Hydrometallurgy, 29 ( 1992 ) 173-189
173
Elsevier Science Publishers B.V., Amsterdam
Acidic dissolution of zinc ferrite
F. Elgersma, G.F. Kamst, G.J. Witkamp and G.M. van Rosmalen
Delft University of Technology, Laboratory for Process Equipment, Leeghwaterstraat 44, 2628 CA
Delft, The Netherlands
(Revised version accepted December l 0, 1991 )
ABSTRACT
Elgersma, F., Kamst, G.F., Witkamp, G.J. and Van Rosmalen, G.M., 1992. Acidic dissolution of zinc
ferrite. In: W.C. Cooper and D.B. Dreisinger (Editors), Hydrometallurgy, Theory and Practice.
Proceedings of the Ernest Peters International Symposium. Hydrometallurgy, 29: 173-189.
The dissolution of synthetic and industrial zinc ferrite was studied in HNOs, HCIO4 and H2504
solutions, at temperatures ranging from 75°C to 95°C and in the presence of an excess of Fe 2+, Fe 3+
or Zn 2÷ ions. The rate constant describing this dissolution process was obtained by using a surfacereaction controlled, shrinking core model, which yielded good results.
The dissolution rate of zinc ferrite depends on the square root ofthe hydrogen ion activity in H2SO4
and HCIO4 solutions. The apparent activation energy for the dissolwtion equals 74 + 2 kJ/mol in H2504,
47 + 22 kJ/mol in HCIO4 and 37 + 16 kJ/moi in HNOs. The order of the rate constant was 0.6 in the
Fe 2+ and - 0 . 5 in the Fe s+ concentration. The rate constant for the dissolution of industrial zinc
ferrite was 20-50% lower than for synthetic zinc ferrite. The presence of 85 g/i Zn 2+ in the leach
solution retarded the dissolution rate of zinc ferrite. A dissolution mechanism is proposed which
qualitatively explains the results obtained.
INTRODUCTION
Hydrometallurgical zinc winning plants frequently use ZnS concentrates
which contain up to 10% Fe. During roasting, iron combines with zinc to
form the spinel ZnO'Fe203 (zinc ferrite). In order not to lose zinc incorporated in ZnO'Fe203, acidic leaching (pH 1, 90°C) is required to dissolve it.
The dissolution yields a Zn/Fe solution from which iron has to be removed.
This is mostly done by operating a jarosite, goethite or haematite process.
This paper aims at studying the factors that affect the dissolution rate of
ZnO'Fe203. Such information is useful for optimizing the current leaching
process and is essential for the design of a simultaneous jarosite precipitation/zinc ferrite dissolution process, which may have considerable advanCorrespondence to: F. Elgersma. Delft University of Technology, Laboratory for Process Equipment, Leeghwaterstraat 44, 2628 CA Delft, The Netherlands.
0304-386X/92/$05.00 © 1992 Elsevier Science Publishers B.V. All rights reserved.
174
E ELGERSMAETAL.
tages compared to the current, subsequent, zinc ferrite dissolution-jarosite
precipitation process.
The dependence of the dissolution rate of ZnO.Fe203 o n the composition
of the leach solution was studied because the solution composition influences
the solution potential, which is one of the rate-determining factors in the dissolution process. The solution composition was varied by choosing different
initial concentrations of Fe 2+ and Fe 3+. Furthermore, the dependency on the
hydrogen activity was determined in H 2 S O 4 and HCIO4 leaching and the kinetics were determined, by varying the leach temperature between 75 °C and
95°C in HNO3, HCIOa and n 2 s o 4. The Z n 2+ and the HSO4- concentration
were varied independently in order to study complexation effects in the leach
solution, which indirectly affect the solution potential. Finally, the dissolution rates of synthetic and industrial ZnO. Fe203 were compared.
Although earlier work [ 1-7 ] on the dissolution rate of ZnO.Fe203 has been
published, no systematic study on the relationship between the solution composition and the dissolution rate, as provided in this paper, was available yet.
EXPERIMENTAL
Synthetic ZnO. F e 2 0 3 w a s prepared in a m o u n t s of 5 or 10 g by the following
procedure: In deionized, distilled water, stoichiometric amounts of
ZnSO4.7H20 (72 g/I) and FeSO4.7H20 (139 g/l) were dissolved at 75°C
after addition of l wt% H2SO4. Upon mixing with a solution containing
(NH4)~C~O4 in 10% excess at 90°C, ZnFe2(C204 ).~ was formed. The precipitate was filtered, washed and subsequently dried for 3 h at l l0°C. It was
calcined at 700°C for 20 h to yield ZnO. Fe.,O3. A sample of 0.25 g ZnO.Fe203
was leached for 2 h at 25°C in 50 ml of a solution, containing 20 g NH4C! and
600 ml 25% NHs per liter, for determining the residual ZnO content of the
synthetic product [ 8 ]. Industrial ZnO. Fe203 was isolated from industrial calcine. This was formed in a 9.5 m diameter fluid bed roaster at 910°C using a
ZnS concentrate containing about 52 wt% Zn and 8 wt% Fe. The isolation
procedure included the dissolution of ZnO in 0.1 M H2SO4 at 20°C until no
t~rther increase in the pH was measured. All chemicals used were analytically
pure.
The synthetic and the industrial ZnO.Fe203 were characterized by scanning electron microscopy (SEM) to establish the crystal size and shape and
by X-ray diffraction (XRD) to confirm the crystal structure. Kryptonporosimetry was applied to determine the specific surface area. The particle
size distribution was determined using a Coulter Counter Multisizer. A sample of 0.1 g ZnO'Fe203 was leached for 15-30 min in l0 ml of 36% HCI at
25 °C. This "total" leach provided the total amount of zinc and iron in the
product.
For determination of the dissolution rate, 0.25 g or 0.5 g ZnO'Fe203 was
175
ACIDIC DISSOLUTION OF ZINC FERRITE
leached by either HNO3, HCIO4 o r H 2 S O 4 at temperatures between 75 °C and
95°C in a 250 ml round-bottomed flask, stirred at 610 rpm. Each 3, 5, 10 or
15 min a liquid sample was taken and the zinc and iron contents of the solution were determined by AAS and/or ICP. I f F e 2+ had been added, the flask
was operated under a N2 atmosphere. The flask was immersed in a thermostatically controlled water bath with a terr,perature accuracy of _ 0.2 ° C. The
experiments lasted between 20 rain and 3 h.
H 2 S O 4 w a s used as the leachant because this reflects industrial practice.
HNO3 and HCIO4 w e r e used, in order to avoid complexation between the
acid anion and Zn 2÷ and Fe 3+. Fe(NO3)3"9H20, ZnSO4"TH20,
Zn (NO3)'6H20, FeCI2.4H20 and FeSO4.7tt20 were used as sources for Fe 3+,
Zn 2+ and Fe 2+ ions, respectively. NaHSO4.H/O was added as a source of
excess H S O 4 - ions in an H2SO4 leach solution.
DATA ANALYSIS
The conversion data for each experiment were characterized by a rate constant for dissolution obtained from a surface reaction controlled shrinking
core model given by:
l - ( l - x ) l / 3 = l /3kSo t
(1)
where: x = t h e conversion; k = t h e rate constant (g/m2.min); So=the specific surface area of the solids (m2/g) at t = 0 ; t = t h e time (min).
In deriving eq. ( 1 ) the following assumptions were made:
( 1 ) the particles dissolve isomorphically;
(2) the number of particles remain constant throughout the process;
(3) the dissolution rate per unit surface area is constant, thus a constant
driving force tbr the dissolution process is required.
The conversion x is defined as the ratio of dissolved iron per gram
ZnO. FeaO3 in a sample over dissolved iron per gram ZnO. Fe203 as obtained
0
~is
•
s
i
lo
20
~m
~'o
Fig. !. A representative volume-based panicle size distribution of synthetic zinc ferrite.
176
F. ELGERSMA ET AL.
Fig. 2. A representative SEM photograpii of a synthetic zinc ferrite particle (a) before and (b)
acidic leaching.
after
during the "total leach". The zinc concentration is used for calculating the
conversion in cases where iron ions were added.
ACIDIC DISSOLUTION OF ZINC FERRITE
177
RESULTS
Characterization of zinc ferrite
Synthetic ZnO. Fe203 was prepared in small charges of 5 or I 0 g, in order
to attain a 100% yield. The free ZnO content varied between 2.5 and 5 wt%
and the free Fe203 content between 2 and 7 wt%. Thus, the purity of zinc
ferrite samples was about 90 wt%. A typical volume-based particle size distribution of the synthetically obtained product is given in Fig. 1. No crystalline
phases apart from zinc ferrite were detected. The specific surface area of the
charges varied between 2 and 7 m2/g (accuracy _+2%). Figure 2 shows representative ZnO. Fe203 particles before and after leaching.
In industrial samples, ZnO'Fe203, Zn2SiO4 and PbSO4 were detected as
crystalline phases. Cd and Cu were detected chemically and are assumed to
be present as CdO.Fe203 (0.4 wt%) and CuO'Fe203 ( 1.4 wt%). The resulting purity of ZnO.Fe203 is about 70%. A good particle size analysis could not
be obtained because the industrial samples coagulated during the size analysis. Figure 3a, however, gives a representative view of industrial ZnO. Fe203,
which consists of even smaller particles than partly leached synthetic
ZnO. Fe203 (Fig. 3b).
Applicability of the shrinking core model
The shrinking core model adequately describes the dissolution process as
long as the conversion increases with time. The model yields a straight line
over the whole time range and the slope of this line yields the rate constant k.
In Fig. 4a the measured conversion data are properly fitted by a conversion
curve which is computed with the k value obtained from the shrinking core
model in Fig. 4b. In Fig. 5a the conversion time plot for a typical experiment
carried out in the presence of initially added Fe 3+, Fe 2+ or Zn 2+ is given.
The shrinking core model no longer yields a straight line when the conversion
becomes practically constant (Fig. 5b) and, therefore, the drawn straight line
was used as the best fit for the period of progressing conversion.
Table 1 shows the results of three experiments carried out to check the reproducibility (which was _+2.5% ) using three different charges. The rate constant is calculated assuming the significant internal surface area of ZnO. Fe203
to be fully wetted. This assumption was checked by comparing the results
obtained with this method with results obtained by rotating disc experiments
with a well-defined exposed surface area [ 3 ]. Since the deviation in the rate
constant for two experiments performed under similar conditions was within
5%, this assumption is justified.
Dependence of the rate constant on the temperature
The dependence of the rate constant on the temperature was determined in
0.5 M H2SO4, 1 M HNO3 and 1 M HCIO4 from experiments performed at
178
F. ELGERSMAET AL.
Fig. 3. A representative SEM photograph of industrial zinc ferrite (a) before leaching, compared to (b) synthetic zinc ferrite after leaching.
temperatures between 75 ° and 95°C. The results are listed in Table 2. Plots
of In k versus 1/ T are shown in Fig. 6. The apparent activation energies calculated from the slope of the drawn straight line equal 74_+2 kJ/mol for
H2SO4, 47 _+22 kJ/mol for HCIO4 and 37 +_16 kJ/mol for HNO3.
179
ACIDIC DISSOLUTION OF ZINC FERRITE
O.t, 1-II-X}
0.8 conversion
0
0
0
0
0.6
1/3
O
O
0
0
0.2
OJ,
0.2
(a)
0.0
0.o
0
50
100
150
200
Ib)
II
200
'
' 150
'
s"
time in minutes
time in minutes
Fig. 4. The measured conversion and the conversion predicted by (a) the shrinking core model
and (b) the linear dependency of the model value with time.
conversion
(-1
(a)
1
~ . / ~ ~ , , , ~
0.8
0.6
0.6¸
1-(1-Xl
I/3
(b)
O.l,,
0,2
- model
0,0
,,
o
lo
'
2'o
'
3'o
'
~'o'
s'o
.
0
.
.
.
20
/,0
6o
time in minutes
time in minutes
Fig. 5. The dissolution grade ofzinc ferrite in 2 M HCIO4at 90°C in the presence of (a) 20 g/i
Fe 3+ and (b) the corresponding shrinking core model value.
TABLE 1
Zinc ferrite dissolution experiments I: reproducibility of the results
Temperature
( °C )
Acid
Acid strength
( mol/! )
Added
ions
Rate constant
(g/m'-. rain )
90
90
90
H2SO4
0.5
0.5
0.5
-
0.0128
0.0133
0.0127
H2SO4
H2504
180
F. ELGERSMA ET AL.
TABI,E 2
Zinc ferrite dissolution experiments II: dependence on the temperature
Temperature
(°C)
Acid
Acid strength
(mol/l)
75
80
85
90
95
80
85
90
95
75
80
90
95
H2SO 4
H2SO 4
H2SO 4
H2SO 4
H2SO 4
HNO3
HNOa
HNO3
HNO3
HCIO4
HCIO4
HCIO4
HCIO4
0.5
0.5
0.5
0.5
0.5
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
-3
Added
ions
Rate constant
(g/m2.min)
0.0043
0.0065
0.0091
0.0133
0.0174
0.0011
0.0013
0.0013
0.0020
0.O004O
0.00095
0.00091
0.00123
tn k
-5
H2S04
•
-'7
.q
0.0030
1/T (l/K)
Fig. 6. The Arrhenius plot for the dissolution of zinc ferrite in 0.5 M sulfuric acid, 1 M nitric
acid and 1 M perchloric acid.
Dependence o f the rate constant on the H + and H S 0 4 - ion activity
The dependence of the rate constant on the H + activity was determined for
dissolution in H 2 S O 4 and HCIO4 solutions, from the experiments described
in Table 3. The H + ion activity was calculated using Pitzer's method [ 9 ]. The
required data were obtained from [ l 0 ]. The dissolution of Fe and Zn ions
during the experiment was neglected in the calculation of the activity coefficients. Fig. 7 shows the plots of log k as a function of log [H + ] obtained. For
both acids, an order of 0.50 in the H + activity was found. The activity cal-
181
ACIDIC DISSOLUTION OF ZINC FERRITE
TABLE 3
Zinc ferrite dissolution experiments !11: dependence on H ÷ activity
Temperature
(° C )
Acid
90
90
90
90
90
90
90
90
85
85
85
85
H2SO4
H2SO4
H2SO4
H2SO4
HCIO4
HCIO4
HCIO4
HCIO4
H2SO4
H2SO4
H2SO4
H2SO4
Acid strength
( mol/! )
Added ion
(g/!)
Rate constant
(g/m2. min )
0.25
0.5
1.0
2.0
0.25
0.5
1.0
2.0
0.5
0.5
0.5
0.5
24.3
48.6
72.9
145.8
0.0078
0.0 i 28
0.0172
0.0249
0.0004
0.0006
0.0009
0.0015
0.01 i 6
0.01 ! 3
0.0112
0.0109
HSO~
HSO~
HSO~
HSO~"
TABLE 4
Zinc ferrite dissolution experiments IV: dependence on the Fe 3+ concentration
Temperature
Acid
(° C )
85
85
85
85
90
t)0
90
90
90
H2SO4
H2SO4
H,SO4
1-12SO4
HCIO4
HCIO4
I-ICIO4
HCIO4
HCIO4
Acid strength
( mol/! )
Added ion
Rate constant
(g/i)
(g/m2. min )
0.5
0.5
0.5
0.5
2.0
2.0
2.0
2.0
2.0
3.8
3.9
5.5
6.4
2.0
4.5
10.0
15.0
20.0
Fe 3÷
Fe ~÷
Fe ~÷
Fe a÷
Fe ~+
Fe 3÷
Fe ~+
Fe "~+
Fe 3+
0.0076
0.0067
0.0066
0.0052
0.0015
0.0017
0.0013
0.00095
0.00080
culations showed that the SO4 2- activity was, surprisingly, independent of
the H E S O 4 concentration. Four experiments were carried out where NaHSO4.HzO was added, in concentrations ranging from 0.25 to 1.5 M. The
addition caused an increase in the rate constant of about 20%, compared to a
similar experiment without HSO4- addition.
Dependence o f the rate constant on the Fe 3+ concentration
The dependence of the rate constant on the Fe 3+ concentration was established for H2SO4 and HCIO4 solutions from the experiments listed in Table
4. The results of the experiments are shown in Fig. 8a, where log k versus the
initial Fe 3+ concentration in log(ppm) is plotted. In 0.5 M H2SO4, the order
182
F. ELGERSMA ET AL.
-2.8
log k
-).C
-3.2
{z)
-3.4
-0.8
-1./,,
•
'
'
_~.l,'
'
'
6
'
'
'
log all+
Io9 k
-1.6-1.8
-2.0
-2.2
(b)
.o',e-d,6-o',~,-o,2
o
o~
o~,
Inq i,i *
Fig. 7. The order of the rate constant for zinc fcrrite dissolution in the hydrogen activity (a) in
HCIO., at 90°C and (b) in aaso4 at 90°C.
in the Fe 3+ concentration equals -0.5+_0.2. in 2 M HCIO4, the order is
-0.50_+0.06. For calculating the order in 2 M HCIO4, the results from the
experiment with 2 g/I Fe 3+ were not used.
Dependence o f the rate constant on the Fe z + concentration
The dependence of the rate constant on the Fe 2+ concentration was determined for dissolution in H2SO4 and HCIO4 solutiens (Table 5). Six experiments were carried out with initial Fe 2+ concentrations between 0 and 800
ppm in a 0.5 M H 2 5 0 4 solution. The rate constants did not differ significantly. The addition of 4 g/l Fe 2+ as FeSO4"7H20, however, increased the
rate constant by 80% compared to a similar experiment in the absence of Fe 2+.
ACIDIC DISSOLUTION OF ZINC FERRITE
-2.0
183
tog k (a)
0
0
-3.2
3.0
tog
-1.0
si0
EFe3,
tog k (b)
t
0
-1.5
-25
3.:
'
316
'
,
,
~.0
---4--
,
,
tog
--, ........
T
4.,4.
CFe2+
Fig. 8. The orderof the rateconstant for ZnO.Fe203 dissolution for (a) Fe3+ concentrationand
(b) Fe2+ concentration in 2 M HCIO4at 90°C and 0.5 M H2SO4at 85°C.
Four experiments were carried out to determine the influence of the
FeCI2.4H20 addition on the rate constant in 2 M HCIO4. The results are shown
in Fig. 8b, where log k is plotted against the initial Fe 2+ concentration in log
(ppm). The order of the rate constant in the Fe 2÷ concentration equals
0.62 _+0.25, if the results are fitted with a straight line.
Dependence o f the rate constant on the
Zn 2+
concentration
The rate constant was measured in a 1.8 M H2504 solution in the presence
of 1.3MZn 2+, added either as Zn(NO3)2 or as ZnSO4. Table 6 shows that, in
the presence of NO3-, the addition of 85 g / l Z n 2 + decreases the rate constant
184
F. ELGERSMAETAL.
TABLE 5
Zinc ferrite dissolution experiments V: dependence on the Fe z+ concentration
Temperature
(-C)
Acid
Acid strength
(mol/l)
Added ion
(g/l)
Rate constant
(g/m-'.min)
85
85
85
85
85
85
90
90
90
90
H2504
H2504
H2SO4
H2SO4
0.5
0.5
0.5
0.5
0.5
0.5
2.0
2.0
2.0
2.0
0.073 Fe 2+
0.164 Fe -'+
0.240 Fe -'+
0.421 Fe 2+
0.806 Fe -'+
4.0 Fe -'+
!.8 Fe -'+
3.9 Fe-"+
8.9 Fe-"+
! 7.6 Fe-"+
0.0080
0.0072
0.0066
0.0075
0.0076
0.0166
0.0104
0.0102
0.0148
0.0455
H2SO4
H.,SO4
HCIO4
HCIO4
HCIO4
HCIO4
TABLE 6
Zinc ferrite dissolution experiments Vl: dependence on the Zn-' ÷ concentration
"i'cmpcrature
( ~(')
Acid
Acid strength
(mol/l)
Added ion
(g/l)
Rate constant
(g/m2,min)
00
~)()
~)()
H:SO4
H:SO,~
H:SO,~
1.8
1.8
1,8
85 Zn-' ÷
85 Zn ~'÷
0.0243"
0.0109 !'
0.0140"
" Estimated value,
"Zinc nitrate was used as zinc source,
' Zinc sulfate was used as zinc source,
by 55%. In the presence of an equivalent amount of
SO42- the rate constant
reduces by 42%.
D(['ferences between synthetic and industrial zinc ferrite
In a 4 M HCIO4 solution the rate constant of industrial ZnO.Fe203 was
about 80% of that of synthetic ZnO. Fe203, as follows from Table 7. In a 0.5
M H 2 5 0 4 solution, the ratio of the rate constants was about 0.5. Thus, the
rate constant for the dissolution of industrial ZnO.Fe,O~ is lower than for
synthetic ZnO. Fe203 which is leached under similar conditions. Since the purity of industrial ZnO'Fe203 is only about 70%, part of the specific surface
should be regarded as "non-active", which partly explains the difference.
An attempt to measure the dissolution kinetics of synthetic ZnO.Fe203 in
an industrial hot acid leach solution was not successful. This was because the
increase in zinc or iron concentration, upon leaching, could not be measured
185
ACIDIC DISSOLUTION OF ZINC FERRITE
TABLE 7
Zinc ferrite dissolution experiments VII: comparison with industrial ferrite
Temperature
( °C )
Acid
Acid strength
( mol/! )
Ferrite type
Rate constant
(g/m 2.min )
80
80
90
90
HCIOa
HCIO4
H2SO4
4.0
4°0
0.5
0.5
synthetic
industrial
synthetic
industrial
0.0026
0.0021
0.0133
0.0066
H2SO4
against the high background values of the iron ( 10 g/l) and zinc ( > 40 g/l)
concentration in the solution.
DISCUSSION AND CONCLUSIONS
It is still unknown whether the dissolution kinetics of ZnO. Fe203 in acidic
media can be fully explained by the dissolution mechanism proposed for haematite by Warren and Devuyst [ 11 ], as suggested by Lu and Muir [2 ]. In
this mechanism two options for desorption are distinguished:
Step 1: surface hydroxylation
I--sFe~lO + H20~b-sFe~ll (OH)2
(2 )
Step 2: surface protonation
t--sFelll-OH + H30 + ~-*[--sFelIIOH2 + + H20
(3 )
Step 3a: direct desorption
I--~FeIIIOH+ ~1--~+ FeIIIOH~,.+
(4)
Step 3b: anion adsorption and subsequent desorption
b-~FeUlOH2 + + X- ~l--s FelllOH2 X
(5)
I--sFelllOH2 X ~b--s+ FeUlOH2X 2+
(6)
Dependence o f the rate constant on the H + and HSO4- activity
The order of the rate constant in the H + activity equals 0.50 in both the
HCIO4 and H2SO4solutions. This is in accordance with a theoretical relationship for the acidic dissolution of a metal hydroxide derived by Vermilyea [ 12 ].
This relation may only be applied when H + is the only complexing agent on
the particle surface. Since metal oxides at pH values below that corresponding
to the zero point of charge (ZnO.Fe203, pHzpc= 3.5 [ 13 ] ) are hydroxylated
(step 1 in the model), the same relationship may also be applied for describing the acidic dissolution of metal oxides:
186
F. ELGERSMA ET AL.
( { n _ k T n _ C } ~ '~+z+~/~'~+l+-°~-z-~
r = n +k + \ {--'ffg-~r-;~ }
(7)
where:C=the H + concentration (tool/l); a + and t~- =the electrochemical transfer coefficients (which are usually equal to 0.5 [ 14 ] ); n = the number ofions/cme; z=the charge number of H +; kr ÷ and kr- =constants; r=the
dissolution rate (mol/l.cm 2 ).
With z+ = 1 for H ÷ and z_ = - 1 for the transferred electron, it becomes
clear that the dissolution rate depends upon the square root of the H +
concentration.
Additions of 0.25-1.5 M HSO4- to 0.5 M H2SO4 solutions all increase the
rate constant for dissolution by 20%; probably as a consequence of the excess
of HSO4- already present.
The dependence of the rate constant on the potential of the solution
The presence of F e 2 + and Fe 3+ directly influences the potential of the solution and, thereby, the dissolvtion rate of ZnO. Fe2Oa, which is determined
by the potential difference between the surface of the particles and the bulk
of the solution. The reductive dissolution, here in the presence of Fe 2+ in
solution, has been described by Lieser [ 15] for the dissolution of anhydrous
Fe2(SO4)3. The rate-determining step was the reduction of a Fe a+ ion in the
surface layer by a hydrogen atom, which, it is claimed, is formed in the double
layer due to the oxidation of a dissolved Fe 2+ ion. If Fe 2+ is present in the
leach solution, in the absence of other reductants it becomes the ion determining the solution potential. The lower solution potential, due to the presence of Fe 2÷, increases the direct desorption rate (step 3a in the model),
because Fe 3+ desorbs more easily from the surface layer, after being indirectly reduced by dissolved Fe 2+. Upon entering the solution the valency of
the desorbed species becomes determined by the solution potential.
Nil and Hisamatsu [ 7 ] presented the dependence of the rate constan~t for
ZnO. Fe203 dissolution, on the addition of 0-1 g/l Fe 2+ as FeSO4.7H20 to a
9 wt% H2SO4 solution at 50°C. The order of the rate constant in the Fe 2+
concentration was 0.5. Our results, however, show that, at 85°C in a 0.5 M
H2SO4solution, the addition of0-1 g/l Fe 2÷ did not change the rate constant,
as is shown in Table 5. This is probably due to the fact that, at higher temperatures, the dissolution proceeds so much faster that a small decrease in solution potential, due to a small Fe 2+ addition, is no longer significant.
In 2 M HCIO4 at 90°C using FeCl2.4H20 the order of the dependency on
the Fe 2+ concentration equals 0.62 _+0.25. It is, however, questionable whether
a linear fit is applicable. In the presence of Cl- and Fe 2+ the direct desorption
rate is increased by the presence of Fe 2+, but indirect desorption via CI- ad-
ACIDIC DISSOLUTION OF ZINC FERRITE
187
sorption (step 3b in the model) may also occur. Since FeCI2.4H20 was used
as the Fe 2+ source, both the Cl- concentration and the Fe 2+ concentration
increased. This probably caused the deviation from the straight line. Further
research on the separate influence of Fe 2+ is, therefore, still to be done.
No experiments have been reported in the literature in which ZnO'Fe203
was dissolved in the presence of Fe 3+. The above results show that the order
of the rate constant in the Fe 3+ concentration equals -0.5. No activities for
the Fe-'+ or Fe 3+ addition experiments were calculated because no appropriate data for Pitzer's models were available. Our experimental results, are in
reasonable agreemerit with a theoretically derived equation, presented by
Gorichev and Kipriyanov [ 16 ], for the dependency of the rate constant for
the dissolution of magnetite (Fe304) on the Fe 2+/Fe 3+ concentration:
0 5
0.5
--0 5
k = k o a f-i + a Fe2 + a Fe3";
(8)
The model of Gorichev and Kipriyanov [ 16 ] contains the same rate-determining step as the model by Lieser [ 15 ] and additionally includes the influence of the presence of Fe 3+ for determining the solution potential.
The solution potential is indirectly influenced by the presence of Zn2+,
SO 4- or CI-. Zn 2 + ions compete with Fe 3+ ions for complex formation with
SO4::-. The decelerating influence of Zn 2+ can be understood by assuming
that the free Fe 3+ concentration increases, and thereby the solution potential,
causing a reduction of the dissolution rate.
Both SO42- and Cl- play a similar role. By complexation of Fe 3+ with either
Cl- or SO42-, the free Fe 3+ concentration decreases, which decreases the solution potential. The complexation stimulates desorption via anion adsorption (step 3b). This explains the considerably lower leaching rate in HNO3
and HCIO4 solutions, when no complexation is expected, compared to the
rate in H2SO4 and HCi solutions. Surprisingly, the apparent activation energies of the dissolution in H2SO4 (74-+ 2 kJ/mol) and HCI (83 kJ/mol [ 4 ] )
are higher than the corresponding values in HNO3 (37_ 16 kJ/mol) and
HCIO4 (47_+22 kJ/mol), which fact points at a less favorable reaction path.
For this apparent discrepancy no explanation is currently available.
The reason for the large degree of uncertainty in the values for the activation energy in HNO3 and HCIO4 is unknown. Thus, further research is required in order to obtain these data with a higher accuracy. Despite the limited accuracy, the experiments clearly show the significant difference in
activation energy for the dissolution of zinc ferrite in H2SO4, compared to the
values in HNO3 and HCIO4.
CONSEQUENCES FOR INDUSTRIAL PRACTICE
The above results indicate that high Fe 3+ concentrations decelerate the dissolution of ZnO.Fe:O3~ However, by dissolving ZnO.Fe203, Fe 3+ is released.
188
F. ELGERSMAETAL.
It is, therefore, desirable to introduce a sink for Fe 3+. A suitable sink for Fe 3+
is jarosite (NHaFe3 (SO4)2 ( O H ) 6 . By simultaneously dissolving Z n O ' F e 2 0 3
and precipitating jarosite, the dissolution of zinc ferrite will not be retarded
by the presence of Fe 3+. Meanwhile, the precipitation of jarosite is taking
place under controlled conditions because the supersaturation for the precipitation is determined by the dissolution of Fe 3+ from ZnO'Fe203. This is beneficial for the produced crystal size and shape and, consequently, improves
the filterability of the solids. A better filterability reduces the zinc losses in
the jarosite residue. Experiments where this simultaneous conversion is studied will be reported in a subsequent paper.
ACKNOWLEDGEM ENTS
The Dutch Ministry of Housing, Physical Planning and the Environment is
acknowledged for sponsoring this reseach. F. van der Ham is thanked for his
contribution to this study.
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