Un estudio del Oxalato férrico para la disolución de Au en soluciones de Tiosulfato

8/8/2019 Un estudio del Oxalato frrico para la disolucin de Au en soluciones de Tiosulfato

1/11


2/11

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


3/11

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


3=4.9 mM 0.312 V

Fe(C2O4)22=0.1 mM

C2O42=0.1 mM


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].

I. Chandra, M.I. Jeffrey / Hydrometallurgy 77 (2005) 191201 193


4/11

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).



5/11

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).



6/11

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).



7/11

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).



8/11

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).



9/11

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).



10/11

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).



11/11

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.

References

Aylmore, M.G., Muir, D.M., 2001. Thiosulfate leaching of goldareview. Miner. Eng. 14 (2), 135174.

Breuer, P.L., Jeffrey, M.I., 2003. Copper catalysed oxidation of

thiosulfate by oxygen in gold leach solutions. Miner. Eng. 16

(1), 2130.

Chandra, I., Jeffrey, M.I., 2004. An electrochemical study of the

effect of additives and electrolyte on the dissolution of gold in

thiosulfate solutions. Hydrometallurgy 73, 305 312.

Chu, C.K., Breuer, P.L., Jeffrey, M.I., 2003. The impact of

thiosulfate oxidation products on the oxidation of gold in

ammonia thiosulfate solutions. Miner. Eng. 16 (3), 265 271.

Dai, X., Breuer, P.L., Jeffrey, M.I., 2005. The development of a flow

injection analysis method of the quantification of free cyanide

and copper cyanide complexes in gold leaching solutions.

Hydrometallurgy 76 (12), 8796.

Dudeney, A.W.L., Tarasova, I.I., 1998. Photochemical decomposi-

tion of trisoxalatoiron(III): a hydrometallurgical application of

daylight. Hydrometallurgy 47, 243 257.

Jeffrey, M.I., Zheng, J., Ritchie, I.M., 2000. Development of a

rotating electrochemical quartz crystal microbalance for the

study of metal leaching and deposition. Meas. Sci. Technol. 11

(5), 560567.

Jeffrey, M.I., Breuer, P.L., Choo, W.L., 2001. A kinetic study that

compares the leaching of gold in the cyanide, thiosulfate and

chloride systems. Metall. Mater. Trans. B 32B, 979 986.

Jeffrey, M.I., Breuer, P.L., Chu, C.K., 2003. The importance of

controlling oxygen addition during the thiosulfate leaching of

gold ores. Int. J. Miner. Process. 72 (14), 323330.

Lide, D.R., 1995. CRC Handbook of Chemistry and Physics. CRC

Press, Boca Raton, Fl.

Nicol, M.J., OMalley, G., 2002. Recovering gold from thiosulfate

leach pulps via ion exchange. JOM 54 (10), 4446.

Nicol, M., Fleming, C., Paul, R., 1987. The chemistry of theextraction of gold. In: Stanley, G.G. (Ed.), The Extractive

Metallurgy of Gold. South African Inst. Min. and Metall,

Johannesburg, S. Africa, pp. 831905.

Page, F.M., 1954. Reaction between ferric and thiosulfate ions:

II. Thiosulfate complexes of iron. Trans. Faraday Soc. 50,

120126.

Page, F.M., 1960. Ferric thiosulfate reaction: III. Mechanism of the

reaction. Trans. Faraday Soc. 56, 398406.

Smith, R.M., Martell, A.E., 1976. Critical Stability Constants. Other

Organic Ligands, vol. 3. Plenum Press, New York.

Tykodi, R.J., 1990. In praise of thiosulfate. J. Chem. Educ. 67 (2),

146149.


Un estudio del Oxalato férrico para la disolución de Au en soluciones de Tiosulfato

Documents

Transcript of Un estudio del Oxalato férrico para la disolución de Au en soluciones de Tiosulfato