaction rate of 1016
molecules cm-2 s-1, the amount of magnesium fluoride (respectively 1.71 and 1.82 g)
should have been completely converted within 4 seconds, which did obviously not
happen.
In Part 2, Messier demonstrates that changing specimen geometry and the formation
of thick magnesium oxide product layers greatly influence the reaction rate. Different
rates were obtained from measurements on different specimen types. Therefore, it is
94
recommended to further study the pyrohydrolysis of the magnesium fluoride
precipitate in order to determine the conditions under which complete conversion can
be realised. Subsequently, it should be investigated whether creating these conditions
is legitimate from an economic point of view.
95
5
REGENERATION OF CHEMICALS
MEMBRANE CELL REACTOR
IN
A
The conversion of magnesium fluoride to magnesium hydroxide by contacting the
precipitate with a sodium hydroxide solution leaves a sodium fluoride solution as a
process residue. Both from an economic as well as environmental point of view, it is
desirable to prevent the discard of this solution. The search for a treatment method for
sodium fluoride solution resulted in the development of an electrodialysis process
route. In Chapter 2, preliminary experiments showed that it is possible to recover zinc
fluoride as well as sodium hydroxide solution from the sodium fluoride solution in a
reactor equipped with ion exchange membranes and an active zinc electrode. In this
chapter, a detailed study of this electrodialysis process is presented and discussed. In a
short literature[77],[78] overview, some fundamentals concerning the principles of
electrodialysis are introduced and discussed in order to make an evaluation of the
proposed process possible with regard to mass transport, membrane functioning and
general energy requirements. The chapter focuses on applicability of the
electrodialysis process and therefore the performance of the reactor over a long time
was studied. Besides, special attention was given to the evaluation of the anion and
cation exchange membranes. It should be realised that the results of this chapter refer
to the chosen membranes. In Section 5.3, the choice of membranes is briefly
elucidated.
5.1
Theory
Electrodialysis is a mass separation process in which charged groups attached to the
polymer backbone of membranes and an electrical potential difference are used to
separate ionic species from an aqueous solution and other uncharged components.[77]
The most important applications of electrodialysis are:[78]
•
•
•
•
•
in the desalination of brackish water;
as a preconcentration step in the production of table salt (Japan);
in chlorine alkaline electrolyses;
as a wastewater treatment method;
in food, drug and chemical process industry.
In other industries, often some hesitation is found concerning the application of ion
exchange membranes, as they sometimes tend to clog with filth from process water.
This clogging is called fouling, which may hamper the functioning of a membrane.
When a precipitate is formed at the membrane surface, the term scaling is often used.
In Figure 5.1, fouling in a membrane is depicted.
Figure 5.1 Schematic representation of fouling on a membrane.[78]
The principles of a conventional electrodialysis process are shown in Figure 5.2.
Feed solution
AEM
CEM
-
+
-
+
-
+
AEM
CEM
-
+
-
+
-
+
+
-
+
Conc entrate
Diluate
Figure 5.2 Schematic diagram of the principles of electrodialysis. [77]
Between the electrodes, a large number of compartments is created by an alternating
series of anion exchange membranes (AEMs) and cation exchange membranes
(CEMs). If for example a sodium chloride solution is pumped through these
compartments and when a potential difference is applied, sodium ions will move
towards the negative electrode, while chloride ions move towards the positive
electrode. On their way, sodium ions easily pass the cation exchange membrane, but
they are retained when they arrive at the anion exchange membrane. Simultaneously,
chloride ions easily pass the anion exchange, but they are retained when they arrive at
the cation exchange membrane. The net result is the development of an alternating
series of enriched and depleted compartments providing diluate and concentrate. In
practical applications, a reactor may consist of hundreds of such compartments.[77],[78]
Most ion exchange membranes are produced as homogeneous polymer films 50 to
200 μm thick. Typically, the membrane is reinforced by casting onto a net or
fabric.[78] In a cation exchange membrane, negatively charged groups (e.g. -SO3-,
-COO-, -HPO2-, -AsO32- or -SeO3-) are attached to the polymer matrix (which can be
for example polystyrene, polyethylene, polysulphone or perfluorinated polymers[77]).
98
Figure 5.3 Structure of a cation exchange membrane with fixed negatively charged groups and mobile
cations.[77]
The fixed anions are in equilibrium with mobile cations present in the interstices of
the polymer matrix; these mobile ions are referred to as counter-ions. Anions are
excluded from the membrane and they are referred to as co-ions. This type of
exclusion is called Donnan exclusion. Similarly, an anion exchange membrane
contains positively charged fixed groups (e.g. -NH3+, -RNH2+, -R2NH+, -R3N+, -R3P+
or -R2S+).[77] The structure of a cation exchange membrane is depicted in Figure
5.3.[77]
Obviously, the driving force for transport in electrodialysis is an electrical potential
gradient. Ion transport (e.g. through one of the ion exchange membranes) can be
approximated by equation 5.1.1 [79]:
Jn = CnmUnm Δф/Δx
(5.1.1)
Since all electrical charge transfer results from ion transport, the mass flux and
electric current are closely related[77]:
i = FΣnznJn
(5.1.2)
The fraction of the current carried by the nth ion is known as the transport number,
which can be represented by equation 5.1.3.
Tn = znJn/ΣnznJn
(5.1.3)
Transport numbers can be both set up when discussing transport in bulk solution as
well as when discussing transport through a membrane. In the latter case, they are
usually called transference numbers. Due to Donnan exclusion, the concentration of
co-ions in membranes is very low. Consequently, their transference numbers are close
to zero. If electrodialysis of a uni-univalent electrolyte such as sodium chloride
solution is considered, the transference number of the counter ion Na+ over the cation
exchange membrane will be close to unity.
The energy requirement in electrodialysis results from both applying the required
electrical potential difference (at a specific current density) as well as from pumping
solution through the compartments.[77] The latter typically accounts for 25 to 50% of
99
energy costs[78]. The electrical potential difference basically has to overcome three
phenomena:
•
•
•
electrode reactions
concentration gradient
ohmic resistance of solution and membranes
As usually hundreds of compartments are placed between anode and cathode, the
energy consumption due to electrode reactions often can be neglected. Also the
concentration potential is usually significantly lower than the potential drop necessary
to overcome the ohmic resistance of the system.[77]
In electrodialysis systems, the ohmic resistance in the compartments is reduced by
vigorously stirring the solutions. However, due to the selective permeation of
(counter)-ions through the ion exchange membrane, inevitably immediately adjacent
to the membrane surface, the solution becomes depleted with respect to these ions. At
the other side, the solution immediately adjacent to the membrane (where the ions
leave the membrane) becomes enriched. In Figure 5.4, this situation is depicted for a
cation exchange membrane.
Figure 5.4 Concentration gradient adjacent to a cation exchange membrane. [78]
This so-called concentration polarisation controls the performance of the
electrodialysis cell. When the applied voltage is increased, the ion flux through the
membrane increases and consequently the region adjacent to the membrane (diffuse
boundary layer, δ) becomes more depleted. There might even be a voltage at which
the ion concentration immediately adjacent to the membrane becomes zero. The
current in the electrodialysis system at this point is called limiting current density.
When the voltage is further increased, the extra current will be carried by other
processes, for example transport of anions through the cation exchange membrane or
even by protons and hydroxyl ions formed by the dissociation of water.[78] The
limiting current density can be experimentally determined by setting up a so-called
Cowen-Brown plot in which the reciprocal current is plotted against the measured cell
resistance. An example of such a plot is depicted in Figure 5.5.
100
Figure 5.5 Example of a Cowen-Brown plot to determine the limiting current density. [78]
5.2
Regeneration of zinc fluoride and sodium hydroxide in an
electrodialysis process
For the treatment of the sodium fluoride solution, a typical electrodialysis cell
arrangement containing multiple cells, such as described above, is not useful. The
objective is to recover zinc fluoride and sodium hydroxide solutions and not to
concentrate a sodium fluoride solution. Therefore, a complete separation of sodium
and fluoride ions is imperative. For this reason, a cell design with bipolar membranes
was considered, which are known for their ability to recover acid and base from the
corresponding salt in wastewater. In bipolar membranes (which are actually laminates
of cation and anion exchange membranes) water is dissociated, thus producing H+ and
OH- for the acid and base recovery. The principles of an electrodialysis cell design,
using bipolar membranes is depicted in Figure 5.6.
NaF
r
o la
Bip
AEM
CEM
r
o la
Bip
Na+
H+
+
OH-
HF
F-
H+
OH-
NaOH
Figure 5.6 The principles of electrodialysis using bipolar membranes.
If the sodium fluoride solution would be treated in such a reactor, HF and NaOH
would be produced. There are, however, a few disadvantages in this set-up:
•
HF cannot be used directly to precipitate magnesium fluoride. The hydrofluoric
acid has to be neutralised with zinc oxide.
101
•
HF is a hazardous (and toxic) chemical of which the use is preferably avoided in
plant operations.
•
The practical application of bipolar membranes has been limited due to their
apparent instability[78].
Instead, a system has been designed in which a zinc fluoride solution is produced
directly by means of an electrodialysis process using anion and cation exchange
membranes. An active zinc electrode replaces the inert anode and stainless steel is
used as cathode. Upon applying a potential difference, the following reactions are
induced:
2H2O + 2eZn
Zn + 2H2O
→
→
→
2OH- + H2
Zn2+ + 2eZn2+ + 2OH- + H2
(cathode)
(anode)
According to the electrochemical series[28], thermodynamically the required potential
difference is only 0.07 V. In Figure 5.7, the principles of this set-up are depicted.
NaF
AEM
Zn
H2
H2
Zn
+
H2
F-
Na+
H2O
Zn
Zn
Zn
OH-
Zn2+
Zn
ZnF2
NaOH
Figure 5.7 Recovery of zinc fluoride and sodium hydroxide by an electrodialysis process.
The obvious advantage of this approach is the direct recovery of both zinc fluoride as
well as sodium hydroxide solution. A disadvantage of this set-up is that it is not
possible to use multiple cells between anode and cathode. Every anolyte compartment
needs an active zinc anode.
In Chapter 2, it was shown that applying a potential difference in the proposed set-up
resulted in the desired ion transport. The objective of this chapter is to acquire further
information regarding the following features:
•
•
•
•
•
transference numbers of sodium and fluoride ions
membrane stability
energy consumption
effect of scaling and fouling
methods to reduce scaling and fouling
102
5.3
Cell design
A series of experiments was carried out with a reactor set-up similar to the one
described in Chapter 2. First, the fixed parts of the reactor are described:
•
Nafion-350® (DuPont) was used as a cation exchange membrane. Nafion has a
perfluorinated backbone and is known for its exceptional chemical stability and
complete inertness to sodium hydroxide solution.
•
Neosepta-AHA® (Tokuyama Soda) was used as anion exchange membrane. The
structure of this membrane is similar to Nafion membranes and thus provides
good chemical stability as well. Furthermore, Neosepta membranes minimise
migration of protons across the membrane[79].
•
Various Perspex reactor parts were used, which provided the possibility to
construct different set-ups. Figure 5.8 shows photographs of the reactor parts and
Table 5.1 provides information regarding the dimensions of the Perspex reactor
parts.
1
2
3
4
Figure 5.8 Photographs of the individual reactor parts.
Table 5.1 Dimensions of membrane cell reactor parts.
Outer length (mm)
Outer width (mm)
Depth (mm)
Inner area (cm2)
Electrode area (cm2)
No. 1
150
150
25
108
-
No. 2
150
150
10
81
-
No. 3
150
150
40
81
-
No. 4
150
150
40
69
62
(zinc anode)
(s.s. cathode)
The membrane reactor is made up of Perspex compartments, two ion-exchange
membranes and two electrodes (attached to a Perpex holder). The threads through
these elements are bolted together with nuts. Between the compartments silicon
rubber seals were placed to get a watertight construction.
103
5.4
Experimental
5.4.1 Experimental set up of the first series
The reactor parts in the first series of experiments were placed in the following order
(the numbers refer to the numbered reactor parts depicted in Figure 5.8 and AEM and
CEM are respectively Anion Exchange Membrane and Cation Exchange Membrane):
4-1-AEM-1-CEM-1-4
Each compartment was filled with solution. In Table 5.2, the initial composition in the
membrane cell reactor compartments of the first series of experiments is given as well
as the applied voltage.
Table 5.2 Initial composition of membrane cell reactor compartments–1st series.
Test
No.
1
2
3
Voltage
(V)
2.5
2.5
2.5
Initial composition
anolyte
0.01M H2SO4
0.01M HCl
0.01M H2SO4
Initial composition
ampholyte
0.20M NaF
0.20M NaF
0.20M NaF
Initial composition
catholyte
0.2M KCl
0.2M KCl
0.2M NaOH
The initial as well as the final zinc and sodium concentration of the compartments was
determined with AAS spectrometry and the initial and final fluoride concentration
was determined with the ion specific fluoride electrode. The use of this ion specific
electrode was justified in Chapter 2. This set of experiments is similar to the
experiment described in Chapter 2 and amongst others provides additional
confirmation of the results of that experiment.
5.4.2 Experimental set up of the second series
In the second series of experiments, the volume of solutions was increased by
coupling the compartments with reservoir tanks. A pumping device accounted for
sufficient mixing of the solutions of the (stirred) reservoir tank and membrane cell
compartment. As a result, the experimental operation time of the electrodialysis
process could be prolonged from 4 to 7 hours. In order to accomplish the circulation
of solutions, the reactor set up had to be changed:
4-3-2-AEM-1-2-CEM-1-2-4
This reactor set-up provides the possibility to install a stirring device in the anolyte
compartment as well as outlets for the compartment solutions to the reservoir tanks. In
Table 5.3, the initial composition in the membrane cell reactor compartments of the
second series of experiments is given as well as the applied voltage. In one of the
experiments, sodium hydroxide was introduced in the ampholyte as the results of
Chapter 4 indicate that the residue of the conversion process will be a mixture of
sodium fluoride and sodium hydroxide solution. Recirculation of solution from the
104
magnesium fluoride precipitation reactor to the anolyte compartment of the membrane
cell reactor implies the presence of a high zinc sulphate concentration and therefore,
experiments with a high zinc sulphate concentration in the anolyte compartment were
carried out. The initial and final compositions of the compartments were determined
with AAS and the ion specific fluoride electrode. Volume changes of the respective
compartments were recorded.
Table 5.3 Initial composition of membrane cell reactor compartments-2nd series.
Test
No.
4
5
Voltage
(V)
2.5
2.5
Initial composition
anolyte
0.01M H2SO4
0.01M H2SO4
6
2.5
0.01M H2SO4
7
2.5
8
2.5
0.01M H2SO4
1.00M ZnSO4
0.01M H2SO4
2.00M ZnSO4
0.50M MgSO4
Initial composition
ampholyte
0.20M NaF
0.15M NaF
0.05M Na2SO4
0.175M NaF
0.025M NaOH
0.20M NaF
Initial composition
catholyte
0.07M NaOH
0.07M NaOH
0.20M NaF
0.07M NaOH
0.07M NaOH
0.07M NaOH
5.4.3 Experimental set up of the third series
In the third set of experiments, a slightly different set up was used:
4-1-2-AEM-2-1-CEM-1-3-4
The compartment solutions were not connected with stirred reservoirs anymore.
Again, the initial composition in the membrane cell reactor compartments as well as
the applied voltage is given in Table 5.4.
Table 5.4 Initial composition of membrane cell reactor compartments-3 th series.
Test
No.
9
10
11
Voltage
(V)
2.5
2.5
2.5
12
2.5
13
2.5
Initial composition
anolyte
0.01M H2SO4
0.01M H2SO4
0.001M H2SO4
2.0M ZnSO4
0.5M MgSO4
0.55M ZnSO4
0.11M MgSO4
0.08M MgF2(s)
0.001M H2SO4
2.0M ZnSO4
0.5M MgSO4
Initial composition
ampholyte
0.2M NaF
0.2M NaF
0.4M NaF
0.025M NaOH
Initial composition
catholyte
0.07M NaOH
0.07M NaOH
0.07M NaOH
0.4M NaF
0.01M NaOH
0.4M NaF
0.01M NaOH
105
In the third experiment, the anolyte compartment was not stirred in order to
investigate scaling effects. The initial and final compositions of the compartments
were determined with AAS and the ion specific fluoride electrode. Again, volume
changes of the respective compartments were recorded.
5.4.4 Experimental set-up of the long-term experiment
These three series of experiments yielded satisfying results, which are discussed
below. However, when electrodialysis processes are discussed, the robustness of the
process (and especially the stability of membranes) over longer operating times is
often questioned. In order to test the long-term stability of the membranes, it was
decided to carry out a long-term experiment of 500 hours.
The membrane reactor was extended with three (10 litre) reservoir tanks. The flow
rates were respectively 150 ml/min for the anolyte compartment and 50 ml/min for
the ampholyte and catholyte compartment. In Figure 5.9, the experimental set-up is
depicted.
V
I
AEM CEM
Figure 5.9 Set-up of the long-term experiment.
Contrary to the previous experiments, a fixed current density was chosen instead of
fixed voltage. When the current density is fixed, the expected compositional changes
can be determined and compared with actual changes. A current density of 50 A/m2
was chosen to avoid depletion of the ampholyte experiment in the course of the
experiment. Table 5.5 provides the initial composition of the solutions in the
respective compartments.
106
Table 5.5 Initial solution composition in the long-term experiment.
anolyte
ampholyte
catholyte
Initial composition
0.55M ZnSO4
0.11M MgSO4
0.08M MgF2(s)
0.9M NaF
0.2M NaOH
Volume (l)
10
10
10
Prior to the experiment, the membranes were loaded in a 0.9M NaF solution. The
anolyte and ampholyte compartments were stirred. In the anolyte compartment, the
objective of stirring the magnesium fluoride containing solution was twofold:
•
to provide sufficient precipitation surface area for transported fluoride in order to
invoke precipitation on magnesium fluoride particles rather than on the
membrane;
•
to scrub the membrane surface with the magnesium fluoride particles.
In the ampholyte compartment, the main objective was to decrease the diffuse
boundary layer. Every six hours, small amounts of HF and H2SO4 were added to the
ampholyte and anolyte compartment respectively in order to counteract the increase
of pH.
Prior and after the experiment, samples of the membranes were taken. The
membranes were cut in rectangular pieces, washed with de-ionised water, dried
overnight at 324K and prepared with standard techniques for electronmicroprobe
analysis. The initial and final solutions were filtered over a membrane filter (pore size
<5 μm) and analysed with Atomic Absorption Spectrometry (AAS) and the ion
specific fluoride electrode. Precipitates on the membrane surface, if present, were also
analysed by XRF. When the amount of precipitate was too small for XRF analysis
(<0.5 gram), X-ray diffraction (XRD) was performed. Every 3-4 days, samples of
50 ml were taken and analysed to determine the concentration change within the
compartments. The final OH- concentration in the catholyte was determined by
titration with a 1M HCl solution. The voltage over the membrane reactor was
continuously monitored and stored every 10 minutes by a computer using
Labnotebook 4.01 software.
5.5
Results
In this section, the current density profiles of experiments 1 to 13 are presented (the
applied electrical potential of 2.5 V is given as a check) as well as the resulting
compositional changes in the anolyte, ampholyte and catholyte compartment.
Combination of the compositional changes with the energy supply will provide
valuable information regarding the performance of the membrane cell reactor and is
discussed in Section 5.6.
107
10.00
35.0
30.0
25.0
20.0
15.0
10.0
5.0
0.0
8.00
6.00
4.00
2.00
0
50
100
150
200
250
Voltage (V)
current density
(A/m2)
In Figure 5.10, the current densities and applied voltages in the first series
(experiments 1 to 3) are depicted.
0.00
300
time (min)
current density-1
voltage-1
current density-2
voltage-2
current density-3
voltage-3
Figure 5.10 Applied voltages and current densities of the first series of experiments.
The average current density was rather low, probably due to the low initial
concentration of dilute sulphuric acid in the anolyte compartment. Generally, the
current densities increased in the course of the batch experiments as the ionic
conductivity of the anolyte compartment increased as a result of fluoride transport and
zinc dissolution. In Figure 5.11, the compositional changes in the first series of
experiments are depicted.
5
4
3
2
1
0
-1
content change (g/l)
6
Na-cat
Na-amp
F-amp
F-ano
Zn-ano
-2
1 2
3 Na-cat
Naamp
Zn-ano
F-amp F-ano
Figure 5.11 Compositional changes of the anolyte, ampholyte and catholyte compartment in
experiments 1 to 3. “Na-cat” represents the change in sodium content in the catholyte
compartment, “Na-amp” represents the change in sodium content in the ampholyte
compartment, “F-amp” represents the change in fluoride content in the ampholyte
compartment, “F-ano” represents the change in fluoride content in the anolyte
compartment and “Zn-ano” represents the change in zinc content in the anolyte
compartment.
108
50.0
10
40.0
8
30.0
6
20.0
4
10.0
2
0.0
0
100
200
300
400
Voltage (V)
current density
(A/m2)
In Figure 5.12, the current densities and applied voltages in the second series of
experiments are depicted.
0
500
time (min)
cur.dens.-4
cur.dens.-5
cur.dens.-6
cur.dens.-7
cur.dens.-8
volt.-4
volt.-5
volt.-6
volt.-7
volt.-8
Figure 5.12 Applied voltages and current densities of the second series of experiments
(cur.dens.=current density, volt.= voltage).
The decreasing current density trend in experiments 4 and 6 can be attributed to
depletion of the ampholyte.
6
4
2
0
-2
Zn-ano
F-ano
F-amp
Na-amp
Na-cat
-4
45
6
content change (g/l)
The compositional changes of experiments 4 to 8 are depicted in Figure 5.13 and
Figure 5.14. In experiments 7 and 8, the change of the zinc concentration is negligible
compared to the initial zinc concentration of the anolyte compartment. Therefore, in
Figure 5.14, the change in zinc content of the anolyte compartment has been omitted.
Na-cat
Na-amp
F-amp
F-ano
Zn-ano
Figure 5.13 Compositional changes of the anolyte, ampholyte and catholyte compartment in
experiments 4 to 6.
109
7
8 Na-cat
content change (g/l)
2
1.5
1
0.5
0
-0.5
-1
-1.5
-2
Na-cat
Na-amp
F-amp
F-ano
F-ano
Na- F-amp
amp
Figure 5.14 Compositional changes of the anolyte, ampholyte and catholyte compartment in
experiments 7 and 8.
Again, the compositional changes are as expected and provide additional confirmation
of the results of Chapter 2.
40.0
10
30.0
8
6
20.0
4
10.0
2
0.0
0
500
0
100
200
300
400
Voltage (V)
current density
(A/m2)
In Figure 5.15, the current densities and applied voltages in the third series of
experiments are depicted.
time (min)
cur.dens.-9
cur.dens-10
cur.dens.-11
cur.dens.-12
cur.dens.-13
volt.-9
volt.-10
volt.-11
volt.-12
volt.-13
Figure 5.15 Applied voltages and current densities of the third series of experiments.
The strong increase of the current density in experiments 12 and 13 could be caused
by the low initial sodium hydroxide concentration in the catholyte compartment.
110
9
10
Nacat
Naamp
Famp
content change (g/l)
6
5
4
3
2
1
0
-1
-2
Na-cat
Na-amp
F-amp
F-ano
Zn-ano
F-ano Znano
Figure 5.16 Compositional changes of the anolyte, ampholyte and catholyte compartment in
experiments 9 and 10.
1.5
1
0.5
0
-0.5
-1
-1.5
13
content change (g/l)
2
Na-cat
Na-amp
F-amp
-2
12
11
Na-cat
F-amp Na-amp
Figure 5.17 Compositional changes of the anolyte, ampholyte and catholyte compartment in
experiments 11 to 13.
5.6
Discussion of the experimental results
5.6.1 Transference numbers
When evaluating this electrodialysis process, one of the most important questions to
be answered is: “To what extent does the supplied energy contribute to the desired
transport of ions?” As discussed in Section 5.1, the answer can be numerically found
in the transference numbers of fluoride and sodium from the ampholyte to
respectively the anolyte and catholyte compartment. Therefore, comparison of the
observed compositional changes with the current densities is imperative. Transference
numbers can be derived from experimental results according to eq. (5.6.1).
t=
F ∆[ F − ] *V
Q
(5.6.1)
111
In principle, the transport or transference number for fluoride can be calculated in two
ways:
•
•
based on the depletion of the ampholyte compartment
based on the increase of the fluoride concentration in the anolyte compartment
However, only transference numbers based on depletion of the ampholyte
compartment were calculated as in some experiments, as transference numbers based
on the enrichment of the anolyte compartment were not useful. In those experiments,
magnesium was present in the anolyte solution, which resulted in the immediate
precipitation of transported fluoride as magnesium fluoride. Consequently, only a
meaningless transference number can be calculated. Furthermore, loading of the anion
exchange membrane with fluoride ions might also yield incorrect transference
numbers. The error in the mass balance is defined as follows:
Error =
([ Amp]t * Vamp ,t + [ Ano]t *Vano ,t ) − [ Amp]0 * Vamp ,0
[ Amp]0 * Vamp ,0
× 100%
(5.6.2)
In Figure 5.18, transference numbers for fluoride are provided based on the depletion
of the ampholyte compartment and depicted as a function of the absolute error in the
mass balance. Transference numbers based on experiments with a fault in the mass
balance of more than 10% were rejected. In some experiments, however, the fault in
the mass balance was rather high due to precipitation of magnesium fluoride in the
anolyte compartment. In those cases, the transference numbers were not rejected.
Furthermore, transference numbers of experiments in which the ampholyte contained
other anions than fluoride ions were not used. These experiments are indicated by
“open” data points in Figure 5.18. This approach yields an average transference
number of 0.93 with a standard deviation of 0.17.
2.5
tamp,F
2
1.5
1
0.5
0
0
10
20
30
40
50
absolute error (%)
Figure 5.18 Fluoride transference numbers based on depletion of the ampholyte compartment as a
function of the absolute error in the mass balance. The open data points indicate
experiments in which fluoride had competition form other anions.
Precipitation did not occur in the catholyte and ampholyte compartment and therefore,
the transference number for sodium can be calculated both based on the depletion of
the ampholyte compartment as well as based on the increase of the sodium
112
concentration in the catholyte compartment. In Figure 5.19, transference numbers for
sodium are provided as a function of the absolute error in the mass balance. When the
fault is >10%, the transference number is rejected for the calculation of the average
transference number.
2.5
tNa
2
1.5
1
0.5
0
0
5
10
15
20
absolute error (%)
ampholyte depletion
anolyte enrichment
Figure 5.19 Sodium transference numbers based on depletion of the ampholyte compartment as well
as on enrichment of the catholyte compartment, given as a function of the absolute error
in the mass balance.
This approach yields an average transference number of 0.90 with a standard
deviation of 0.25 for the depletion of sodium from the ampholyte and a transference
number of 0.92 with a standard deviation of 0.11 for the enrichment of the catholyte.
The average transference numbers of fluoride and sodium correspond quite well and
with experimental values of 0.90 and 0.92 respectively, it is shown that the proposed
process makes rather efficient use of the supplied charge.
The two experiments in which the ampholyte was a mixture of NaF and NaOH
indicated that hydroxide is preferentially transported from the ampholyte to the
anolyte compartment. Table 5.6 shows the initial and final [F-]/[OH-] ratio in the
ampholyte compartment.
Table 5.6 Preferential transport of hydroxide over fluoride.
Initial
[F-] amp
Final
[F-] amp
Initial
[OH-] amp
Final
[OH-] amp
0.080
0.24
0.064
0.21
5.5*10-3
3.2*10-3
4.3*10-4
8.5*10-4
Initial ratio
[F-]/[OH-]
(ampholyte)
15
75
Final ratio
[F-]/[OH-]
(ampholyte)
149
247
The remarkable increase of this ratio makes clear that when a mixture of NaF and
NaOH is fed to the membrane cell reactor, a neutralisation step is imperative, as
hydroxide will cause problems in the anolyte compartment. Neutralisation with zinc
fluoride or hydrofluoric acid could be an option, as this way the hydroxide is replaced
by fluoride and thus a ‘pure’ sodium fluoride solution is obtained.
113
5.6.2 Water transport
The ampholyte compartment experienced water transport to the anolyte compartment.
In the second and third series (experiments 4 to 13), volume changes were recorded
(except in experiment 6 and 10) and in Figure 5.20, the depletion of the ampholyte
compartment is depicted as a function of the current density.
0
water transport (ml/s)
0
10
20
30
40
50
-0.0005
-0.001
-0.0015
-0.002
-0.0025
average current density (A/m2)
Figure 5.20 Water depletion of the ampholyte compartment a function of average current density.
The catholyte compartment also experienced (to a lesser extent) water transport to the
ampholyte compartment, possibly caused by the depletion of the ampholyte
compartment. The volumes of the respective solutions were recorded before and after
each experiment. Thus, volume changes were calculated, which were subsequently
corrected for evaporation. By simply measuring volume changes of pure water in the
respective compartments of the membrane cell reactor over a period of a few hours
(without applying an electrical potential difference), the evaporation (ml/h) was
determined and extrapolated to the experiments. Possible volume changes due to
changes in concentration in the compartments (due to the transport of sodium and
fluoride) were assumed to be negligible. The results are summarised in Figure 5.21.
40
ml water
20
0
-20
cat
-40
amp
ano
4 5
7 8
experiment no.
9 11
12 13
cat
amp
ano
-60
Figure 5.21 Volume changes of the respective compartments (anolyte = ano, ampholyte = amp and
catholyte = cat) in various experiments.
114
In membrane applications, water transport is usually associated with the following
phenomena[78]:
•
•
Cations permeating the membrane carry water molecules in their hydration shells.
Osmotic transport of water may occur as a result of a concentration difference.
It seems legitimate to exclude the first option as the water transport is from the
ampholyte to the anolyte compartment. Cations are excluded from the anion exchange
membrane, which separates the anolyte and ampholyte. Therefore, it is considered
likely that osmotic transport is responsible for the observed transport of water.
The depletion in the catholyte and ampholyte compartment is not always balanced by
the increase in the anolyte compartment. One of the reasons could be the formation of
hydrated salts in the anolyte compartment during some of the experiments.
Furthermore, evaporation in all three compartments disturbs the mass balance.
5.6.3 Active zinc anode
In this section, the increase of the zinc concentration in the anolyte compartment in
various experiments is compared with the supplied charge. Thus, a so-called apparent
anodic zinc consumption is defined as the ratio of the amount of charge entering the
solution due to zinc dissolution and the charge delivered by the power supply:
Consapp
F∆[ Zn 2 + ]ano V
=
2Q
(5.6.3)
consumption (%)
In Figure 5.22, the apparent anodic zinc efficiency (as defined above) is depicted as a
function of pH.
500
450
400
350
300
250
200
150
100
50
0
apparent
anode
consumption
corrected
anode
consumption
0
0,5
1
1,5
2
2,5
pH
Figure 5.22 Apparent anode consumption as a function of pH.
Below pH≈2, the apparent consumption increases linearly with decreasing pH. The
most likely causes are the following:
115
•
Dissolution of the zinc anode according to the following reaction:
Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)
•
(5.6.4)
As a result of this acid attack, small grains of the anode can be liberated.
The following observations prove that some zinc is consumed without contributing to
the transference of fluoride and sodium ions:
•
•
•
•
gas evolution at the anode
increasing pH of the anolyte
roughening of the anode surface
presence of zinc anode particles in solution (a diffractogram of these particles was
identical to the diffractogram of the zinc anode).
pH
In Figure 5.23, the monitored pH of the anolyte compartment is provided for tests 4 to
8 and test 10 and 11. From the increase of pH, the amount of zinc dissolved according
to reaction (5.6.4) can be calculated. Subsequently, a corrected apparent zinc anode
consumption is obtained, which is also shown in Figure 5.22. The results seem to
suggest that particles are not liberated from the zinc anode above pH=2.0.
5
4,5
4
3,5
3
2,5
2
1,5
1
0,5
0
4
5
6
7
8
10
11
0
100
200
300
400
500
time (min)
Figure 5.23 pH of the anolyte compartment as a function of time in various experiments.
Figure 5.22 and Figure 5.23 strongly emphasise that keeping the pH above 2 is
imperative in order to prevent unwanted zinc dissolution. It can be recommended to
keep the pH between 4-4.5. At higher pH, unwanted zinc hydroxide precipitation
might occur.
5.6.4 Scaling
Some experiments experienced scaling in the anolyte compartment on the anion
exchange membrane. Not only was magnesium fluoride found, but also (hydrated)
zinc sulphates and traces of zinc fluoride and magnesium sulphate. Therefore, it was
116
deduced that local supersaturation near the membrane surface, as a result of (zinc)
cations migrating in the direction of the cathode, caused the scaling. One of the
measures to prevent this scaling could be vigorously stirring the compartment or (in
case of magnesium fluoride) to present magnesium fluoride seed particles in the
anolyte solution. The scaling did not influence the selectivity of the membranes in the
course of the experiments. It is expected that scaling would reduce the average current
density. However, this is not observed in the experiments, probably as a result of the
short operating times in the experiments.
5.6.5 Maximum current density
Resistance (Ohm)
The maximum current density was determined in a series of identical experiments in
which the only variable was the ampholyte (NaF) concentration. The result of these
experiments is the Cowen-Brown-plot depicted in Figure 5.24. From the CowenBrown-plot it can be deduced that at an ampholyte concentration of 0.3-0.4M, a
current density of approximately 150 A/m2 can be reached in this experimental set-up.
At an ampholyte concentration of 0.5M NaF, the left upward part of the data series
could not be measured due to limitations of the available equipment.
20
18
16
14
12
10
8
6
4
2
0
0.05 M NaF
0.10 M NaF
0.20 M NaF
0.50 M NaF
0
5
10
15
20
25
Reciprocal current (1/A)
Figure 5.24 Cowen-Brown plot in which the maximum current density is depicted as a function of the
NaF concentration in the ampholyte compartment.
5.6.6 Long-term experiment
A preliminary long-term experiment revealed the following features:
•
•
•
Without remedial measures, the reactor voltage continuously increases due to
scaling of the anion exchange membrane in the anolyte compartment.
Scaling of the membrane surface can be avoided by (periodic) sulphuric acid
addition (thus keeping the pH of the anolyte compartment below 4.5).
Brushing the anion exchange membrane restores the reactor voltage to the
calculated potential.
117
It will be clear that when the scaling is not removed, energy costs significantly
increase and the process could even entirely stop. Therefore, the above mentioned
measures were implemented in the long-term experiment as described in Section
5.4.4. Accurate pH-control in the industrial system could be difficult. However, it
should be realised that the acceptable range (pH=2 to 4.5) is rather broad and it should
be possible to keep the pH in this range by controlled acid additions. Besides, as
described above and in Section 5.7 there are other methods to oppose scaling. The
voltage-time plot as well as the pH-time plot of the long-term experiment is depicted
in Figure 5.25.
10
10
8
6
6
4
4
2
2
0
0
600
0
100
200
300
400
500
pH
Volts
8
time (hours)
measured voltage
pH anolyte
pH ampholyte
Figure 5.25 Voltage-time plot of the long-term experiment.
In the first fifty hours of the experiment a sharp increase of the pH of the ampholyte
as well as the anolyte compartment was observed. It seemed reasonable to assume that
this was due to hydroxide ions originating from the catholyte compartment.
Therefore, the functioning of the cation exchange membrane was questioned. After
replacement of the cation exchange membrane, the pH of both the ampholyte as well
as the anolyte compartment was soon restored by means of periodic acid addition.
The first sharp increase of the voltage, after approximately 70 hours, seemed to be the
result of clogging of the membranes with fine particles. Visual inspection of the
experimental set-up revealed that the anolyte mixer transported very fine droplets
from the anolyte to the ampholyte and even catholyte compartment, thus causing
scaling and fouling. The topside of the reactor was covered with a thin layer of salts.
Furthermore, it was observed that the stainless steel cathode suffered from corrosion.
The membranes were scrubbed and the stainless steel cathode was replaced.
Consequently, the reactor voltage dropped to the initial value.
The second increase of the reactor voltage (approximately after 200 hours) was
caused by temporarily failure of the acid addition pumps. As a result, the pH of the
anolyte compartment increased, which again caused scaling of the anion exchange
membrane. After repairing the pumps, the membranes were not scrubbed in order to
investigate if acid addition could solely control the reactor voltage. The reactor
voltage indeed dropped to a lower value and subsequently a fluctuating pattern around
4 V was observed. The slight increase of the trend of the reactor voltage towards the
end of the experiment can be attributed to the depletion of the ampholyte
compartment, which lowered its conductivity.
118
The anolytic side of the anion exchange membrane experienced magnesium fluoride
and zinc fluoride scaling.
FZn2+, Mg2+
AMPHOLYTE
ANOLYTE
Figure 5.26 Schematic representation of the cleaning effect of the stirring device.
The middle part of the membrane did not experience scaling due to the turbulent
environment created by the mixer as schematically depicted in Figure 5.26. The low
turbulent regions near the corners were not kept clean. The precipitates could be
removed by washing with dilute (0.001M) sulphuric acid.
Contrary to the short-term experiments, the cation exchange membrane experienced
scaling as well. The composition of the very loose scale was a combination of
NaMgF3 and ZnO. The latter is clearly the result of the of fine droplets of (zinc and
magnesium containing) anolyte solution as described above. As the pH of the
ampholyte was above 5, zinc hydroxide precipitation is to be expected, which
converts to ZnO upon drying of the membrane. NaMgF3 could be formed as a result
of evaporation of the magnesium containing (due to the fine droplets) sodium fluoride
solution. The loose scale easily could be removed by addition of 0.1M H2SO4
solution. In practical application, the dispersal of droplets will not occur as the
compartments will be closed. Therefore, scaling of the cation exchange membrane is
not to be expected.
In order to get an idea of the susceptibility of the cation and anion exchange
membrane to fouling as well as of the stability of the selective layer, secondary
electron images of both membranes were recorded and visually inspected.
119
Figure 5.27 Secondary electron images of an unused (left) and 500 hours tested (right) cation
exchange membrane.
In Figure 5.27, cross sections of both an unused membrane as well as the used cation
exchange membrane, which functioned properly for 500 hours, are depicted.
Examination of the cation exchange membrane with electron microprobe, before and
after the long-term experiment, does not reveal any deterioration of the membrane. In
both images, the selective layer (clearly visible as a bright, thin layer) is intact.
Contrary to the anion exchange membrane, the cation exchange membrane was not
exposed to (scrubbing) particles. Therefore, it is no surprise that the cation exchange
membrane lacks any sign of fouling. Apparently, the observed scaling did neither
influence the selective layer nor the backbone of the membrane.
The stability of the selective layer as well as susceptibility to fouling of the anion
exchange membrane was investigated similarly. Secondary electron images of the
anion exchange membrane before and after the long-term experiment are depicted in
Figure 5.28.
Figure 5.28 Secondary electron images of the unused (left) and 500 hours tested (right) anion exchange
membrane.
The functioning of the anion exchange membrane remained stable in the course of the
experiment, which is in agreement with the intact selective layer in the right-hand
image. Despite this proper functioning, fouling of the anion exchange membrane can
be observed close to the selective layer. Figure 5.29 provides a closer look on the
fouling of the membrane.
120
Figure 5.29 Detailed view on the fouling of the anion exchange membrane.
The composition of the fouling was investigated by means of an element plot (XRF
analysis) of which the results are depicted in Figure 5.30. As zinc fluoride seemed to
be a likely fouling agent, zinc and fluoride were selected for the element plot.
Figure 5.30 Electron microprobe element plots of the fouling of the anion exchange membrane (the
fouled area is bordered by two bright lines).
Zinc is present in the fouling but fluoride is not. Therefore, zinc fluoride is not a
fouling agent, but some other zinc compound is. As the pH of the ampholyte was
rather high, possibly zinc hydroxide was formed. The anisotropic character of the
anion exchange membrane implies that the position of the selective layer (facing the
anolyte solution or the ampholyte solution) could play a role in the fouling
mechanism. It is not possible to identify the side of the selective layer when the anion
exchange membrane is placed in the reactor. Therefore, the following two fouling
mechanisms could have occurred.
•
The selective layer faced the anolyte solution. The selective layer was not
completely impermeable for zinc cations.
•
The selective layer faced the ampholyte solution, thus zinc cations were able to
penetrate the membrane’s supporting fabric. Due to a local increased hydroxide
concentration, zinc hydroxide precipitation could occur.
These two options are schematically depicted in Figure 5.31.
121
Zn anode
AEM
Zn anode
Zn2+
OH-
OR
AEM
Zn2+
OHF-
Fselective
layer
fouling
supporting
fabric
selective
supporting
fabric fouling layer
Figure 5.31 Schematic representation of two fouling mechanisms of the anion exchange membrane.
The fouling could be removed by rinsing the anion exchange membrane with a 2M
sulphuric acid solution. However, the selective layer of the rinsed membrane could
not be detected anymore, which might suggest that the strong acidic solution
destroyed the selective layer.
Transference numbers of this experiment could be calculated as described before, by
comparing the compositional changes with the supplied amount of charge. The results
are summarised in Table 5.7.
Table 5.7 Transference numbers in the long-term experiment.
Na+
F-
tamp
tano or tcat
1.08
1.03
0.91
0.96
fault in mass
balance (%)
-11.5
-6.6
When the formation of scale is kept in mind, the transference numbers are remarkably
close to unity. The zinc anode consumption was also evaluated by comparing the
weight loss of the zinc anode with the supplied amount of charge. The results are
summarised in Table 5.8.
Table 5.8 Zinc anode consumption in the long-term experiment.
weight loss zinc anode
total zinc in solution (AAS)
expected amount of zinc
based on supplied charge
added amount of H2SO4
zinc dissolved by H2SO4
grams
298
285
220
110.3
73.6
122
The results indicate that addition of dilute sulphuric acid causes an excess zinc
consumption of 33%. As explained before, excess zinc consumption is twofold:
•
•
direct dissolution of the zinc anode
liberation of small grains
As the difference between the weight loss of the zinc anode and the amount of zinc
found in solution is very small (4.4%), it can be deduced that the number of liberated
zinc grains as well as the amount of zinc fluoride scaling is also small. The main
cause for excess zinc consumption is therefore direct dissolution of the zinc anode.
80
70
60
50
40
30
20
10
0
0
100
200
300
400
500
9
8
7
6
5
4
3
2
1
0
600
[Mg2+] and [F -] (g/l)
[Zn2+] (g/l)
The concentration changes of the anolyte compartment, with respect to the zinc,
magnesium and fluorine content, are provided in Figure 5.32. Obviously, the fluoride
concentration starts to increase after approximately 200 hours when all magnesium is
consumed (precipitation of magnesium fluoride). As a result of the increase of both
the zinc as well as the fluoride concentration, the solubility product of zinc fluoride
could have been exceeded towards the end of the experiment.
time (hours)
Zn conc.
Mg conc.
F conc.
Figure 5.32 Compositional change of the anolyte with respect to the Zn-, Mg- and F-content.
The initial and final composition of the (scrubbing) solids in the anolyte compartment
were analysed as well by means of XRF measurement of which the results are
provided in Table 5.9.
Table 5.9 Initial and final composition of particles in the anolyte.
Mg
Zn
F
S
LOI
rest
initial (wt.%)
27.2
6.4
41.3
1.4
22.2
1.5
final (wt.%)
29.3
9.4
39.5
3.7
17.4
0.7
123
The composition remained more or less constant, except for the slight increase of the
zinc content due to zinc fluoride precipitation. Comparing the increase in the
hydroxide concentration of the catholyte compartment with the supplied amount of
charge yields a cathode efficiency of 99%.
5.7
Membrane performance in a larger scale experiment
Based on the results of the long-term experiment, a new experimental set-up was
developed with the following objectives:
•
study membrane performance in a larger scaled experimental set-up
•
create a turbulent environment by means of increased fluid flow instead of stirring
•
reduce energy consumption by reducing the reactor width (and thus ohmic
resistance by the solutions)
In the larger scale experiment, the principles of the reactor set-up largely remained the
same (similar to the set-up depicted in Figure 5.9). A new reactor was constructed
with the following (inner) dimensions for each compartment: width=10cm, length=50
cm and depth=0.5 cm. Especially the last number is of interest as the reduced depth
brings along a reduction of the voltage drop due to ohmic resistance. The area of
electrodes was 10x30 cm, which implies that the lower and upper part of each
compartment does not face an electrode. Combined with the multiple inlet (and outlet)
construction of the compartments, entrance and exit effects are largely reduced. For
stability reasons, the membranes were placed between holders, which results in slight
reduction of the active area to two times 10x14 cm. Figure 5.33 depicts the electrode
holder as well as the membrane holder of the electrodialysis reactor.
Each compartment was coupled to a reservoir tank of 20 l. As the flow rates of the
solutions were relatively high, 4.5 to 12 l/min, neither the compartments nor the
reservoir tanks were stirred. The experimental set-up can be considered as a
combination of three separate fluid circulation systems. Figure 5.34 depicts one of
three fluid recycling lines, illustrating the possibility to continue fluid circulation
when the electrodialysis reactor is dismantled by opening valve 2 and closing valve 3.
124
10 cm
10.5 cm
Electrode
area
30 cm
14 cm
membrane
area
1 cm
10 cm
10 cm
14 cm
membrane
area
10 cm
10.5 cm
Electrode holder
Membrane holder
Figure 5.33 Dimensions of the electrode and membrane holders of the electrodialysis reactor.
3
ED
reactor
compartment
2
Flow meter
(l/min)
1
pH
valve
measurement
storage tank
tap point
tap point
Figure 5.34 Schematic view of 1 of 3 fluid recycling lines.
Figure 5.34 refers to the anolyte recycling line, but the only difference with the
ampholyte and catholyte is the pH-measurement in the anolyte line, which is not
carried out in the other lines. The pH of the anolyte compartment was kept constant
125
by controlled dilute sulphuric acid addition. Contrary to the previously described
long-term experiment, the pH control of the ampholyte solution with HF was not
carried out. The voltage over the membrane reactor as well as the pH of the anolyte
were continuously monitored and stored every 10 seconds by a computer using
Labnotebook 4.01 software. The voltage-time plots are discussed below. Figure 5.35
provides a photographic impression of the experimental set-up.
Figure 5.35 Photograph impression of the up-scaled membrane cell reactor.
Although this thesis does not treat the details of fluid dynamics in the membrane cell
reactor, the Reynolds numbers in the respective compartments were determined using
equations as described in the standard text book “Electrochemical Process
Engineering” (Goodridge and Scott)[80]. The maximum fluid velocities that could be
obtained using the most powerful pumps available were 0.4 m/s for the anolyte
compartment and 0.15 m/s for the ampholyte and catholyte compartment.
Combining these fluid velocities with data concerning cell dimensions, fluid density
and viscosity yields the following Reynolds numbers:
•
•
•
anolyte compartment
Re = 2787
ampholyte compartment Re = 1096
catholyte compartment Re = 1096
Usually, Reynolds numbers up to 2000 indicate laminar flow, while Reynolds
numbers above 2000 indicate a turbulent regime.[80] Therefore, the flow through the
anolyte compartment is just turbulent and the flow in the ampholyte compartment is
laminar.
Initially, it was attemted to test the membrane reactor in another long-term
experiment. Therefore, a fixed current density of 100 A/m2 was chosen in
experiments. This current density was significantly higher than in previous
experiments, but not so high that the ampholyte solution would deplete too quickly.
Again, Nafion® and Neosepta® membranes were used.
In preliminary experiments with the larger scale reactor, scaling of the anion
exchange membrane and a continuous rise of the reactor potential was observed. It
126
can be concluded that the environment was not sufficiently turbulent to prevent
scaling. Furthermore, it was deduced that addition of large pulses of dilute sulphuric
acid result in sharp decreases of the required reactor voltage. Although the reactor
voltage could be reduced by these acid pulses, the original voltage was not reached
and the rise of reactor voltage was only delayed.
46.55
44.33
42.12
39.90
37.68
35.47
33.25
31.03
28.82
26.60
24.38
22.17
19.95
17.73
15.52
13.30
11.08
8.87
6.65
4.43
2.22
10
9
8
7
6
5
4
3
2
1
0
0.00
Voltage (volts)
The settings of the pH-control system were changed and the pH was kept at 3.5 by a
continuous series of very small dilute sulphuric acid pulses. Despite this pH-control,
the reactor voltage continuously increased, as can be seen in Figure 5.36. Apparently,
pH control does not prevent the formation of scaling, but dilute sulphuric acid
additions can detach the scaling layer if these additions are sufficiently large.
Time (hrs)
Figure 5.36 Voltage-time plot of the first experiment with the up-scaled reactor.
At the end of one of the experiments, an interesting phenomenon was observed. Upon
the closure of valve 3 and the opening of valve 2 of the anolyte line, the reactor
voltage returned to its original value. The following mechanism might explain the
observed reduction of the reactor voltage: closure of valve 2 and opening of valve 3
creates a pressure difference in the membrane cell, which results in the bulging of the
anion exchange membrane. The bulging causes the breakdown of the scaling layer.
This mechanism is schematically visualised in Figure 5.37.
AEM
CEM
+
-
Na+
NaF
Zn2+
Mg2+
H2
F-
MgF2 scaling
NaOH
127
AEM
CEM
+
-
Na+
NaF
Zn2+
H2
NaOH
Figure 5.37 Exaggerate representation of the bulging of the anion exchange membrane, causing
breakdown of the scaling layer.
6.40
6.00
5.60
5.20
4.80
4.40
4.00
3.60
3.20
2.80
2.40
2.00
1.60
1.20
0.80
0.40
10
9
8
7
6
5
4
3
2
1
0
0.00
Voltage (volts)
An experiment was carried out with the objective to use the observed phenomenon to
control the reactor voltage. Figure 5.38 depicts the voltage-time plot of this
experiment.
Time (hrs)
Figure 5.38 Voltage-time plot of the second experiment with the up-scaled reactor.
It was possible to keep the reactor voltage below 3.5 V during 7 hours by periodically
creating a pressure difference over the anion exchange membrane. As it was not
possible to automate this application overnight, a long-term experiment could not be
performed. However, in a 28 hours experiment the reactor voltage was controlled
below 4 V for approximately 8 hours. During evening and night, the reactor voltage
increased to 10 V. Upon applying the pressure difference, the other morning, the
reactor voltage could be restored to its original value and controlled below 4 V
(Figure 5.39). Clearly, the scaling on the anion exchange membrane is a very loose
precipitate, which can easily be removed. Bulging of the membrane as a result of the
application of a pressure difference could be a handy tool in this respect. A
disadvantage of this method could be size of the particles resulting from the breakdown of the scaling. Therefore, it would be interesting to determine the optimal
frequency to apply the pressure difference and thus control particle size. It could be
128
29.45
27.90
26.35
24.80
23.25
21.70
20.15
18.60
17.05
15.50
13.95
12.40
10.85
9.30
7.75
6.20
4.65
3.10
1.55
10
9
8
7
6
5
4
3
2
1
0
0.00
Voltage (volts)
possible that the scaling layer is not thick enough to break down if the interval
between the pressure drops is too short. On the other hand, when the interval is too
long it could be possible that the scale particles are too large to leave the reactor.
Additional research and development efforts concerning this subject in order to make
the method applicable in industrial systems is therefore recommended. Although it
was not possible to carry out a mass balance, magnesium fluoride seemed to have
accumulated in the anolyte compartment. It should be realised that this could also be
the result of the low pumping capacity.
Time (hrs)
Figure 5.39 Voltage-time plot of the third experiment with the up-scaled reactor.
When the difference between surface fouling and internal membrane fouling, as
depicted in Figure 5.1, is kept in mind while interpreting Figure 5.39, it will be clear
that surface fouling almost exclusively causes the increase of reactor voltage. Even
when the reactor voltage has increased to its maximum value of 10 V (equipment
limitations), a single pressure drop could cause the scaling layer to break down and
the reactor voltage to return to its normal value. Internal membrane fouling does not
seem to contribute to the increasing reactor voltage, as it is not expected that internal
membrane fouling can be removed by bulging the membrane.
In Table 5.10, the initial composition of the anolyte, ampholyte and catholyte
solutions are provided for the three experiments described above. When it became
clear that it would not be possible to perform a long-term experiment, it was decided
to decrease the concentration of the ampholyte solution. Furthermore, the particle
concentration of magnesium fluoride was increased from 1 g/l to 5 g/l in order to
stimulate the precipitation onto particles rather than on the membrane surface. In the
second experiment the influence of hydroxide in the ampholyte was investigated.
129
Table 5.10 Initial compositions of membrane cell reactor compartments.
Initial composition
anolyte
0.3M MgSO4
0.4M ZnSO4
5 g/l MgF2
0.3M MgSO4
0.4M ZnSO4
5 g/l MgF2
0.3M MgSO4
0.4M ZnSO4
5 g/l MgF2
1
2
3
Initial composition
ampholyte
0.45M NaF
Initial composition
catholyte
0.1M NaOH
0.45M NaF
0.05M NaOH
0.1M NaOH
0.5M NaF
0.1M NaOH
After each experiment, samples of the scaling were dried overnight (at 378K) and
analysed with X-ray fluorescence spectrometry and X-ray diffraction. The results are
summarised in Table 5.11 and Table 5.12 respectively. The fluoride content of the
anolyte was measured at the end of each experiment (not during experiments) and
appeared to vary between 150 and 250 ppm.
Table 5.11 XRF analyses of the scaling of the anolyte compartment (wt.%).
1
2
3
Mg
29.9
27.3
29.6
F
49.6
37.6
40.8
Zn
4.6
13.4
9.3
Sx
0.4
2.0
0.8
Ca
0.9
0.9
0.7
K
0.4
0.7
0.6
Mn
0.1
0.2
0.2
Na
0.3
0.5
0.5
LOI
13.6
17.3
17.4
Table 5.12 Phases present in the scaling of the anolyte compartment.
1
2
3
phases present (XRD)
MgF2, ZnOHF, KZnF3, NaZnF3
MgF2, zincite, Mg(OH)2, KMgF3
MgF2, zincite, ZnOHF, KMgF3
Not surprisingly, magnesium fluoride is the major component in the scaling.
However, ZnOHF and zincite (ZnO) were also present in the scaling. In experiment 2,
even magnesium hydroxide has been found. As for magnesium hydroxide to
precipitate, pH>9 is a requirement, it can be concluded that pH control (set point
pH=4 was installed) was not successful. Obviously, local supersaturation could
develop in the membrane cell, which resulted in the precipitation of magnesium
hydroxide and zinc hydroxide compounds. It seems legitimate to assume that an
increase of the turbulence (e.g. Re>10000) would result into more adequate mixing of
solution in the anolyte compartment. In that case, precipitation of magnesium
hydroxide and zinc hydroxide compounds is not expected anymore. The compositions
of these (unwashed) precipitates once again clearly indicate that a dilute sulphuric
acid wash of the magnesium fluoride precipitate is imperative in order to reduce the
130
zinc content. The occurrence of small amounts of NaZnF3 and KZnF3/KMgF3 can be
attributed to the storage of the membranes in KCl or NaCl solution. As the potassium
and sodium concentration in zinc electrolyte is usually very low, these precipitates are
not expected in plant operations.
5.8
Conclusions
The following conclusions could be drawn from the detailed experimental
investigation of the proposed electrodialysis process:
•
It is possible to use zinc electrolyte containing magnesium fluoride particles as
anolyte solution in the proposed electrodialysis process.
•
Transference numbers for sodium and fluoride ions across the membranes are
typically 0.90 tot 0.92, which indicates that the proposed process makes rather
efficient use of the supplied charge.
•
When a mixture of sodium fluoride and sodium hydroxide is fed to the ampholyte
compartment, hydroxide ions are preferentially transported. As hydroxide ions
cause scaling problems in the anolyte compartment, neutralisation of the
NaF/NaOH mixture with hydrofluoric acid should be considered.
•
The zinc anode suffers from acidic dissolution according to the following
reaction:
Zn + 2H+ → Zn2+ + H2↑
The acid attack also causes the release of small zinc grains, although this is a
minor effect. In order to counteract this process, the pH should be controlled
above 2. On the other hand, the pH should be kept below 4.5 in order to prevent
the precipitation of zinc hydroxide compounds. The excess zinc consumption was
33% under experimental conditions used.
•
Water is transported from the ampholyte to the anolyte compartment.
•
The maximum current density in the electrodialysis process largely depends on
the ohmic resistance of the solutions. Therefore, the distance between electrodes
and membranes should be minimal.
•
Without opposing measures the anion exchange membrane tends to scale due to
the development of local supersaturation , thus increasing the required potential
difference. The following measures could be identified to oppose scaling:
- create a turbulent environment by vigorously stirring the anolyte solution;
- create a pressure difference over the membrane in order to break down the
scaling layer;
- brush the membranes;
- supply dilute sulphuric acid pulses.
131
•
Creating a turbulent regime with Re=2787 is insufficient for the prevention of the
formation of scaling. The Reynolds number must be increased (by increased fluid
flow) and probably additional measures (as described above) remain necessary.
•
The reactor voltage could be controlled at 3.5 V at a current density of 100 A/m2.
•
The anion exchange membrane experiences fouling, although the fouling did not
influence the functioning of the membrane for at least 500 hours.
132
6
CONCLUSIONS AND PROCESS
INTEGRATION
The objective of this chapter is to integrate the results of the previous chapters into a
comprehensive flowsheet of the proposed magnesium removal process. Some
important process assumptions and experimental findings of the previous chapters are
summarised below and serve as the basic principles for the construction of the
flowsheet.
6.1
Chapter 1
•
Electrolytic zinc plants tend to experience accumulation of dissolved magnesium
in process liquors, as the natural magnesium bleed usually does not balance the
input from concentrates.
•
High magnesium levels are unacceptable as they hamper the dissolution of calcine
in leach solutions and furthermore contribute to an increased voltage drop in the
cellhouse and increased viscosity (and thus pumping requirements) of process
liquors.
•
A widely employed method to force magnesium bleed is the so-called basic zinc
sulphate precipitation process, which produces a large amount of waste gypsum.
•
If the dumping of gypsum on a zinc plant site would be prohibited, the
development of an alternative magnesium removal process is imperative as
gypsum disposal in (equivalents of) C2-deponies is very expensive.
•
The first zinc plant that faced such a ban on the dumping of gypsum, Pasminco
Budel Zink, was able to integrate magnesium removal with an existing biological
wastewater treatment facility. However, the investment costs of such a facility are
very high and will not be made solely for the removal of magnesium. Besides, the
operating costs of magnesium removal in the biological wastewater plant are also
high.
•
Literature descriptions of alternative magnesium bleed methods (especially
carrying out the Basics Process with calcine instead of slaked lime) require very
specific plant operations, which might be impossible to realise. Besides,
investment costs of these processes are expected to be high.
6.2
•
Chapter 2 and 3
It is possible to remove magnesium from zinc electrolyte by selective precipitation
of magnesium fluoride (sellaite), which is brought about by the slow addition of
zinc fluoride solution.
•
Optimal conditions are between pH=4-4.5 and near the boiling temperature of the
solution.
•
After solid/liquid separation, the wet magnesium fluoride contains entrapped
electrolyte solution. If the precipitate is not washed (with dilute sulphuric acid),
zinc and sulphate will contaminate the precipitate.
•
The fluoride concentration of the resulting magnesium depleted solution is too
high for immediate application in an electrolytic zinc plant. Experiments as well
as thermodynamic modelling indicate fluoride levels of more than 100 ppm in
return solutions, which is unacceptable. High fluoride levels might corrode the
aluminium starting sheets in the cellhouse, which in turn could cause problems
with the stripping of zinc. Acceptable fluoride levels are below 10 ppm.
•
The optimal magnesium removal percentage, with respect to minimising residual
fluoride, is approximately 70%.
•
It is possible to reduce the residual fluoride concentration by addition of reagents,
such as jarosite, goethite, (freshly precipitated) aluminium hydroxide, aluminium
anodising waste sludge or titanyl sulphate dehydrate.
•
For environmental reasons, the destination of the fluoride containing waste
products should be taken into account when selecting a reactive agent.
•
Discarding the fluoride containing solution to jarosite or goethite precipitation,
followed by treatment of the fluorinated iron compound (possibly with HF
recovery) could be considered the most attractive option.
•
Magnesium fluoride experiences hydroxyl incorporation, which influences the
details of the chemical requirements in the flowsheet.
6.3
Chapter 2 and 4
•
It is possible to convert magnesium fluoride (sellaite) to magnesium hydroxide
(brucite) by contacting it with sodium hydroxide solution.
•
However, when a stoichiometric amount of NaOH is added, a complete
conversion is not realised. The magnesium hydroxide product still contains
approximately 10 wt.% fluorine. Even if a large excess is used (Mg:OH=1:50), the
magnesium hydroxide still contains approximately 2 wt.% fluorine.
•
It is not possible to reduce the residual fluoride content in the magnesium
hydroxide product to <1 wt.% by optimising operating conditions such as stirring
velocity, reaction temperature and leach concentration.
•
Magnesium hydroxide has a wide range of applications. Perhaps, the converted
product can be sold despite the residual fluoride content.
134
•
The residual fluoride is incorporated into the brucite crystal structure. X-ray
diffractograms do not show any sign of sellaite (MgF2).
•
It is possible to calcine the (fluoride containing) magnesium hydroxide to
magnesium oxide (with <1 wt.% F) at temperatures of >1273K. Thus, a saleable
product can be obtained. The gaseous HF should be recovered and recycled to the
magnesium removal process.
•
Both from the literature and from preliminary experiments it is inferred that direct
conversion of magnesium fluoride to magnesium oxide under the continuous
addition of gaseous water is possible.
•
As the reaction rate is greatly influenced by specimen geometry, the determination
of the required conditions to accomplish complete conversion still requires further
research.
6.4
Chapter 2 and 5
•
Zinc fluoride and sodium hydroxide solutions can be regenerated from the
resulting sodium fluoride solution in a membrane cell reactor.
•
In order to counteract scaling of the anion exchange membrane, excess hydroxide
in the ampholyte solution can be neutralised with hydrofluoric acid, which has to
be added anyway in order to maintain the mass balance.
•
It is possible to use zinc electrolyte containing magnesium fluoride particles as
anolyte solution.
•
Due to acidic dissolution, consumption of zinc anodes was 33% higher than
stoichiometrically would be expected, under used experimental conditions.
•
Water is transported from the ampholyte to the anolyte compartment.
•
The distance between electrodes and membranes should be minimised in order to
reduce ohmic resistance of the reactor.
•
As the anion exchange membrane tends to scale due to the development of local
supersaturations, opposing measures have to be applied, such as
- vigorously stirring the anolyte solution;
- periodically creating a pressure difference over the membrane in order to
break down the scaling layer;
- periodically brushing the membranes;
- periodically supplying dilute sulphuric acid pulses to the anolyte solution.
•
Creating a turbulent regime with approximately Re=2800 (by increased pumping)
is insufficient for the prevention of the formation of scaling. The Reynolds
135
number must be increased and probably additional measures (such as described
above) remain necessary.
•
The reactor voltage can be controlled at 3.5 V at 100 A/m2.
•
The anion exchange membrane experiences fouling, although the fouling did not
influence the functioning of the membrane for at least 500 hours.
6.5
Process Integration
A zinc plant of the size of Pasminco Budel Zink processes 400 000 tpa of concentrate,
which implies that with an average magnesium content of 0.2 wt.% (Table 1.1) 800
tpa of magnesium has to be treated. However, as mentioned in Section 1.5, the
biological desulphurisation (BDS) unit of Pasminco Budel Zink only treats 0.5 m3/h
magnesium bleed. This implies that the amount of removed magnesium is only 50 tpa.
Although the total magnesium bleed is higher as a result of selling purified solution,
representatives of Pasminco Budel Zink[22] argued that 800 tpa magnesium removal is
far too high. Therefore, a flowsheet for the magnesium removal process has been
designed for an electrolytic zinc plant assuming a magnesium bleed of 250 tpa.
Integration of the three unit processes in the above-mentioned flowsheet is depicted in
Figure 6.1.
2.2 m3/h 0.1 M NaF
0.018 M ZnF2
2.63 M ZnSO4
0.158 M MgSO4
25 g/l MgF2-particles
175 ppm F
purified
solution
2.9 m3/h
2.1 M ZnSO4
0.56 M MgSO4
Acid addition
0.5 kg HF/h
(5 tpa)
63 m3/h
AEM
F101 kg Zn/h
MgF2
precipitation
2.9 m3/h
2.5 M ZnSO4
0.158 M MgSO4
175 ppm F
fluoride
removal
Fluorinated
product
Na+
NaF
zinc
2.2 m3/h
1 M NaF
2.63 M ZnSO4, 0.158 M MgSO4
175 ppm F
2.2 m3/h
0.82 M NaF
0.18 M NaOH
L
S
water
(e.g. 1 m3/h)
MgF1.9OH0.1
72 kg/h
recycled
wash water
CEM
NaOH
ZnF2
63 m3/h suspension
(25 g/l MgF2)
Purified solution
back to plant
(<10 ppm F)
H2
2 kg/h
Washing
Wash water (e.g. 1 m3/h)
to fluoride removal or
MgF2 precipitation
2.2 m3/h
0.9 M NaOH
0.1 M NaF
L
Conversion
S
3
Water (e.g. 0.5 m /h)
MgOH1.65F0.35
68 kg/h
Washing
wash water (e.g. 0.5 m3/h)
to ampholyte
HF(g)
8.1 kg/h
Burning
>10000C
MgO
47 kg/h
Figure 6. 1 Flowsheet of the proposed magnesium removal process.
Below, some important aspects and possible variations of this flowsheet are
discussed:
136
•
The zinc electrolyte input for the magnesium fluoride precipitation is modelled as
2.1M zinc sulphate (137 g/l Zn) and 0.56M magnesium sulphate (13.6 g/l Mg),
which is in accordance with XRF solution analyses and information from
representatives of Pasminco Budel Zink.[22]
•
The fluoride input is modelled as 0.018M zinc fluoride in the slurry recycle
(which further has the composition as provided in Figure 6.1). However, it should
be realised that this modelling is hypothetical and is only done for reasons of
clarity. As long as there is dissolved magnesium present in the recycling slurry,
magnesium fluoride immediately will start to precipitate.
•
The slurry anode concept (i.e. recycling a large volume of zinc electrolyte from
the magnesium fluoride precipitation to the anolyte compartment of the membrane
cell reactor) prevents the continuous inlet of water into the anolyte compartment.
Thus, a possible available evaporation excess of the zinc plant does not have to be
“used”. The recycle stream is large in order not to exceed the solubility products
of zinc sulphate and zinc fluoride when zinc of the zinc anodes enters the solution
in the anolyte compartment. However, the precise volume of the proposed recycle
stream was not optimised and was arbitrarily chosen.
•
It will be clear that it is also possible to operate the membrane cell reactor with a
continuous intake of water if the “use” of the evaporation excess is permitted.
Thus, scaling of the anion exchange membrane could be prevented. When a
concentration of 0.12M ZnF2 is selected, a continuous water intake (and
evaporation excess) of 9.1 m3/h is required for 250 tpa magnesium removal. With
this type of operation, the slurry anode concept is not employed and the flowsheet
changes to Figure 2.7.
•
The fluoride content of the magnesium depleted solution is assumed to be
175 ppm. The thermodynamic model described in Section 3.2 predicts 75 ppm,
while the experimental determined solubility of magnesium fluoride in zinc
electrolyte indicates 250 ppm F (Section 3.2.3).
•
The magnesium depleted, residual fluoride containing solution is sent to “fluoride
removal”, for example jarosite precipitation. In Section 3.3, the removal efficiency
for jarosite particles was determined between 0.6 and 1.5 ppm. Although it may be
expected that the efficiency of fluoride incorporation into the jarosite structure
will be much higher when the fluoride containing solution is directly fed to the
jarosite precipitation, an average removal efficiency of only 1 mg F/g jarosite is
assumed. Then, an annual amount of 6300 tpa jarosite is required, which is far less
than the usual annual jarosite product of an electrolytic zinc plant.
•
Obviously, recalling Section 3.3, there are other options such as sending the
fluoride containing solution to goethite precipitation or adding reactive
compounds (e.g. aluminium hydroxide, aluminium anodising waste sludge etc.).
•
The selection of the fluoride removal method should, amongst others, depend on
the destination of the resulting fluoride containing product. Preferably, the
incorporated fluoride should be recovered.
137
•
For example, a suitable jarosite treatment method should be available, as it is
unlikely that dumping of gypsum would be prohibited, while dumping of jarosite
would still be tolerated.
•
The washing of magnesium fluoride precipitate with dilute sulphuric acid also
requires some water intake and use of evaporation excess. Although it was
demonstrated that reducing the zinc and sulphate content of the magnesium
precipitate is possible by washing with dilute sulphuric acid, the minimal amount
of wash liquor (required per kg removed magnesium) was not determined. In the
flowsheet, the arbitrary amount of 1 m3/h is suggested.
•
The amount of zinc and sulphate, which is recycled from the washing of
magnesium fluoride to either the fluoride removal or magnesium fluoride
precipitation is not modelled in the flowsheet of Figure 6.1.
•
If it would be possible to sell the magnesium fluoride, the operating costs would
significantly be reduced as the investment and operating costs for the conversion
step and the membrane cell reactor are thus avoided. In this respect, in an
interesting paper, Layne et al.[81] describe the use of (molten) magnesium fluoride
to extract aluminium from scrap alloy. Firstly, gaseous AlF and Mg are produced
by reacting the alloy with molten MgF2. Upon cooling, immiscible molten layers
of purified Al and MgF2 are formed:
2Al(alloy) + MgF2(l)
2AlF(g) + Mg(g)
→ 2AlF(g) + Mg(g)
→ 2Al(l) + MgF2(l)
This approach would provide an attractive waste-to-waste technology. However,
the hydroxyl incorporation into the magnesium fluoride could be a problem
(aluminium oxide formation) and the purity of the magnesium fluoride should
probably be improved, for example by more thorough washing of the precipitate.
•
Reportedly, some aluminium producers add small amounts of magnesium fluoride
(up to 6%) to the cryolite electrolyte in the conventional Hall-Héroult process.[82]
However, again it should be realised that the presence of oxygen could limit this
application. Even if all residual entrapped solution is evaporated, the incorporated
hydroxide (MgF1.9OH0.1) still contributes to 2.5 wt.% oxygen in the magnesium
fluoride product. Besides, it will be clear that any residual zinc and sulphate
should be removed from the magnesium fluoride product by thorough washing.
•
Although dumping of magnesium fluoride is not an environmentally sound
solution, it should be realised that the amount of waste is dramatically reduced
when compared with the gypsum producing process.
•
If complete conversion of magnesium fluoride to magnesium oxide can be
realised under realistic conditions (i.e. in terms of economic consequences of
required reaction temperature, reaction time and adjustment of specimen
geometry), the production of magnesium oxide could be a viable process route.
138
•
If these conditions are unrealistic or unknown or if applications for magnesium
fluoride cannot be found and dumping is also considered unacceptable, it could be
decided to further process the magnesium fluoride as depicted in Figure 6.1. The
composition of the magnesium hydroxide conversion product is based on the
conversion experiments described in Chapter 4.
•
The converted product is washed to remove, amongst others, residual sodium
traces. As unwashed magnesium hydroxide contains a relatively small amount of
sodium (Table 4.8), the amount of wash liquor is assumed to be less than the wash
liquor required after magnesium fluoride precipitation. An arbitrary amount of 0.5
m3/h has been selected for the flowsheet depicted in Figure 6.1. The NaF/NaOH
containing wash liquor can be recycled to the ampholyte compartment.
•
Perhaps an application for fluoride containing magnesium hydroxide can be
found, which would prevent the necessity to burn the brucite to periclase.
•
If it is not possible to sell fluoride containing magnesium hydroxide, calcining to
magnesium oxide with HF-recovery is proposed. Recovered HF can be used to
neutralise the residual NaF/NaOH mixture of the conversion unit process.
•
Although it was observed that the ampholyte compartment suffers from water
depletion, the amount of water transported is very small (typically 10 l/h when test
results are extrapolated to flowsheet requirements) and therefore has been omitted
in the flowsheet.
6.6
Recommendations
From this set of conclusions it can be deduced that the proposed process (as suggested
in Figure 6.1) is technically feasible. At this point, the following step would be the
construction of a pilot plant. In order to decide whether such an investment is
acceptable, a rough economic evaluation of the process is imperative. A conceptual
process design of the flowsheet depicted in Figure 6.1 was performed[83]. Although
the process design study was based on a lot of assumptions and uncertainties, the
result can be considered to be a useful indication of the order of magnitude of
investment and operating costs. The total investment costs were estimated at
€ 3,500,000.00 and the operating costs at € 8.60/kg removed magnesium. This
calculation was carried out using standard cost estimation methods such as described
by Peters and Timmerhaus.[84] In Appendix A, these figures are justified by providing
the working method as well as key tables of the process design study. However, it is
stressed once again that these costs are indicative and only provide an idea of the
order of magnitude of investment and operating costs.
In order to justify the construction of a pilot plant, it was considered necessary to
recover the investment costs within three years. Based on the investment costs of
€ 3,500,000.00 and the amount of magnesium to be removed (250 tpa), it can be
deduced that the newly proposed process should save € 4.70 per kg removed
magnesium. Compared to the operating costs of the biological desulphurisation unit of
Pasminco Budel Zink, this condition could not be fulfilled. Therefore, Pasminco
139
Budel Zink decided not to further investigate replacement of the BDS with the
proposed magnesium removal process.
Other zinc plants usually employ the basic zinc sulphate process. The typical
operating costs of the basic zinc sulphate process are estimated at € 3.00 per kg
removed magnesium[22] if the disposal of gypsum on the plant site is permitted.
Therefore, the newly proposed process is not of interest for replacement the basic zinc
sulphate process at other zinc plants if dumping of gypsum on the plant site is
allowed.
However, as international environmental awareness is still growing, it may be
expected that, like Pasminco Budel Zink, other zinc plants will face a ban on the
dumping of gypsum at the plant site at some point in future. At that time, it seems
legitimate to assume that disposal costs for gypsum have increased to the level of the
Dutch dump sites (“C2-deponies”), which are at present € 658.00/tonne[85]. As the
removal of 1 kg magnesium typically creates 27 kg of waste gypsum[22], the operating
costs will increase significantly as a result of increased dumping costs. An amount of
27 kg waste gypsum has to be disposed of in a C2-depony at a cost of € 17,80. Thus
operating costs are increased to € 19,80 per kg removed magnesium. In this situation
the proposed magnesium removal process (€ 8.60/kg removed magnesium) becomes
an interesting alternative. It can easily be calculated that in this situation the
investment costs of € 3,500,00.00 are recovered after 15 months. Therefore, in this
situation it can be recommended to build a pilot plant to further investigate the
operating conditions of the process.
A variant of the newly proposed process is to dump magnesium fluoride at a relatively
high cost (and omit the conversion step as well as the membrane cell reactor).
Although this is not the most attractive option from an environmental point of view, it
should be realised that the amount of material to be dumped is far less than in the
(gypsum producing) basic zinc sulphate process route. Consequently, operating costs
are also lower. In Appendix A, investment costs as well as operating costs are derived
for this process variant. Not surprisingly, the investment costs are significantly lower
when the conversion of magnesium fluoride to magnesium hydroxide and subsequent
pyrohydrolysis as well as regeneration of chemicals in the membrane cell reactor are
left out of the flowsheet: € 975,741.00. The operating costs are € 6.15 per kg removed
magnesium. The operating costs are mainly determined by the costs for disposal of
magnesium fluoride in a C2-depony and the purchase costs of fluoride. Under these
conditions, the operating costs of this process variant are lower than the process in
which magnesium fluoride is converted to magnesium hydroxide and subsequently
magnesium oxide with regeneration of chemicals. However, the difference is small
and it should be realised that dumping fees of C2-deponies have increased rapidly
over the past few years.
Another promising variant is to directly convert magnesium fluoride to magnesium
oxide. Both from the literature and from preliminary experiments it is inferred that
direct conversion of magnesium fluoride to magnesium oxide under the continuous
addition of gaseous water is possible. However, the reaction rate is greatly influenced
by specimen geometry. Therefore, it should be investigated under which conditions
bulk amounts of magnesium fluoride can be converted to magneisum oxide with
140
accpetable residual fluoride content. Subsequently, it should be investigated whether
creating these conditions is legitimate from an economic point of view.
Summarising, the newly proposed process is not of interest for replacement of the
existing facility at Pasminco Budel Zink nor for replacing present Selective
Precipitation (basic zinc sulphate) processes at other zinc plants. However, when
dumping of gypsum at other zinc plant sites becomes prohibited, the Selective
Precipitation process becomes too expensive and (one of the variants of) the new
process becomes a promising alternative and it can be recommended to build a pilotplant to determine the site specific operating conditions.
141
APPENDIX A
Investment costs and operating costs of the proposed process.
A conceptual process design study[83] was carried out in order to indicate the
investment costs as well as operating costs of the proposed process. A block scheme
of the process was translated into unit operations from which a list of major
equipment requirements was derived. A summary of this list as well as the associated
purchase costs[86] are provided in Table A.1.
Table A.1 Major equipment costs.
Unit
Storage Tanks
Cyclone
Sedimentation centrifuge
Filtration centrifuge
Fluidised bed
Membrane reactor:
- Crane
- Housing
- Foundation
- pumps
- additional (10%)
- Total membrane reactor
Rotary kiln
Thickeners
Pumps
Mixers
Transport screw
Total (2000)
Inflation Rate (4%/yr)
Total (2002)
Price (Euro)
21,870.00
10,244.00
32,520.00
22,683.00
7,050.00
(110,678.00)
(24,863.00)
(3,770.00)
(10,182.00)
(3,881.00)
153,374.00
408,862.00
76,000.00
146,682.00
92,602.00
4,649.00
976,536.00
79,686.00
1,056,222.00
A factorial method, known as the Lang method[84], was used to derive other fixed
capital costs (such as buildings, storage, piping) from the estimation of the costs of
major equipment. Also the working capital was derived as a fraction of fixed capital
costs. Thus, total investment costs are obtained as summarised in Table A.2.
Table A.2 Total investment costs for the proposed process.
Cost centre
Major equipment
Lang factor
Fixed capital costs
Working capital (15% of fixed capital costs)
Total investments
Price
€ 1,056,222.00
2.9
€ 3,063,044.00
€ 459,457.00
€ 3,522,501.00
Some of the operating costs (maintenance, miscellaneous materials, taxes and
royalties) are also derived as a fraction of the total investment costs. Operating costs
are specified in Table A.3. For labour costs, it was assumed that one person is
continuously needed for process control. In a continuously operating zinc plant this
implies that five employees are needed. It was assumed that the operator is assisted
by, on average, 2.4 persons. Supervision is assumed to be 20% of the costs of
operating labour. Chemical requirements are mainly zinc anodes and some
hydrofluoric acid. For zinc anodes, only the production costs of producing the
required amount of zinc from purified zinc sulphate solution are incorporated. Anion
and cation exchange membranes can be purchased at a cost of € 218.00 per square
meter.[87] It was assumed that the service life of a membrane is eight months when the
membrane cell reactor is operated at 3.5 V and 100 A/m2. Energy costs are mainly
based on requirements of the membrane cell reactor.
Table A.3 Operating costs of the proposed process.
Cost centre
Chemicals, membranes and energy
Maintenance (8% of total investment)
Misc. materials (10% of maintenance)
Labour costs:
- process control (5 workers)
- assistance (12 workers)
- supervision (20% of operating
labour)
Total labour costs
taxes (2% of total investment)
royalties (1% of total investment)
total operating costs per year
costs per kg Mg
Price
€ 1,009,000.00
€ 281,800.00
€ 28,180.00
(250,000.00)
(360,000.00)
(122,000.00)
€ 732,000.00
€ 70,000.00
€ 35,000.00
€ 2,155,980.00
€ 8,60
Investment costs and operating costs of magnesium fluoride precipitation, followed by
dumping the magnesium fluoride.
A variant of the newly proposed process is to dump magnesium fluoride at a relatively
high cost (and omit the conversion step as well as the membrane cell reactor). In order
to estimate the investment costs and operating costs of this process variant, the
equipment list of Table A.1 is used. All equipment for converting magnesium fluoride
to magnesium hydroxide and the subsequent pyrohydrolysis to magnesium oxide is
left out as well as all equipment associated with the membrane cell reactor. The
storage capacity (for zinc fluoride) was increased as the recycling line through the
anolyte compartment for fluoride addition disappeared. The list of major equipment
required for the precipitation of magnesium fluoride from purified zinc sulphate
solution is provided in Table A.4.
144
Table A.4
Major equipment costs for the process variant with dumping of
magnesium fluoride.
Unit
Storage Tanks
Sedimentation centrifuge
Thickeners
Pumps
Mixers
Total (2000)
Inflation Rate (4%/yr)
Total (2002)
Price (Euro)
61,601.00
32,520.00
76,000.00
44,566.00
55,816.00
270,503.00
22,073.00
292,576,00
Again, the Lang method is used to derive other fixed capital cost as well as the
working capital. Thus, total investment costs are obtained as summarised in Table
A.5.
Table A.5
Total investment costs for the process variant including dumping
magnesium fluoride.
Cost centre
Major equipment
Lang factor
Fixed capital costs
Working capital (15% of fixed capital costs)
Total investments
Price
€ 292,576.00
2.9
€ 848,470.00
€ 127,271.00
€ 975,741.00
The fluoride content of the precipitate is lost in this process variant. Therefore,
chemical costs increased significantly. The removal of 1 kg of magnesium typically
creates 2.5 kg of (dry) magnesium fluoride precipitate, which has to be dumped in a
C2-depony at a cost of € 658,00/tonne[85]. For the proposed process (removing 250 tpa
Mg) this implies an amount of magnesium fluoride of 641 tpa to be dumped. In Table
A. 6, the operating costs of this process variant are listed.
Table A. 6 Operating costs for the process variant including dumping
magnesium fluoride.
Cost centre
Chemicals (ZnF2) and energy
Maintenance (8% of total investment)
Misc. materials (10% of maintenance)
Total labour costs
taxes (2% of total investment)
royalties (1% of total investment)
dumping costs
total operating costs per year
costs per kg Mg
Price
€ 750,000.00
€ 78,059.00
€ 7,059.00
€ 250,000.00
€19,515 .00
€ 9,757.00
€ 421,743.00
1,536,132.00
€ 6.15
145
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153
Summary
Magnesium removal in the electrolytic zinc industry
J.L. Booster
Electrolytic zinc plants need to take measures to control the magnesium content in
their process liquors, because the natural magnesium bleed does not balance the input
from concentrates. The most important disadvantage of too high a magnesium content
is that less zinc can be dissolved in the process liquor, which contributes to an
increase in energy costs. A common method to bleed magnesium is selective
precipitation of the process liquor as basic zinc sulphate. In this process, magnesium
remains dissolved. This magnesium-containing solution, generally after a cleaning
step, can be discarded. A by-product of the process is gypsum, usually dumped on the
plant site. In The Netherlands, dumping of gypsum on plant sites is prohibited and it
seems reasonable to assume a similar ban on the dumping of gypsum in a lot of other
countries in the near future. Contrary to the removal process described above,
Pasminco Budel Zink, the only zinc refinery of The Netherlands, has integrated its
magnesium removal with an existing biological wastewater-treatment facility. The
investment costs as well as the operating costs of such a facility are very high and will
not be made solely for the removal of magnesium. It is unlikely that in future other
zinc plants will apply such an integration of removal processes. Therefore, the
development of an alternative process for the removal of magnesium from zinc
sulphate solutions is useful.
It proved to be possible to remove magnesium from zinc electrolyte by selective
precipitation of magnesium as magnesium fluoride. The precipitate is obtained by the
slow addition of zinc fluoride solution to the process liquor. Optimal process
conditions are between pH=4.0-4.5 and just near the boiling temperature of the
solution. During precipitation, the magnesium fluoride appears to substitute fluoride
ions partly by hydroxide ions. If the precipitate is not washed with dilute sulphuric
acid, zinc and sulphate from the pore water will remain behind, permanently
contaminating the product.
The fluoride content of the residual solution was found to be too high for immediate
application in the zinc plant. Both experiments as well as thermodynamic modelling
indicate that after precipitation the fluoride content will never be lower than 100 ppm.
Such a fluoride concentration results in the corrosion of aluminium starting sheets in
the cellhouse of the zinc plant. The optimal magnesium removal percentage with
respect to minimising residual fluoride appears to be approximately 70%.
It is possible to remove fluoride from process liquor by addition of compounds
capable of adsorbing fluoride such as jarosite, goethite and aluminium hydroxide. For
environmental reasons, the subsequent treatment method of the fluoride-containing
product should play an important role in selecting the compound. A process whereby
hydrofluoric acid can be regenerated is the most logical option.
How to deal with the magnesium fluoride? In the literature, applications for
magnesium fluoride in the aluminium industry are described. However, these are not
standard applications and therefore it is not sure if bulk amounts of magnesium
fluoride can be disposed of. Although the mass of the magnesium fluoride is
considerably less than the mass of the gypsum produced in the conventional process,
dumping of magnesium fluoride is undesirable for environmental reasons and
possibly too expensive as well. Therefore, a treatment method was developed to
convert the magnesium fluoride into a saleable product.
It proved to be possible to convert magnesium fluoride to magnesium hydroxide by
contacting it with sodium hydroxide solution. Magnesium hydroxide, if pure enough,
has a wide range of (industrial) applications. The conversion process produces a
sodium fluoride solution as process residue. Complete conversion is not realised,
however, when a nearly stoichiometric amount of sodium hydroxide is added. Even
when a large excess of sodium hydroxide is added, a few wt.% of fluorine remain in
the magnesium hydroxide product. X-ray diffraction analysis indicates that the
fluoride is incorporated in the crystal structure of brucite (magnesium hydroxide), as
magnesium fluoride is not found as a separate phase anymore. It appears to be
possible to calcine the fluoride-containing magnesium hydroxide at a high
temperature (>1000°C) to sufficiently pure periclase (magnesium oxide). In principle,
the hydrofluoric acid produced in this calcination can be captured and recycled to the
magnesium removal process.
With a modified electrodialysis reactor, the sodium fluoride solution obtained from
the conversion process described above, can be converted into reagents, i.e. zinc
fluoride and sodium hydroxide required for the previous process steps. The reactor
consists of three compartments separated from each other by an anion and a cation
exchange membrane. When a potential difference is applied over the electrodes of this
membrane reactor, fluoride ions will move towards the anode and pass the anion
exchange membrane. The fluoride ions end up in the compartment in which the
anode, a zinc sheet, is present. The zinc sheet slowly dissolves as a result of the
passage of current. Thus, a zinc fluoride solution is produced. Sodium ions move
towards the negative cathode and pass the cation exchange membrane. They reach the
compartment where the cathode is present. At this negative cathode, hydroxide ions as
well as hydrogen gas are produced. Thus a sodium hydroxide solution is produced. By
side reactions the consumption of zinc anodes is approximately 33% higher than
would be expected stoichiometrically.
Without appropriate measures, at the anion exchange membrane a transport-limiting
effect called scaling occurs. In order to reduce this scaling, the sodium fluoride
solution, which always contains some sodium hydroxide, should be neutralised with
hydrofluoric acid. In addition, any of the following measures appears to be effective
in reducing the scaling of the anion exchange membrane:
•
•
•
creating a strong turbulent regime in the anolyte compartment;
periodically creating a small pressure difference over the membrane;
periodically brushing the membrane.
156
Especially, periodic application of a small pressure difference seems to be a relatively
easy and cheap method to maintain the reactor voltage constant at a current density of
100 A/m2. It is recommended to further investigate this effective low-tech measure to
counteract scaling in membrane processes.
The research has proved that the new process is technically feasible. An indication of
the economic feasibility is obtained by a rough economic evaluation of the process.
The operational costs of the process appear to be in the same order of magnitude as
the operating costs of the present process applied by Pasminco Budel Zink whereby
magnesium removal is integrated with an existing biological wastewater-treatment
device. The operational costs of the new process are significantly higher than the costs
of the conventional basic zinc sulphate process. However, the operational costs of this
basic zinc sulphate process will increase significantly when the expected ban on the
dumping of gypsum becomes operative and dumping fees must be paid similar to
present costs of disposal in Dutch C2-deponies. In that case, the estimated operational
costs of the new process are expected to be significantly lower than the costs of the
basic zinc sulphate process. Anticipating the more stringent international rules, more
detailed elaboration of the process is therefore recommended.
157
Samenvatting
Magnesiumverwijdering in de elektrolytische zinkindustrie
J.L. Booster
Elektrolytische zinkfabrieken moeten maatregelen nemen om het magnesiumgehalte
in hun procesvloeistoffen onder controle te houden. De natuurlijke magnesiumspui
weegt niet op tegen de input uit concentraten. Het belangrijkste nadeel van een te
hoog magnesiumsulfaatgehalte in de procesvloeistof is dat er minder zink kan worden
opgelost in die procesvloeistof, waardoor de energiekosten blijken toe te nemen. Een
veel gebruikte methode om magnesium te spuien is het selectieve neerslaan van de
procesvloeistof als basisch zinksulfaat. Magnesium blijft bij dit proces achter in de
procesoplossing die eventueel na een zuiveringsstap geloosd mag worden. Bij dit
proces wordt ook gips geproduceerd, dat gewoonlijk op het fabrieksterrein gestort
wordt. In Nederland is het storten van gips op fabrieksterreinen inmiddels verboden
en het lijkt redelijk te veronderstellen dat een zelfde verbod in veel andere landen op
korte termijn ook ingevoerd zal worden. Pasminco Budel Zink, de enige zinksmelter
in Nederland, heeft in afwijking van het hierboven geschetste verwijderingsproces de
magnesiumverwijdering geïntegreerd met een biologische afvalwaterzuiveringsinstallatie. De investeringskosten en operationele kosten van een dergelijke installatie
zijn echter dermate hoog dat ze niet voor magnesiumverwijdering alleen gemaakt
zullen worden. Het valt niet te verwachten dat in de toekomst een zelfde integratie
van verwijderingsprocessen in andere zinkfabrieken toegepast zal worden. Het is
daarom zinvol om een alternatief proces voor de verwijdering van magnesium uit
zinksulfaat- oplossingen te ontwikkelen.
Het is mogelijk gebleken magnesium uit zinkelektrolyt te verwijderen door selectief
neerslaan van dit magnesium als magnesiumfluoride. Het neerslag wordt verkregen
door het langzaam toevoegen van een zinkfluoride-oplossing aan de procesvloeistof.
Optimale procescondities liggen tussen pH=4.0-4.5 en dichtbij het kookpunt van de
procesoplossing. Bij de vorming van het magnesiumfluoride blijkt een gedeeltelijke
substitutie van fluoride door hydroxyl-ionen op te treden. Als het neerslag niet wordt
uitgewassen met verdund zwavelzuur blijft zink en sulfaat uit het poriewater erin
achter, waardoor het neerslag blijvend vervuild wordt met zink en sulfaat.
De oplossing, die na het neerslaan van magnesiumfluoride overblijft, heeft een te
hoog fluoride-gehalte en kan daardoor niet direct ingezet worden in de zinkfabriek.
Zowel uit experimenten als uit een thermodynamische modelstudie valt af te leiden
dat na het neerslagproces nimmer een fluoride-gehalte van minder dan 100 ppm
behaald kan worden. Deze te hoge fluoride-concentratie zou leiden tot corrosie van de
aluminium startplaten in het celhuis van de zinkfabriek. Het blijkt dat de laagste
hoeveelheid fluoride in de resterende oplossing wordt verkregen als ongeveer 70%
van het magnesium uit de procesvloeistof wordt neergeslagen.
Het is mogelijk fluoride uit de procesvloeistof te verwijderen door toevoeging van
stoffen die fluoride aan zich kunnen binden zoals jarosiet, goethiet en aluminium-
hydroxide. Uit milieu-oogpunt zou de methode van opwerken van het op deze wijze
verkregen fluoride bevattende product bij de keuze van de stof waaraan het fluoride
wordt gebonden een belangrijke rol moeten spelen. Een opwerkroute waarbij
waterstoffluoride kan worden teruggewonnen is daarbij de meest logische optie.
Wat te doen met het verkregen magnesiumfluoride? In de literatuur worden
toepassingen van magnesiumfluoride bij de aluminiumproductie beschreven. Het
betreffen hier echter geen standaardtoepassingen en het is daarom onzeker of hier
grote hoeveelheden magnesiumfluoride kunnen worden afgezet. Hoewel de
hoeveelheid magnesiumfluoride in massa aanzienlijk minder is in vergelijking met de
hoeveelheid gips geproduceerd in het conventionele proces, is storten van
magnesiumfluoride uit milieu-oogpunt ongewenst en mogelijk ook te kostbaar. Er is
daarom gezocht naar een manier om het verkregen product te veranderen in een wel
direct afzetbaar product.
Het blijkt mogelijk magnesiumfluoride om te zetten in magnesiumhydroxide door het
in contact te brengen met natronloog. Magnesiumhydroxide, mits voldoende zuiver,
kent een zeer groot aantal (industriële) toepassingen. Bij deze omzetting wordt een
natriumfluoride-oplossing verkregen. Bij een (bijna) stoichiometrische toevoeging
van natronloog wordt er echter geen volledige omzetting gerealiseerd. Zelfs bij een
grote overmaat natronloog blijven er enkele gewichtsprocenten fluoride in het
magnesiumhydroxideproduct achter. Röntgendiffractie-onderzoek wijst uit dat het
fluoride ingebouwd wordt in de kristalstructuur van bruciet (magnesiumhydroxide).
Er wordt namelijk geen magnesiumfluoride als afzonderlijke fase meer gevonden. Het
is mogelijk gebleken dit fluoride bevattende magnesiumhydroxide bij hoge
temperatuur (>1000°C) uit te stoken, zodanig dat voldoende zuiver magnesiumoxide
(periclaas) wordt verkregen. Het bij dit stookproces geproduceerde waterstoffluoride
kan in principe afgevangen worden voor hergebruik.
Met een gemodificeerde elektrodialysereactor kan de bij de hierboven beschreven
omzetting verkregen natriumfluoride-oplossing omgezet worden in voor eerdere
processtappen vereiste reagentia: zinkfluoride en natriumhydroxide. De reactor
bestaat uit drie compartimenten, die van elkaar gescheiden zijn door een anion en een
kation uitwisselingsmembraan. Wanneer een potentiaalverschil over de electrodes van
deze membraanreactor wordt aangebracht, bewegen de fluoride-ionen zich naar de
anode, de positieve elektrode, en passeren daarbij het anion uitwisselingsmembraan.
Ze bereiken dan het compartiment waarin zich tevens de anode, een zinkplaat, bevindt
die tengevolge van de stroomdoorgang langzaam oplost en daarbij zinkionen
produceert. Er wordt simpelweg een zinkfluoride-oplossing geproduceerd.
Natriumionen bewegen naar de negatieve elektrode en passeren het kation
uitwisselingsmembraan. Ze bereiken dan het compartiment waarin zich de kathode
bevindt. Aan deze negatieve elektrode worden zowel hydroxide-ionen als
waterstofgas gevormd, zodat simpelweg een natriumhydroxide-oplossing onstaat.
Door nevenreacties wordt ongeveer 33% meer zink geconsumeerd dan
stoichiometrisch verwacht zou worden.
Zonder passende tegenmaatregelen treedt aan het anion uitwisselingsmembraan een
transport belemmerend effect op, dat met “scaling” wordt aangeduid. Om dit
“scalingproces” tegen te gaan, dient de altijd wat loog bevattende natriumfluorideoplossing in het middelste compartiment van de elektrodialysereactor geneutraliseerd
160
te worden met waterstoffluoride. Verder blijken de volgende aanvullende maatregelen
het terugdringen van “scaling” tot een aanvaardbaar niveau te bewerkstelligen:
•
•
•
het creëren van een sterk turbulent regime;
het periodiek aanbrengen van een klein drukverschil over het membraan;
het periodiek schoonborstelen van het membraan.
Met name het periodiek aanbrengen van een klein drukverschil lijkt een relatief
eenvoudige en goedkope methode om bij een stroomdichtheid van 100 A/m2 het
reactorvoltage op de gewenste waarde te houden.
Het onderzoek heeft uitgewezen dat de nieuwe route technisch haalbaar is. Een
indicatie van de economische haalbaarheid van dit proces is verkregen door een ruwe
economische evaluatie. De operationele kosten van het proces blijken van dezelfde
orde van grootte te zijn als de kosten van het huidige proces toegepast door Pasminco
Budel Zink, waarin magnesiumverwijdering geïntegreerd wordt met een biologische
afvalwaterzuiveringsinstallatie. De operationele kosten van het nieuwe proces zijn
aanzienlijk hoger dan die van het conventionele basisch zinksulfaatproces. De kosten
van het basisch zinksulfaat proces stijgen echter enorm wanneer het te verwachten
verbod op het storten van gips op fabrieksterreinen van kracht wordt en stortkosten
betaald moeten worden die vergelijkbaar zijn met de huidige Nederlandse C2deponiekosten. In dat geval worden de geschatte operationele kosten van het nieuwe
proces juist aanzienlijk lager dan die van het basisch zinksulfaatproces. In afwachting
van de te verwachten strengere internationale regelgeving moet nog meer
gedetailleerde uitwerking van het proces daarom raadzaam geacht worden.
161
ACKNOWLEDGEMENTS
I would like to thank my promotor Markus Reuter for his supervision during this
research project. By stimulating me to consider the project from an engineering
perspective, he not only enhanced the quality of my thesis, but also broadened my
view on science. I am very grateful to Ton van Sandwijk for his contributions to the
principles of the proposed process and pleasant guidance from the beginning to the
end of my Ph. D. study. Together, they really gave me the opportunity for scientific as
well as personal development during the past four years.
I wish to thank Pasminco Budel Zink for supporting this project. I was allowed to use
industrial zinc electrolyte and laboratory equipment. The constructive discussions we
had with Cris Copini and Wiel Vermeulen further contributed to the standard of the
research project. I also would like to thank my new employer, GeoDelft, for giving
me the opportunity to finish my Ph.D. thesis.
The effort of Jack Voncken in characterising the magnesium fluoride precipitate
(resulting in a paper in Powder Diffraction) is gratefully acknowledged. I really
enjoyed the guidance of three graduation students: Guillo Schrader, Mustafa Bayirli
and Gonzalo Rivera. Their Master’s theses were essential contributions to my Ph. D.
thesis. Also the conceptual process design study carried out by Koen Pepels, Selvedin
Telalovic, Vincent Toussaint, Steven van Vegten and Maarten Wichhart formed a
strong basis for the economic evaluation of the proposed process.
I am greatly indebted to the people of the "Analytical Service": Theo Verkroost,
Pascal Visser, Ellen Meijvogel-de Koning, Ron Penners, David Muiderman, Jolanda
van Haagen-Donker and John van den Berg, amongst others for numerous XRF- and
XRD-analyses as well as many fluoride analyses.
A thank you to my fellow AIO’s, especially to my roommates Steven, Willemijn en
Jim, for the great time we had the last couple of years. I already miss room 123 and its
nice discussions over a cup of coffee brewed in an ancient fume cabinet.
Finally, I would like to thank all my friends for always asking: “Are you finished
yet?” and my parents and Iris for their unconditional love and support.
Curriculum Vitae
Surname
First names
Date of Birth
Place of Birth
Booster
Jacco Leendert
December 22th 1973
Maassluis, The Netherlands
2002 – date
GeoDelft
Strategic Research
Project leader
1998 – 2002
Delft University of Technology
Ph. D. student
Faculty of Applied Earth Sciences
Raw Materials Technology
1994 – 1997
University of Utrecht
Student
Faculty of Geographical Sciences
Environmental Sciences
graduation: 31-10-1997
1992 – 1997
University of Utrecht
Student
Faculty of Earth Sciences
Geochemistry
graduation: 29-09-1997
1986 – 1992
Marnix Gymnasium, Rotterdam