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Solubility of Manganese Sulfate Monohydrate in the Presence of Trace Quantities of Magnesium Sulfate Heptahydrate in Water.

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Dev. Chem. Eng. Mineral Process., 11( 5 4 , pp. 423-435.2003.
Solubility of Manganese Sulfate
Monohydrate in the Presence of Trace
Quantities of Magnesium Sulfate
Heptahydrate in Water
Kouji Maedal*, Eric Sukma Chandra', Ha-Ming
Ang3, Moses TsdC3,Wataru Matsuoka' and Keisuke
Fukui'
'Department of Chemical Engineering, Himeji Institute of
Technology, 2167 Shosha, Himeji, Hyogo 671-2201, Japan
*AJ Parker CRCfor Hydrometallurgy, Murdoch University, Perth,
6001, Western Australia
3Department of Chemical Engineering, Curtin University of
Technology, GPO Box U1987,Perth, Western Australia
The solubility of aqueous manganese sulfate monohydrate has been measured
in the presence of magnesium sulfate as the second salt. Over the temperature
range from 273 to 353 K, aqueous manganese sulfate on its own exhibited an
unusual behavior in that its solubility reached a maximum at an intermediate
temperature of 293K. The solubilities of manganese suIphate under and over
the maximum solubility temperature were first predicted individually by the
electrolyte NRTL model. The binary interaction parameters and solubility
parameters were universal parameters that were applied to any salts
consisting of manganese, magnesium and sulfate ions, and they were newly
obtained from the binary solubility data. The predicted solubility curves could
represent the experimental data quantitatively, and adequately described the
effect of co-salt of magnesium. The experimental solubility data were also
satisfactorily correlated with the electrolyte NRTL model.
* Author for correspondence (maeda@mech.eng.himeji-tech.ac.jp).
423
K. Maeda et al.
Introduction
Crystallization of inorganic salts is a common purification technique used in
industry, and crystallization is usually operated under high concentrations of
aqueous electrolyte solution. The aqueous solution system containing
manganese sulfate and magnesium sulfate is one example in electrolytic
manganese dioxide (EMD). Crystallization is a suitable alternative to separate
manganese from magnesium. Manganese sulfate forms the monohydrate while
magnesium sulfate forms the heptahydrateo. However, once these crystals
dissolve in aqueous solution, the salts dissociate in anion and cations. Their
solubilities in aqueous solution are the most fundamental and important
property to quantify before actual crystallization could be studied.
The solution models to express the excess Gibbs free energy such as the
local composition model have been well developed in the last fifty years [l].
In particular, the handling of organic mixtures in the petroleum industry has
promoted the application of the solution models, and these models are used for
the design and operation of distillation, extraction and absorption processes
for the organic chemical industries. In early studies, Pitzer [2] explained the
thermodynamic equation for activity coefficient of electrolyte solutions with
the Debye-Huckel model [3, 4, 51, and his co-workers have applied several
modified models to the systematic electrolyte solutions [6]. Since the 1980s,
Chen and Evans [7] have applied the local composition concept to electrolyte
solutions combined with the Debye-HUckel model. Their model took into
consideration both short- and long-range interactions, and resulted in good
correlation of activity coefficient of relatively high concentrations of salt in
aqueous solution.
Theoretically, the Debye-Hiickel approach (DH) [2] and mean sphere
approximation (MSA) [8, 9, 101 are often used to express the activity
coefficients of electrolyte solute in aqueous solution. The local composition
model is an interesting method to reproduce many electrolyte solution systems
in practice. Chen and his co-workers [ 11, 121 used the NRTL (non-randomtwo-liquid) model [13] for the short range interaction, and they have applied
the electrolyte NRTL model to a variety of solutions including some
electrolytes. Lu and Maurer [14] have also proposed the electrolyte
UNIQUAC (universal quasi chemical) model [15] to be applied to mixed
electrolyte systems. Thomsen et al. [16] used the UNIQUAC model to
describe phase equilibrium including electrolyte systems. Loehe and Donohue
[ 171 recently reviewed the thermodynamic modelling of electrolyte solutions
with respect to an engineering approach.
424
Solubility of Manganese SuIfate Monohydrate in Water
Some examples which present the effect of electrolytes on liquid-liquid
equilibrium, vapor-liquid equilibrium and surface tension have been published
recently [18-211. However, the solubility of solids or crystals of salts is
important for many inorganic industries, as crystallization is a common
separation or purification technique. Kim and Myerson [22] described the
different driving forces of crystallization that were based on common
concentration and chemical potential. To define the chemical potential
difference, the knowledge of the activity coefficient of electrolytes in aqueous
metastable solution is required.
The objectives of this study are to measure the solubility of manganese
sulfate monohydrate in aqueous electrolyte solution, and application of the
electrolyte-NRTL model to predict and correlate our data. The predicted
ternary solubility phase diagram for crystallization of manganese sulfate
monohydrate was proposed.
Experimental Details
(i) Materials
The crystals used in the experiments were basically sulfate crystals, and the
sulfate ions are considered large size ions. The size factor has made most of
the sulfate crystals dissolve readily in water [23]. Hence the solubility of the
MnS04.H20and MgS04.7H20are relatively large.
The manganese (11) sulfate monohydrate used in the experiment has a
purity of 99% by weight. The MnS04.H20 used is an analytical grade,
obtained from the Univar chemical company with molecular weight of 169.01
g/gmol. The MgS04.7H20 used as the second salt is also analytical grade
obtained from the same company. The MgS04.7H20 is described as
colourless, efflorescent crystals with molecular weight of 246.47 g/gmol.
(ii) Procedure
The experimental apparatus in this study is illustrated in Figure 1. The
experiments were designed to investigate the solubilities of manganese sulfate
monohydrate in distilled water with or without magnesium sulfate
heptahydrate as an impurity, in the range of temperature: 275 K, 283 K, 293
K, 303 K, 313 K, 323 K, 333 K, 343 K and 353 K. Four different amounts of
impurity (MgS04.7H20) were examined for this experiment: 0, 3.69, 15 and
30 grams. The solubility data is estimated to have an absolute maximum error
of 0.5 gram of MnS04.H20 (or a standard error of 0.03 mol/kg of water) that
425
K.Maeda et al.
1.
2.
3.
4.
5.
6.
7.
8.
9.
Heaterbath
Temperature
controller
Plexigiass sheet
Flask
Rubberstopper
Stirrerrod
Condenser
Stand
Cooler bath and
temperature
control
I
Figure 1. Experimental Apparatus.
was insignificant experimentally. The experimental apparatus for the solubility
is a three-necked flask mounting a thermometer, condenser and agitator. The
temperature of the solution in the flask is controlled in a thermostatic water
bath. The condenser minimises water evaporating from the flask. Initially, the
mass of the three-neck flask was weighed, then lOOg of distilled water was
added into the three-neck flask and weighed. The total mass and the mass of
distilled water were recorded. The flask with water was then placed into a
water bath that was set to the desired temperature. Before starting the
experiment, the temperature of the solution was maintained at the desired
temperature. A certain known amount of MgS04.7H20 as an impurity was fust
put into the flask. The solution was agitated until all the MgS04.7H20 was
dissolved. A certain amount of MnS04.H20 was then added into the flask and
the solution was agitated. After 10-15 minutes, the remaining MnS04.H20
crystals at the bottom of the flask were investigated visually to judge whether
all MnS04.H20 crystals had dissolved. When all these MnS04.H20 crystals
had dissolved, 0.5 g of MnS04.H20 was continually added to the agitated
solution until no more could dissolve. The total amount of MnS04.H20
crystals was weighed, and the solubility was computed.
Results and Discussion
The experimental data are listed in Table 1. Four different series of the
maximum solubilities of MnS04.H20crystals in aqueous solution, each having
426
Solubility of Manganese Sulfate Monohydrate in Water
a different MgS04.7H20 concentration were measured over a wide
temperature range. It is obvious that the solubility of MnS04.H20decreased as
the MgS04.7H20 content increased in the aqueous solution. The solubility
data seemed to indicate that the presence of MgSO4.7Hz0 did not shift the
temperature of the maximum solubility. If we consider the selective
crystallization of MnS04.H20 from the mixed salts with MgS04.7H20, it is
inevitable that the thermodynamic model to express the effect of co-salts for
each solubility is required. The thermodynamic solution model gives us the
exact solubility having thermodynamic consistency. It will be possible to find
some important solubility points, e.g. the maximum solubility, the eutectic
solubility, and further prediction of solubility with other salts.
Table 1. Ternary solubility data of MnS04.H20 crystals in aqueous solution
containing MgS04.7H20 (ml and m2 are for concentration of MgS04.7HzO and
MnSOI.H20, respectively).
T[K]
275.15
283.15
293.15
303.15
313.15
323.15
333.15
343.15
353.15
275.15
283.15
293.15
303.15
313.15
323.15
ml[mol/kgH20] mz[mol/kgHzO]
ml[mol/kgH20]
mz[mol/kgHzO]
0.00
3.75
4.04
0.15
0.15
3.55
3.97
0.00
0.00
0.00
0.00
0.00
0.00
4.26
4.13
3.88
3.71
3.41
3.11
0.15
0.15
0.15
0.15
4.22
3.97
3.77
3.56
0.15
0.15
3.41
3.02
0.00
2.82
0.15
2.67
0.61
0.61
0.61
0.61
0.61
0.61
3.15
3.77
4.01
3.77
3.58
3.38
1.22
1.22
1.22
1.22
1.22
1.22
2.83
3.08
3.41
3.28
3.08
2.88
333.15 0.61
343.15 0.61
353.15 0.61
3.24
2.86
1.22
1.22
2.69
2.5 1
2.53
1.22
2.36
0.00
42 7
K.Maeda et al.
(i) Solubility model
To study the effect of MgS04.H20 on the solubility of MnS04.7H20 in the
aqueous solution, the thermodynamic equation should be considered for our
data. The electrolyte solutions are more difficult to model than molecular
solutions. Piker [2] described the electrolyte solution model with the DebyeHuckel theory, and succeeded in expressing the activity coefficient of the
dilute electrolyte solution [3]. The activity coefficients of dilute electrolytes in
aqueous solution depend only on ion length. However, the solution treated in
crystallization processes contains highly concentrated salts. Chen et al. [7]
first introduced the local composition model of the NRTL combined with the
Pitzer-Debye-Huckel model to the electrolyte solutions. This model took not
only long-range interactions but also short-range interactions of ions into
account, and provided the possibility of applying it to many types of
electrolyte solutions in the inorganic chemical industries. The modified
electrolyte NRTL-model used here did not consider the electroneutrality
condition in order to use the more universal parameters. The universal
interaction parameters made it possible to predict the activities of many types
of electrolyte solutions.
Alternatively, the molarity unit (m) of our solubility data can be converted
to true mole fraction unit (z) in electrolyte aqueous solution as follows:
x+ =
u .m
9
1000
-++m(u+
+u-)
4
x- =
urn
1000
-+ m(u+ + u - )
Ms
The rigorous solubility product (K,) of salt in the solution is given by:
which is a function of unsymmetrical activity coefficients (y*) and mole
fractions of ions in the solution; v+- is the stoichiometric number of
dissociation equilibrium. We define the solubility product as a function
temperature of the van't Hoff ty-pe as:
B
R In K,,= A + T
428
(3)
Solubility of Manganese Sulfate Monohydrate in Water
where R 18.314 JlmoIJK] is the gas constant, and A and B are parameters of
solubility product for pure salts. These parameters are quite constant in the
solution. Activity coefficients of ions are significant variables to present the
difference of solubility with or without the second salt. The unsymmetrical
activity coefficient is:
The first term on the right hand side of Equation (4) is the contribution of
long range interactions of ions, and is given by the Pizter-Debye-Huckel
equation:
lnY,:@ - -(1000/Ms)1~2A)((2z;
/ p ) l n ( l +PI'/*)
+(z;1;12 - 2
9 /(1+
When the closest approach parameter (p) is constant, only ion strength (I)
is the essential variable in Equation ( 5 ) . The value of the Pitzer-Debye-Huckel
(A,) parameter given in Chen's paper [7, 113 was used in this study. Ion length
is given by:
I=
l N
-c
2 i
z2xi
where, z is ion charge. Only the Pizter-Debye-HBckel term is sufficient to
predict the activity coefficients of dilute electrolytes. The second term on the
right hand side of Equation (4) is the contribution of short-range interactions
of any species in the solution, and is given by the NRTL equation. The
unsymmetrical activity coefficient can be expressed by:
The NRTL equation [ 131 is given by:
429
K.Maeda et al.
G, = exp(-az,),
z, = g , I RT
where ci is the non-randomness parameter, and 0.2 is used in this study as
well as in Chen's report [7, 111; gij is the most important parameter to present
the interaction of i-j combination for any species. These values could express
the effect of co-salt on the solubility, and change the solubility behaviour for
each other.
Table 2. Predictive NRTL parameters g,. [J/moll of the binary pair of i-j
species.
Water
MgZC
Mn2+
so-:
Water
0
64 8
4130
-4560
Mgz+
- 16800
Mn2+
.O
-20500
0
0
85 1
87.8
0
sodz-18500
-3450
2560
0
(ii) Prediction and correlation
The binary NRTL parameters were obtained from the binary activity data book
[24] as shown in Table 2. For the single MnS04.H20solubility in water, there
seems to be two different solubilities under and over the temperature between
290 K and 310 K in Table 1. The solubility of single MgS04.7H20also shows
the maximum values [24]. The solubility data could be separated into two
regions. The pure parameters of solubility product for MgS04.7H20 and
MnS04.H20 considering the electrolyte NRTL solution model were
determined in this study. These parameters for two regions are listed in Table
3. The binary NRTL parameters resulted in the activity coefficients of
electrolytes by Equations (4)-(8), and the pure parameters of salts resulted in
the solubility as a function of temperature by Equations (2) and (3). Here, we
could predict the MnS04.H20 solubility in aqueous solution containing
MgS04.7H20, as shown in Figure 2. Over the critical temperatures, the
solubilities were well predicted by the model, and the effect of MgS04.7H20
content on the solubility was satisfactorily described by the model. Under the
critical temperature, the model could not quantitatively predict the usual
430
Solubility of Manganese Sulfate Monohydrate in Water
I
5.0
I
4.0
c?
3.0
'6
I+
0
2.0
OI"
a m
'5z
50
I
1.0
I
0.1
A 0.6
0 1.2
0.0
270 290 310 330 350 37(
temperature [K]
Figure 2. Predicted solubility of MnS04.H20in the presence of MgS04.7H20
by the electrolyte NRTL model.
solubility, but presented the decrease of solubility with MgS04.7H20content
qualitatively.
The correlated parameters of the NRTL model are shown in Table 4. The
correlation cannot significantly improve this model to represent the solubility
data. It is thought that the model having the predicted NRTL parameters will
be useful for crystallization process of MnS04.H20 in the entire temperature
range.
Table 3. Pure parameters of solubility products for MnS04.H20 and
MgSO4.7HzO.
MnS04.Hr0 (low)
MnS04.Hj0 (high)
MgS04.7H20
(low)
MgSOI.7HtO
A [Jmou
-1 17
-187
-82.8
B [J/mol/K]
4825
26043
-21 1
37312
-6129
(iii) Eutectics between MnS04.H20 and MgS04.7H20solubilities in ternary
system
For selective crystallization of MnS04.H20, it is necessary to find the eutectic
concentrations where both salts may co-crystallize. The concentration of the
feed solution for crystallization should be controlled in the supersaturated
431
K. Maeda et al.
-.
.......
/
K
293 K
323K
333 K
..*
0.0 1.0 2.0 3.0 4.0
molarity of MgS04.7H20
[mol/kgH20]
5.0
Figure 3. Predicted ternary solubility of MnS04.H20and MgS04.
7H20 by the
electrolyte NRTL model.
region for MnS04.H20 crystals relative to MgS04.7H20. The predicted phase
diagram for the system (water + MgS04.7H20 + MnS04.H20) is shown in
Figure 3. The diagonal line represents the aqueous solution containing 50150
mixtures of MgS04.7H20 and MnS04.H20 salts. The eutectic concentration
places the MnSO4.HZO-rich region at relatively low temperatures, and is
shifted closer to the diagonal as temperature increases. However, the eutectic
concentration suddenly changes to the MgSO4.7H20region as the temperature
increases above 323 K. This novel behaviour of the solubility change is caused
by the unusual solubility of MnS04.H20that is distinguished from the normal
solubility at low temperatures. Above 323 K, the concentration field to
crystallize MnS04.H20 expands beyond 50/50 concentration ratios of Mn to
Mg. Preferable operation of selective crystallization for MnS04.H20 crystals
should be at high temperature.
Conclusions
The solubilities of MnS04.H20 crystals were measured in aqueous solution
containing MgS04.7H20 in the temperature range from 283 K to 353 K. The
solubility of MnS04.H20 passed through a maximum and could be separated
into two different solubility curves for the entire temperature range. These
solubilities were predicted and correlated by the electrolyte NRTL model, and
the effect of MgS04.7H20 as the second salt on the solubility of MnS04.H20
was considered. It was found that MgS04.7H20 and MnS04.H20 have a
432
Solubility of Manganese Surfate Monohydrate in Water
capacity of salting-in for the different salts. In predicting the isothermal
ternary phase diagram of the (water + MgS04.7H20+ MnS04.H20) system, it
was evident that a higher temperature will be preferable for the selective
crystallization of MnS04.H20.
Table 4. Correlative NRTL parameters gV [J/mol/ of the binary pair of i-j
species.
Water
0
MgZ+
4 1700
Mn 2+
so42
Water
19000
15800
Mg2’
2300
0
6300
1800
Mn”
5270
3 1400
0
108
SO:
4570
2400
3900
0
Nomenclature
Parameter of van’t Hoff equation, J/mol
Debye-Htickel constant for the osmotic coefficient
4
Parameter
of van’t Hoff equation, J/mol/K
B
G
Paramter of NRTL equation
Ionic strength in mole fraction scale
1,
Solubility
product based on activity
KO
Ms
Solvent molecular weight, g/mol
Gas constant, = 8.3 14 J/mol/K
R
T
Temperature, K
a
Activity
NRTL
interaction parameter for all species; molecules and ions,
gij
J/mol
m
Molality, gmoYkg solvent
True liquid phase mole fraction based on all species; molecules and
X
ions
Greek letters
NRTL non-randomness parameter, = 0.2
a
Y
Activity coefficient
The closest approach parameter of Pitzer-Debye-Huckel equation, =
P
A
14.9
V
Stoichiometric number
433
K.Maeda et al.
NRTL interaction parameter for i-j pair of any species; molecules
and ions
Superscripts
i
Any species in solution
pdh
Pizter-Debye-Huckel equation
OD
infinite dilution
*
unsymmetric convention
+
cation species
anion species
r
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