Un estudio del Oxalato férrico para la disolución de Au en soluciones de Tiosulfato
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Transcript of Un estudio del Oxalato férrico para la disolución de Au en soluciones de Tiosulfato
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decrease due to the homogenous reaction between
copper(II) and thiosulfate. In addition, the generation
of polythionates from the oxidation of thiosulfate was
found to have a negative impact on both golddissolution (Chu et al., 2003; Jeffrey et al., 2003),
and the recovery of gold using ion exchange resins
(Nicol and OMalley, 2002). The problems associated
with this system have motivated the study of alternative
oxidants to dissolve gold in thiosulfate solutions. To
date, very few studies have been attempted on the use
of other oxidants apart from copper and ammonia in
thiosulfate solutions.
Thermodynamically the dissolution of gold in a 1
M thiosulfate solution requires an oxidant with a
reduction potential greater than 0.15 V, which is thestandard potential for the gold oxidation reaction
(Nicol et al., 1987):
Au 2S2O23 YAuS2O3
32 e
1
In this work, using 0.1 M thiosulfate solution, it
will be shown that gold dissolution in solutions
containing no added gold(I) requires a potential
greater than 0.1 V.
Iron(III) was selected as an oxidant because it is
inexpensive and it has high enough potential to
effectively oxidize gold. It is well known that ligands
can be used to stabilize ferric ions in solution, and
hence minimize its reaction with reductants such as
thiosulfate. In this study oxalate was chosen as the
ligand for the following reasons:
a. Oxalate forms a strong complex with iron(III): logb
[Fe(C2O4)+]=7.58; log b [Fe(C2O4)2
]=13.81; and
log b [Fe(C2O4)33]=18.6 (Smith and Martell,
1976).
b. Sodium oxalate is produced as a waste product
from alumina refineries.c. Ferric oxalate undergoes photocatalytic decompo-
sition to form ferric/ferrous oxide, CO2 and CO32
(Dudeney and Tarasova, 1998). Thus it could be
discharged to a tailings dam without causing a
serious threat to the environment.
The speciation of ferric in oxalate containing
solutions can be estimated using the equilibrium
equations and stability constants for the various
complexes coupled with the mass balance for ferric
and oxalate. Fig. 1 shows the speciation for 5 mM total
ferric as a function of total oxalate concentration. When
the total oxalate concentration is 15 mM (mole ratio
Ox/Fe=3), the predominant species present isFe(C2O4)3
3. As the mole ratio of Ox/Fe increases
above 3, the free oxalate concentration increases with
little change to the ferric speciation. For Ox/Fe=2,
Fe(C2O4)2 will be the major species, but there is some
Fe(C2O4)33 and Fe(C2O4)
+, whilst for Ox/Fe=2.5, the
solution will contain almost equal amounts of
Fe(C2O4)2 and Fe(C2O4)3
3.
The standard reduction potential for the ferric
oxalate complexes can be calculated using the
stability constants quoted above, and the stability
constant for the ferrous oxalate complex, log b[Fe(C2O4)2
2]=5.15 (Smith and Martell, 1976). For
the reduction of Fe(C2O4)2 as shown in Eq. (2), given
that E8 (Fe3+/Fe2+)=0.771 V (Lide, 1995), the E8 is
calculated to be 0.258 V, and thus we can expect
Fe(C2O4)2 to oxidize gold to gold thiosulfate. For the
more stable Fe(C2O4)33 complex, the reduction can
be represented by Eq. (3), for which the E8 is
calculated to be 0.025 V, which is substantiallymore negative than standard reduction potential of the
Fe(C2O4)2 complex.
Fe(C2O4)2
+e
YFe(C2O4)22
(2)
Fe(C2O4)33+eYFe(C2O4)2
2+C2O42 (3)
It must be remembered that the standard potentials
are calculated assuming unit activity of each species.
5 10 150
1
2
3
4
5Fe(C
2O
4)+
Fe(C2O
4)-2
Fe(C2O
4)3-3
Concentratio
n/mM
Total Oxalate / mM
Fe3+
C2O
42-
Fig. 1. Calculated speciation diagram for the ferric oxalate
complexes. [Total Fe(III)=5 mM].
I. Chandra, M.I. Jeffrey / Hydrometallurgy 77 (2005) 191201192
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As will be shown later, the Fe(C2O4)33 complex is
quite stable in thiosulfate solutions, and hence it
would be expected that for an Ox/Fe mole ratio of 3,
the ferrous and free oxalate concentrations will be
significantly lower than the ferric concentration. The
calculated potentials for the reduction of Fe(C2O4)33
at different activities of each species involved in the
half cell reaction are shown in Table 1. Even when the
solutions contain oxalate to ferric ratios higher than
3:1, the potential is positive enough that one expects
ferric oxalate to be a strong enough oxidant to
dissolve gold.
2. Experimental methods
All experiments were carried out using solutions
prepared from analytical reagents and Millipore water.
The ferric oxalate complex was prepared by mixing
ferric nitrate and sodium oxalate. Unless specified, all
experiments were carried out at 20 8C, at a rotation
rate of 300 rpm, and at the natural pH of the solution,which was typically pH 4.56. All electrochemical
experiments were carried out using a Model 362
potensiostat (EG and G PAR). All potential measure-
ments were made with respect to a saturated calomel
reference and corrected to the standard hydrogen
electrode (SHE). A platinum wire was used as the
counter electrode, and the potential sweep rate was 1
mV s1. Mass changes were measured using the
rotating electrochemical quartz crystals microbalance
which is described elsewhere (Jeffrey et al., 2000).
Prior to each gold leaching experiment, the gold was
electroplated onto the quartz electrode at 25A m2
from a solution containing 0.02M potassium dicyano-
aurate, 0.23 M potassium cyanide, and 0.086 Mpotassium carbonate. The thiosulfate concentration in
the solution was measured using the flow injection
system developed recently (Dai et al., 2005), and by
the traditional iodine titration. Before using the iodine
titration, the ferric and ferrous were precipitated from
solution using sodium hydroxide, and then the
solution was filtered.
3. Results and discussions
3.1. Gold dissolution kinetics
The kinetics of gold dissolution in the thiosulfate
solutions were measured using the Rotating Electro-
chemical Quartz Crystal Microbalance (REQCM).
This instrument has already been used extensively to
measure the kinetics of gold dissolution in both
cyanide and thiosulfate solutions. Fig. 2 shows the
type of information which can be obtained using the
REQCM. The top half of this figure shows the mass
which is measured using the REQCM. The solid line
shows the dissolution behaviour of an air saturated
Table 1
Calculated potentials for the reduction of the ferric oxalate
complexes at various activities
Reductionof complex
Activities of species Calculatedpotential
Fe(C2O4)2 Fe(C2O4)2
=1 M 0.258 V
Fe(C2O4)22=1 M
Fe(C2O4)33 Fe(C2O4)3
3=1 M 0.025 VFe(C2O4)2
2=1 M
C2O42=1 M
Fe(C2O4)33 Fe(C2O4)3
3=4.9 mM 0.312 V
Fe(C2O4)22=0.1 mM
C2O42=0.1 mM
Fe(C2O4)33 Fe(C2O4)3
3=4.9 mM 0.211 V
Fe(C2O4)22=0.1 mM
C2O42=5.1 mM
-30
-25
-20
-15
-10
-5
0
E/mV
addition of Fe(III)
m/
g
no TUwith TU
0 500 1000 1500 2000 2500
100
200
300400
500
600
t / s
Fig. 2. Change in gold mass versus time in an air saturated
thiosulfate solution in the presence and absence of thiourea and
effect of Fe(III) nitrate addition. Figure also shows the mixed
potential measured for each system. [Conditions: 0.1 M
(NH4)2S2O3, 5 mM thiourea, 5 mM Fe(NO3)3, 20 8C].
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thiosulfate solution prior to and after the injection of
Fe(III) into solution. It can be seen that the mass of the
electrode decreases as the metal is dissolved and the
decrease in electrode mass is related to the golddissolution rate, r, through Eq. (4),
rDm
AMDt4
where r is the oxidation rate (mol m2 s1), A is the
electrode surface area (m2), M is the relative atomic
mass of the metal (g mol1), m is the mass of the
electrode (g) and t is the time elapsed (s).
Fig. 2 initially shows that the gold dissolution rate
in air saturated thiosulfate solutions is very slow in the
absence of Fe(III) and thiourea catalyst; and a rate of3.7107 mol m2 s1 can be calculated from theslope of the line. Thus, under these conditions, it is
unlikely that very high extraction of gold from ores
can be obtained using an air saturated thiosulfate
system. However the figure also shows a dramatic
increase in mixed potential and gold dissolution rate
was obtained when 5 mM iron(III) nitrate was injected
into the thiosulfate solution. Upon the addition of the
ferric ions into the thiosulfate solution, a strong violet
colour was produced. This colour is due to the
formation of intermediate complex between ferric
ions and thiosulfate, FeS2O3+ (Page, 1954, 1960).
However, as can be seen from Fig. 2, the mixed
potential and gold dissolution rate both decay rapidly,
indicating that the Fe(III) is rapidly reduced to Fe(II)
by thiosulfate. This reaction is known to occur
according to the following equation (Tykodi, 1990):
2Fe3 2S2O23 XFeS2O
3Y2Fe
2 S4O26 5
This is supported by the observation that the
violet colour fades as the reaction proceeds. After
s te ad y s ta te i s r ea ch ed t he m ix ed p ot en ti alapproaches the value that was obtained prior to
iron(III) addition.
It has been previously reported that thiourea
greatly enhances the gold oxidation half reaction in
thiosulfate solutions (Chandra and Jeffrey, 2004).
Fig. 2 (dashed lines) also shows the dissolution of
pure gold in solution containing thiosulfate and
thiourea (TU). Initially for the air saturated solution
in the presence of 5 mM thiourea, the dissolution of
gold is approximately 3 times faster than in the
absence of thiourea. The gold dissolution rate is
calculated to be 1.1106 mol m2 s1, althoughthis value is substantially lower than the rates
measured in the copperammoniathiosulfate orcyanide systems (Jeffrey et al., 2001). A significant
increase in the gold dissolution rate and mixed
potential was obtained upon the addition of ferric
ions. However the mixed potential and gold dis-
solution rate both decay in a similar manner to that
in the absence of thiourea due to the homogeneous
reduction of Fe(III). In this instance the Fe(III) may
be reacting with thiosulfate and/or thiourea. In
summary, it can be concluded from the data shown
in Fig. 2 that the presence of a ligand is required to
complex with the iron(III) in order to increase itsstability in thiosulfate solution.
Experiments were conducted using oxalate as a
ligand for the ferric ion. The data shown in Fig. 3
were obtained in a similar manner to Fig. 2, with the
exception that the ferric solution injected contained
oxalate at a mole ratio of 2.5. Upon the addition of the
ferric oxalate complex, the gold dissolution rate and
mixed potential increased, and did not decay rapidly
as was observed in Fig. 2. In the absence of thiourea, a
steady gold dissolution rate of 5.5106 mol m2 s1
was obtained. However, in the presence of thiourea
the gold dissolution rate was significantly higher
(24106 mol m2 s1).
-60
-40
-20
0
Ferric oxalate added
E/mV
m/
g
no TU5mM TU
0 250 500 750 1000 1250 150060
120
180
240
300
360
t / s
Fig. 3. Dissolution of gold and mixed potential in air saturated
thiosulfate with thiourea and effect of Fe(III) oxalate addition.
(Conditions as Fig. 2 with 12.5 mM oxalate added).
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3.2. Effect of oxalate concentration
Oxalate is important in the dissolution of gold in
thiosulfate using iron(III) as an oxidant as it com-
plexes the ferric ions and hence reduces its reactivity
with thiosulfate. Therefore changing the oxalate
concentration will affect the stability of the iron(III)
complex, and thus alter the potential of the iron(II)
iron(III) redox couple. The influence of oxalate on the
dissolution of gold was investigated by varying the
oxalate concentration in the range of 10 to 17.5 mM
for a fixed Fe(III) concentration of 5 mM. This will
result in a range of distribution of Fe(III) complexes
from mainly Fe(C2O4)2 to mainly Fe(C2O4)3
3 with
free oxalate. Fig. 4 shows the effect of oxalate
concentration on the gold dissolution rates. When
the Ox/Fe mole ratio was 2 or 2.5, the gold dissolution
was initially very fast. However, the gold dissolution
rate decreased during the first 5 h. However, between
5 and 24 h, there is li tt le change in the gold
dissolution rate. It is believed that the reduction in
dissolution rates with time is a result of the slowreduction of iron(III) to iron(II) by thiosulfate and/or
thiourea. This is supported by the fact that after 6 h,
the solution potential (EH) also dropped from 0.36 to
0.23 V for Ox/Fe=2, and from 0.33 to 0.22 V for Ox/
Fe=2.5. This aspect is discussed in more detail below.
The results from the experiments conducted at
higher oxalate concentrations are also shown in Fig. 4,
from which it can be seen that the gold dissolution
rate was relatively constant over 24 h for Ox/Fe=3 and
3.5. The initial dissolution rate obtained was lower
than the initial value obtained at lower oxalate
concentrations. However, Fig. 4 shows that after 24
h, the gold dissolution rates obtained for 10, 12.5, and
15 mM oxalate are essentially the same. For theremainder of the present paper, a Ox/Fe ratio of 2.5
was chosen as a compromise between solution
stability and fast gold dissolution kinetics. As will
be shown in Section 3.8, under these conditions, the
gold dissolution rate is higher than that obtained for
the copperammonia system under similar conditions,
and comparable to the rates obtained for dissolution of
gold/silver alloys in cyanide solutions.
3.3. Electrochemistry of ferric oxalate complexes
3.3.1. Voltammetry studies
In order to gain further understanding about the
ferric oxalate complexes, electrochemical experiments
were performed using a platinum electrode in sol-
utions containing ferric oxalate, thiosulfate and
thiourea. Since platinum is inert in thiosulfate media,
its use enables the study of the cathodic half reaction
of gold dissolution, i.e. ferric oxalate reduction. The
first set of experiments were performed using freshly
prepared solutions, and the polarisation curves
obtained for ferric oxalate reduction at different ferric
to oxalate mole ratios are shown in Fig. 5. At an Ox/
Fe ratio of 3, the reduction commences at 260 mV,
and approaches a limiting current of 8.3 A m2 at100 mV. The shape of the polarisation is typical for
0 5 10 15 20 250.0
0.5
1.0
1.5
2.0
2.5
3.0
105r/molm-2
s-1
t / hr
10.0 mM oxalate12.5 mM oxalate15.0 mM oxalate17.5 mM oxalate
Fig. 4. Effect of oxalate concentration on the gold dissolution
kinetics. (Conditions: as Fig. 2 with oxalate added).
-200 -100 0 100 200 300 400-10
-8
-6
-4
-2
0
i/Am-2
E / mV
10mM12.5mM15mM
Fig. 5. Polarisation curves for the reduction of fresh solutio ns of
ferric oxalate with mole ratio Ox/Fe=2, 2.5, 3. (Conditions: as Fig. 2
with oxalate added).
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the reduction of other ferric complexes, such as
ferricyanide. The limiting current for ferric reduction
for Ox/Fe=2.5 is similar to Ox/Fe=3, although for Ox/
Fe=2, the limiting current is slightly lower. Inaddition, at the lower Ox/Fe ratios, the shape of the
polarisation curves are substantially different. There
appears to be two reduction waves, and the limiting
current for the first wave is higher at the lower Ox/Fe
ratio. These results are not surprising when it is
considered that at Ox/Feb3, there are the two
dominant ferric oxalate complexes, Fe(C2O4)2 and
Fe(C2O4)33. The first wave in the voltammogram
corresponds to Fe(C2O4)2 reduction, whilst the 2nd
wave corresponds to reduction of the more stable
Fe(C2O4)33
.It is important to note that when Ox/Fe=2, the
majority of the ferric iron exists as Fe(C2O4)2.
However, the voltammogram in Fig. 5 shows that
the reduction wave for Fe(C2O4)33 is larger than the
reduction wave for Fe(C2O4)2. Such a result is
consistent with a change in speciation equilibrium at
the interface during the reduction reaction. The
reduction of Fe(C2O4)2 shown in Eq. (2) forms the
ferrous complex Fe(C2O4)22. It is proposed that
oxalate exchange from the ferrous Fe(C2O4)22 to
the ferric Fe(C2O4)2 occurs at the interface according
to Eq. (6). The equilibrium constant for this reaction
can be estimated from the stability constants quoted
earlier, along with the stability constant for FeC2O4,
log b=3.05, and the value calculated is K=490. Thus
thermodynamically this reaction lies towards the right
hand side. This reaction is presumably responsible for
the generation of the Fe(C2O4)33 at the interface, and
hence the occurrence of the large Fe(C2O4)33
reduction wave in the voltammogram for Ox/Fe=2.
Thermodynamically it is also possible for ligand
exchange from ferrous Fe(C2O4) to Fe(C2O4)2, as
shown in Eq. (7); the equilibrium constant for thisreaction is calculated to be 55.
FeC2O422 FeC2O4
2YFeC2O4 FeC2O4
33
6
FeC2O4 FeC2O42YFe
2 FeC2O433 7
The occurrence of Eqs. (6) and (7) can be verified
with some simple experiments utilizing ferrous
oxalate. A solution was prepared containing 2.5 mM
ferrous, 5 mM oxalate, 100 mM thiosulfate and 5 mM
thiourea and allowed to stand. After 7 min, a yellow
precipitate formed, which is the well known ferrous
oxalate dihydrate. A second solution was preparedwith the same composition, with the exception that
after mixing, 5 mM ferric complexed to 10 mM
oxalate was added. In this instance, the precipitate was
not observed even after standing for 5 h. A third
solution was prepared containing 2.5 mM ferrous, 100
mM thiosulfate and 5 mM thiourea, and 5 mM ferric
complexed to 15 mM oxalate was added after mixing.
The polarisation curve for the reduction of ferric from
the second and third solution were then obtained and
found to be identical. These results indicate that the
speciation of the second and third solutions are thesame, which is consistent with the Fe(C2O4)2
in the
second solution taking up the oxalate associated with
the ferrous oxalate as shown in Eqs. (6) and (7).
3.3.2. Electrochemical reduction of Fe(III) oxalate
aged in thiosulfate solution
Electrochemistry can also be used to understand
why the gold dissolution rate shown in Fig. 4
decreases with time. The polarisation curves for the
reduction of ferric oxalate on platinum in solutions
containing ferric, oxalate, thiourea and thiosulfate
which had been aged for 5 h are shown in Fig. 6. At
Ox/Fe=3, the polarisation curve is almost identical to
that obtained for the fresh solution shown in Fig. 5.
Therefore it can be concluded that the Fe(C2O4)33
-200 -100 0 100 200 300-10
-8
-6
-4
-2
0
i/Am
-2
E / mV
10mM12.5mM15mM
Fig. 6. Polarisation curves for the reduction of fresh solutions of
ferric oxalate aged for 5 h with mole ratio Ox/Fe=2, 2.5, 3.
(Conditions: as Fig. 2 with oxalate added).
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complex, which is dominant at Ox/Fe=3, is unreactive
towards thiosulfate and thiourea. However when the
Ox/Fe b3, the limiting current for ferric reduction is
lower for the aged solution, indicating that some ofthe ferric has been homogeneously reduced while the
solution was allowed to age. Based on the limiting
currents it can be calculated that for Ox/Fe=2.5, the
ferric concentration had decreased to 4.2 mM, whilst
for Ox/Fe=2, it had decreased to 3.2 mM. Thus b2
mM thiosulfate or thiourea is oxidized.
The other interesting aspect of the data shown in
Fig. 6 is that apart from the limiting current, the
polarisation curves are similar for each of the Ox/Fe
ratios. In particular, the two reduction waves, which
were prominent in Fig. 5, are much less apparent forthe aged solutions. Such a result indicates that the
Fe(C2O4)33 is a major species present in the aged
solutions, even when the initial Ox/Fe ratio was 2.
The most likely mechanism by which Fe(C2O4)33 is
formed is that shown in Eqs. (6) and (7). Therefore it
is not surprising that the gold dissolution rate in aged
solutions is very similar when the Ox/Fe ratio is
between 2 and 3.
3.4. Re-oxidation of Fe(II) oxalate by oxygen
Although the data shown above have indicated that
Fe(C2O4)33 is not readily reduced by thiosulfate,
when leaching a highly reductive gold ore, it is likely
that the ferric will be reduced to ferrous. Thus it is
worth considering the re-oxidation of ferrous to ferric
using oxygen. It is generally accepted that the reaction
between ferrous and oxygen is slow. The oxalate
system has the extra complication that ferrous oxalate
dihydrate can precipitate given the correct conditions.
When solutions containing 1 mM ferrous and 10 mM
oxalate were prepared, it was found that precipitation
did not occur, and the solution potential increasedwhen exposed to oxygen. The solution EH of the
ferrous oxalate was initially 50 mV, increasing to 160
mV after 15 min, and to 230 mV after 30 min, which
was close to the steady state value of 238 mV.
Therefore under appropriate conditions, ferrous oxa-
late can be oxidized to ferric oxalate, but whether this
can occur during the leaching of highly reductive ores
will depend on a number of factors and is beyond the
scope of the present paper. However it should be
pointed out that if the ore is not highly reductive, then
the regeneration of ferric oxalate from ferrous may not
be necessary.
3.5. Thiosulfate consumption
The main reason the ferric oxalate system was
investigated was the high thiosulfate consumption
which is obtained when using the traditional cupric
ammine system. This is due to the redox cycle
between Cu(I) and Cu(II), with both the copper(II)
reduction and copper(I) reactions being rapid. It was
envisioned that due to the slower reactivity of ferric
oxalate, and low reactivity of ferrous oxalate towards
oxygen, there would not be a redox cycle involving
the oxidation of thiosulfate for the ferric oxalatesystem. This was confirmed by measuring the
thiosulfate consumption when bubbling pure oxygen
through a solution containing Ox/Fe ratios between 2
and 3. In each instance, after 5 h, the change in
thiosulfate concentration was within the detection
limit (2%) of the experimental method. Therefore it
can be concluded that thiosulfate consumption is
negligible for the ferric oxalate system.
3.6. Effect of rotation rate on gold dissolution
Fig. 7 shows a Levich plot of gold dissolution rate
as a function of x1/2, from which it is possible to
establish whether the reaction is diffusion or chemi-
cally controlled. Initially when the Ox/Fe ratio was
2.5, the gold dissolution rate shows some dependence
0 60
1
2
3
105r/mol
m-2s
-1
1/2/ s1/2
Fe:ox = 1:2.5
Fe:ox = 1:2.5 after 5 hrs
Fe:ox = 1:3
2 4 8
Fig. 7. Levich plots for solutions containing mole ratio Ox/Fe=2.5
and 3. (Conditions: as Fig. 2 with oxalate added).
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on the rotation rate. Such a result implies that the
reaction under these conditions is at least mixed
controlled. Fig. 7 also shows the data for the same
solution after 5 h ageing and for the Ox/Fe ratio of 3,
for which the gold dissolution rate is largely
unaffected by rotation rate. Such a result implies that
under these conditions, gold dissolution is mainly
chemically controlled. The shift from mixed control at
Ox/Fe=2.5 to chemical control at Ox/Fe=3 can be
readily explained by overlaying the polarisation
curves for the reduction of ferric oxalate with the
polarisation curve for the oxidation of gold in
thiosulfate solutions containing thiourea. This Evans
diagram is shown in Fig. 8. For the solution with Ox/
Fe=3, and aged solutions with Ox/Fe=2.5, the
intersection of the anodic and cathodic polarisation
curves clearly occur in the chemical control region.
However for fresh solutions with Ox/Fe=2.5, the
polarisation curves intersect in the region in which the
first reduction wave of Fe(C2O4)2
is approaching adiffusion limiting value prior to the second reduction
wave of Fe(C2O4)33. Thus the reduction of ferric will
be affected by the mass transfer of Fe(C2O4)2, and
hence the gold dissolution rate shows a dependence
on rotation rate.
Another interesting aspect of the data shown in
Fig. 7 is the apparent non-zero intercept. This is
typical for chemically controlled reactions, as they do
not show any dependence on rotation rate. However
when the Ox/Fe=2.5, the non-zero intercept is also
evident. At very low rotation rates, the rate of the
diffusion limiting process giving rise to mixed control,
viz. Fe(C2O4)2 reduction, becomes very small.
However at low rotation rates, gold dissolution canstill occur via the reduction of Fe(C2O4)3
3 which is a
chemically controlled process.
3.7. Effect of ferric and thiosulfate concentrations
Fig. 9 shows the effect of iron concentration on
the initial gold dissolution rate in ferric oxalate
thiosulfate solutions. Also shown are the gold
dissolution rates after 5 h and 24 h from the time
leach solutions were prepared. Increasing iron(III)
concentration from 1 mM to 10 mM results in alinear increase in the initial gold dissolution rates,
indicating a 1st order dependence of reaction rate on
Fe(III) concentration. However, after 5 h the gold
dissolution rate is lower for each of the experiments,
and there is not such a marked influenced of iron
concentration on the gold dissolution rate. Allowing
the solution to stand from 5 to 24 does not result in a
significant decrease in the gold dissolution rate. The
non-zero intercept of the data may partially attributed
to the fact that in the absence of ferric oxalate, gold
leaches at a low rate in oxygenated solutions
containing thiosulfate and thiourea.
Fig. 10 shows a plot of gold dissolution rates as a
function of thiosulfate concentration. The thiosulfate
concentration was varied in the range from 0.05 M to
0.4 M, while the Fe(III), oxalate and thiourea
0 100 200 300 4000
2
4
6
8
i/Am-2
E / mV
Ox:Fe=2.5 fresh reduction
Ox:Fe=3 fresh reduction
Ox:Fe=2.5 5 hrs reduction
gold oxidation
Fig. 8. Evans diagram for the reduction of ferric oxalate for mole
ratio Ox/Fe=2.5 and 3, and the oxidation of gold. (Conditions: as
Fig. 6).
0 100
1
2
3
4
5
105r/molm-2s
-1
[Fe3+] / mM
fresh
5 hours
24 hours
2 4 6 8
Fig. 9. Effect of Fe(III) oxalate concentration (@ Ox/Fe=2.5) on the
gold dissolution kinetics. (Conditions: as Fig. 3).
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concentrations were kept constant. It is surprising that
increasing thiosulfate concentration to 0.4 M does not
increase the reaction rate significantly, since it is well
known that increasing the thiosulfate concentration is
known to increase the rate of gold oxidation in
thiosulfate+thiourea solutions (Chandra and Jeffrey,
2004). However it should be remembered that at
higher thiosulfate concentrations, the reaction
between the ferric di-oxalato species Fe(C2O4)2 and
thiosulfate is likely to occur at a higher rate. This was
confirmed by measuring the mixed potential, which
was found to decrease with increasing thiosulfate
concentration. For 50 mM thiosulfate solutions, the
mixed potential was initially 0.285 V, whilst for the
400 mM thiosulfate, the mixed potential was initially
0.17 V. Hence it appears as though the gain in gold
oxidation rate at higher thiosulfate concentrations is
offset by the lower potential of the Fe(III)Fe(II)
couple. Since the kinetics of the gold dissolution are
largely independent of thiosulfate concentration
between 0.05 M to 0.4 M, it was thus decided touse 0.05 M thiosulfate for the remainder of the present
study.
3.8. Comparison with the copperammoniathiosul-
fate system
A set of experiments were conducted to compare
the ferric-oxalate leaching with the traditional cop-
perammonia thiosulfate system. In these experi-
ments, both solutions contained 50 mM (NH4)2S2O3,
and they were carried out at 20 8C. The concen-
tration of the oxidant added was 5 mM [copper(II) or
iron(III)], whilst the concentration of ligand for the
copperammonia system (100 mM ammonia) wassignificantly higher than the concentration of ligand
for the ferricoxalate system (12.5 mM oxalate).
During leaching, the solutions were left open to air
in a beaker to allow the mass transfer of oxygen into
solution. Table 2 shows the gold dissolution rate as a
function of time for both of these systems, and it can
be seen that the rate obtained in the ferric oxalate
solution is considerably higher than that obtained for
the aged copperammonia system. These results
highlight the major problem with the copper
ammonia system in that copper(II) is quite reactivetowards thiosulfate so that the majority of copper is
present as Cu(I). It has been shown previously that
the copper(II) to copper(I) ratio is low, even in the
presence of oxygen, and efforts to improve mass
transfer of oxygen such as sparging or increasing the
oxygen partial pressure, do not result in significant
changes in the copper(II) concentration (Breuer and
Jeffrey, 2003). Hence the results presented in Table 2
indicate that in terms of gold dissolution kinetics, the
ferricoxalate oxidant in the presence of thiourea
catalyst is a viable alternative to the copper
ammonia oxidant.
The gold dissolution rate in ferric oxalate solutions
can also be compared to the published leach rates for a
gold/silver alloy containing 2% silver in 20 mM air
saturated cyanide solutions, 3.8105 mol m2 s1.The initial leach rate in the ferric oxalate system,
2.4105 mol m2 s1, is similar to that for cyanidesolutions, although the leach rate does decrease with
time.
Table 2
Comparison of gold leach rates for the ferricoxalatethiourea andcopperammonia systems after ageing
Time Gold leach rate in 0.05 M
(NH4)2S2O3 at 20 8C
5 mM Fe(III),
12.5 mM oxalatea5 mM Cu(II),
100 mM ammonia
10 min 2.4105 mol m2 s1 0.75105 mol m2 s1
5 h 1.1105 mol m2 s1 0.15105 mol m2 s1
24 h 0.94105 mol m2 s1 0.075105 mol m2 s1
The oxidant concentration given is based on the total amount added
to solution.a Also contains 5 mM thiourea catalyst.
0 100 200 300 4000.0
0.5
1.0
1.5
2.0
2.5
3.0
105r/molm-2s
-1
[S2O
32-]/ mM
Fresh5 hours24 hours
Fig. 10. Effect of thiosulfate concentration on the gold dissolutionkinetics. (Conditions: as Fig. 3).
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3.9. Effect of cations in thiosulfate salts
It has been shown previously that gold oxidizes
much more readily in ammonium and potassiumthiosulfate solutions than in sodium thiosulfate
solutions (Chandra and Jeffrey, 2004). Therefore, the
effect of changing the cation in the thiosulfate salt on
gold dissolution in the ferric oxalate system was
investigated. Fig. 11 shows comparison between gold
oxidation rates with time in ammonium, sodium and
potassium thiosulfate solutions. There is very little
difference between the gold dissolution kinetics
obtained with the three different thiosulfate salts.
Such a result is important, as it means that sodium
thiosulfate can be used instead of ammonium thio-sulfate, and hence the environmental issues of
ammonium discharge are avoided. At present, the
cost of sodium thiosulfate is significantly more than
ammonium thiosulfate, but if a sodium-based process
were adopted industrially, the cost of sodium thio-
sulfate may decrease.
3.10. Effect of pH
The data in Fig. 12 show the effect of pH on the
rate of gold dissolution in the presence of thiourea
catalyst. In these experiments, 10 mM potassium
hydrogen phthalate was added as a buffer. The pH was
varied from 4.0 to 6.2 by the addition of sodium
hydroxide. It was found that pH 4.0 was the lowest
pH that can effectively be investigated as below this
pH, thiosulfate decomposes to sulfur. At pH values
above 6.2, iron(III) hydroxide begins to precipitate
from the solution. Fig. 12 shows that pH has some
effect on the initial gold dissolution rates, with lower
dissolution rates being obtained at higher pH values.
However, the gold dissolution rates obtained after 24
h are relatively unaffected by change of pH. These
results are consistent with the Fe(C2O4)2 complex
being less oxidizing at the higher pH values. The fact
that the mixed potential for the fresh solutions
decreased as the pH was increased supports this view.
One possible explanation for this may be the com-
plexation of hydroxide to the Fe(C2O4)2 at higher pH
values.
4. Conclusions
The dissolution of gold in thiosulfate solutions
containing thiourea catalyst using iron(III) oxalate has
been investigated by use of the REQCM. It was foundthat the presence of thiourea is required in order to
obtain high leach rates. The ratio of ferric to oxalate is
important in the process as this determines the EH and
the dominant ferric complex. At an Ox/Fe mole ratio
of 3, the gold leach rate was essentially constant over
24 h. However, at lower Ox/Fe ratios, the Fe(C2O4)2
complex formed is more reactive towards thiosulfate
and/or thiourea, and hence the gold leach rate and
solution potential, which are initially very high,
decrease with time. Studies undertaken with an Ox/
0 5 10 15 20 250.0
0.5
1.0
1.5
2.0
2.5
3.0
105r/molm
-2s
-1
time / hr
Na2S
2O
3
K2S
2O
3
(NH4)2S
2O
3
Fig. 11. Effect of alkali cation on the gold dissolution kinetics.
(Conditions: as Fig. 3 except 0.05 M thiosulfate).
4.0 4.4 4.8 5.2 5.6 6.0 6.40.0
0.5
1.0
1.5
2.0
2.5
3.0
fresh24 hrs
105r/molm-2s
-1
pH
Fig. 12. Effect of pH on the gold dissolution kinetics. (Conditions:
as Fig. 3 except buffered with 10 mM potassium hydrogenphthalate).
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Fe ratio of 2.5 showed that initially the gold leach rate
was first order with respect to Fe(III) concentration
and was at least mixed controlled. However, after
allowing the solution to age, the reaction becamechemically controlled and essentially independent of
the Fe(III) concentration. The gold leach rates of the
ferricoxalate oxidant compare very favorably to the
traditional copperammonia oxidant under similar
conditions. Ferrous oxalate oxidation by air appears
to be slow and thiosulfate consumption very low.
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