solution chemistry of iodide leaching of gold

Hydrometallurgy, 32 (1993) 143-159 143 Elsevier Science Publishers B.V., Amsterdam

Solution chemistry of iodide leaching of gold

A. Davis, T. Tran and D.R. Young Department of Mineral Processing and Extractive Metallurgy, School of Mines,

The University of New South Wales, Kensington, NSW, Australia

(Received June l, 1990; revised version accepted March 4, 1992 )

ABSTRACT

Davis, A., Tran, T. and Young, D.R., 1993. Solution chemistry of iodide leaching of gold. Hydrome- tallurgy, 32: 143-159.

The thermodynamic equilibria and kinetic aspects of gold dissolution in iodide electrolytes have been studied with emphasis on the effect of different oxidants on the system. In conjunction with kinetic measurements, the CHEMIX computer program was used to predict the concentration profiles of the predominant species at equilibrium in different solution conditions for the systems Au-I - -12- H20 and A u - I - - O C I - - H 2 0 .

The thermodynamic study showed that I~- is the predominant oxidant species in both systems. However, if the concentrations of OC1- and 1 are equal, solid iodine is formed. In these systems iodide ( I - ) is used to form I~- (responsible for the gold oxidation) and more free iodide needed for the gold complexation is destroyed in the I - -OC1- system than the I -12 system. The formation of solid Aul also explains the lower rate of gold dissolution determined for certain conditions in the kinetic study.

The thermodynamic modelling supports the kinetic measurements which show that, although the I - - O C I - system has a higher oxidation capacity, it does not extract gold as well as the I--12 system. In all cases there exist optimum oxidant/iodide ratios for achieving maximum gold extraction rates. A mixture which has the highest I~- and free I concentration will attain the best gold extraction rate.

INTRODUCTION

Iodide-iodine has been suggested in various patented processes for extract- ing gold from scrap materials [1-3] or in-situ operations [4-5 ]. This lixi- viant system possesses certain advantages such as low oxidizing potential and high solution stability compared with other halide systems [ 6 ]. The kinetics of gold dissolution under different conditions of pH, concentrations of iodide and the oxidants (iodine and hypochlorite) has been studied [ 7 ]. It has been established that, at ambient temperature, gold dissolution by a mixture of iodide and iodine is strongly dependent on the iodide concentration (con-

Correspondence to: Dr. T. Tan, Univ. of New South Wales, School of Mines, Dept. of Mineral Processing and Extractive Metallurgy, P.O. Box 1, Kensington, N.S.W. 2033 Australia.

0304-386X/93/$06.00 © 1993 Elsevier Science Publishers B.V. All rights reserved.

1 4 4 A. DAVIS ET AL.

trolling the oxidation half-cell) and the iodine concentration (controlling the reduction half-cell). The gold oxidation rate decreases at a pH above 10 (depending on the iodide concentration). Iodine is a good reagent as long as the pH is below 11, where the predominant oxidizing species is tri-iodide. On a molar basis, iodine performs better than hypochlorite because higher gold dissolution rates can be achieved over a wide pH range. Fundamental kinetic and electrochemical studies by other researchers [ 8 ] have also established the effects of agitation, temperature, pH, iodide and iodine concentration on the rate of gold dissolution. Cyclic voltammograms of gold in the presence of the halogens were used to show that, for a given potential, the rate of gold oxidation is much faster in iodide solution than in either chloride or bromide.

In order to understand the solution chemistry of iodide leaching of gold, the thermodynamic constraints affecting the speciation of aqueous iodide solutions in the presence of gold need to be studied to determine the soluble species which are likely to be generated in the system. The CHEMIX computer program [ 9 ] has been used to determine the equilibrium conditions that pre- vail in aqueous iodide leaching of gold at atmospheric conditions. The program calculates the particular species and phases and their quantities which give the minimum total Gibbs energy for the system at a given temperature and pressure by the Eriksson minimization method [ 10 ].

In this paper the equilibria of the iodide-oxidant and gold-iodide-oxidant systems are presented in terms of the predominant species concentration versus pH diagrams. Such diagrams are used to elucidate the results of gold dissolution studies in different aqueous iodide-oxidant systems. The two oxidants evaluated in the study are iodine and hypochlorite.

E X P E R I M E N T A L

Reagents

All reagents used in the study were of analytical grade. High-purity gold strips were used in dissolution experiments.

Apparatus and equipment

Gold strips of known surface areas were used in dissolution tests which were conducted with different stirred mixtures of iodide and an oxidant (iodine and hypochlorite). The average rate of gold dissolution (measured as mg/cm2h) was determined from the gold weight loss measured with a micro- balance (Mettler M3 ) and confirmed by solution chemical analysis using a Varian atomic absorption spectrophotometer. The reduction rate of various oxidants was determined from current-potential relationships using a rotat- ing disc electrode system (Pine Instrument RDE4).

SOLUTION CHEMISTRY OF IODIDE LEACHING OF GOLD [ 4 5

CHEMIX computer program

The CHEMIX program is part of a THERMODATA package developed by the CSIRO Division of Minerals Products (Melbourne, Australia). In this study the program was run on an Olivetti M386 personal computer and the results extracted from the calculations were plotted with a Roland DX990 digital plotter.

RESULTS AND DISCUSSION

Thermodynamic considerations

All gold dissolution processes rely on a proper choice of suitable oxidants (oxygen, hydrogen peroxide, etc.) and the complexants (thiourea, cyanide, etc. ). Gold is very stable in an aqueous system unless a complexant is added to govern its stability in a potential range within the water region. To achieve this an oxidant is simultaneously added to bring the gold surface to a mixed potential and dissolve it as a complex. The dissolution of gold in an iodide electrolyte is an electrochemical process which can be illustrated as:

Au+2I -~AuI~£ + e - (anodic) (1)

and

I3 + 2 e - - ~ 3 I - (cathodic) (2)

The Nernst equations related to these half-cells are expressed as:

Standard reduction potential, V versus SHE (standard reduction potentials calculated with standard free-energy data in Appendix 1 )

El =E? + (RT/F)In{[AuIS- ]/[I- ]2} 0.578V (3)

and

E2=ET+(RT/2F)ln{[ly]/[I-] 3} 0.535V (4)

The overall reaction to form aqueous gold iodide complex, AuI2 is therefore:

2 A u + I ~ + I - ~ 2 A u I £ (5)

To dissolve gold in an iodide solution there seem to be several oxidants that can be used. The effect of oxidants, iodine and hypochlorite on the solution chemistry of the aqueous gold-iodide system was investigated. As the chemistry of the iodide-oxidant system is complex, with different species existing in the solution, their concentration distributions were modelled using CHEMIX in an at tempt to characterize these mixtures.

146 A. DAVIS ET AL.

Thermodynamic data and computational methods

The equilibrium concentrations of the predominant species presented in this paper were generated with the aid of thermodynamic data taken from the literature [ 9-12 ] and the CHEMIX computer program [ 9 ]. Detailed tabula- tion of these data are presented in Appendix 1. All possible compounds which are known to exist in the system and for which thermodynamic data could be found were considered. A discussion on the thermodynamic concepts behind the Gibbs energy minimization method has been presented by Eriksson [ 10 ] and will not be repeated here. All the diagrams were constructed for 25 °C and at atmospheric pressure. The symbols (g), (aq) and (s) designate gas- eous, aqueous and solid phases, respectively.

It was found that the gas phase input for CHEMIX is very sensitive to the final equilibrium results. If air is left to equilibrate with the solution; that is, the solution potential is equivalent to the oxygen potential of air (at 0.21 atm oxygen ), all iodide in the solution will be oxidized to iodine thermodynam- ically. Kinetically, however, this oxidation should be slow, especially as the solubility of oxygen from saturated air is less than l0 ppm. The solution chemistry at equilibrium therefore has to be treated independently from the atmosphere/solution equilibration. In this study the solutions are equili- brated with an inert nitrogen gas atmosphere.

The I--I2-H20 system

As shown in a previous study [7 ] the equilibrium involved in the iodide- iodine solution can be predicted by establishing three reactions involving final species such as I - , I~-, I2, IO- and HIO. These reactions yield equations which relate concentrations of these species and hydrogen ion to their appro- priate equilibrium constants. To allow the calculation of analytical results the system has to be simplified by not considering IO7 and the input concentrations are chosen such that iodine concentration is much smaller than iodide concentration. The analytical solution of these equations yields tri-iodide as the predominant species over a pH range from 0 to 11. In these calculations since IOn- is not considered IO- is the predominant species at pH values higher than 11.

Using CHEMIX the concentrations of predominant species versus pH for the

SOLUTION CHEMISTRY OF IODIDE LEACHING OF GOLD 147

system I - - I 2 - H 2 0 at 25°C with initial input conditions [ I - ] = 0 . 1 M and [I2] = 50 mMwere calculated with I~-, I2, IO- , HIO and IOn- as aqueous and I2 as solid species. Another polyiodide, I ; , was not considered due to the lack of free energy data for this species. The exclusion of I~ should not alter the equilibrium concentration profile as the I2 species (needed to form I5 according to the reaction I~-+I2-,I~-) exists at a much lower concentration. The predominant oxidant species indicated in Fig. 1 are I~- and IO; . At a pH below 9.5, I ; (tri-iodide) is the predominant species while above pH 9.5, IOn- (iodate) is the only predominant species in the system. The same concentration profiles were observed with 0.5 and 10 m M I2 in 0.1 M I- system.

The thermodynamic equilibria have shown that I~- (tri-iodide) is the predominant oxidant in an iodide-iodine mixture responsible for gold dissolution at ambient temperatures. The equilibrium responsible for the formation of tri-iodide is:

I - + I 2 = I j - (6)

The I--OCI --H20 system

When hypochlorite is added into an iodide solution it is expected that OI- is first formed according to the reaction:

OCI- + I - ~ O I - +C1- (7)

The solution re-equilibrates to form I2 and then I; according to the following equations:

2H ÷ + O I - + I - =I2 + H 2 0 (8)

8O

z 6 0 r c) H ! L

50 o o - o o 02 I-- z 4O uJ Q3 Z O 3O Q}

U0 LU 20 LJ LU £L 10 (/3

0 I 14

o 1 3 - (IRQ)

zs I 2 (AG)

O I0 3 - (RG)

I 2 ( S )

o ~ o o

\ o ~ / o - O . . . .

/ o A-A--A A - - - ~ - - Z X I Z ~ /

2 3 4 5 6 7 8 9 10 11 f2 13

p H

Fig. 1. Concentration profiles of different species at 0.1 M I-, 5 0 mM total 12 and pH 0-14.

148 A. D A V I S E T AL.

I -+I2=I~- (9)

The overall reaction to form I~- is, therefore:

OC1- + 2 H + + 31- =I~- + H 2 0 + C 1 - (10)

According to this stoichiometry the formation of I~- is solely from I- and the free I- left will be less than for similar conditions illustrated in eq. (6) . The thermodynamic modelling confirms these reaction paths, which show similar concentration profiles of the oxidizing species to the I - - I2 -H20 system at low OC1- concentrations of 0.5 and 10 m M (Fig. 2a) but differ signif- icantly at 50 m M OCl- , as shown in Fig. 2b. At this level of hypochlorite, solid 12 is the predominant species at pH values lower than 8 and I~- only

0

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12 ( S )

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i i ©

S /Q-<> I "\

( b ) 0 . . . . ~ ~ - ~ ' ~ o ~ ,~..~s_°~._o___~ I 2 3 4 5 6 7 8 9 !0 !i 12 13

p H

Fig. 2. (a) Concentration profiles of different species at 0.1 M I- , 0.5 mM OCI and pH 0- ! 4. (b) Concentration profiles of different species, 0.1 M I- , 50 mM OCI- and pH 0-14.

SOLUTION CHEMISTRY OF IODIDE LEACHING OF GOLD [ 49

exists in the solution in the pH range 6-9 which passes through a maximum at a pH around 8. Around this pH, I~- is also the predominant species. IO~ is still the predominant species at pH values higher than 10.

A significant difference in the two systems was observed in the I (iodide) concentration profile as shown in Fig. 3. The free iodide concentration for the iodide-hypochlori te system is always lower than for the iodide-iodine system over the entire pH range from pH 2 to 12. In the iodide-iodine system, at a pH lower than 8 the concentration of iodide is reduced by the exact stoichiometry of 1 : 1 of the reaction with iodine to form tri-iodide; whereas at a pH, higher than 10, the initial concentration of iodide at 0.1 M is increased by the decomposit ion of iodine to form IO£ (iodate). In all calculations the mass balance of species I (iodine, iodide, tri-iodide and iodate) is preserved.

For the iodide-hypochlori te system, the stoichiometry of the reaction between iodide and hypochlorite is 3: 1 in the cases where tri-iodide is formed. At 50 m M hypochlorite, the concentration of iodide at pH values lower than 7 is reduced to nearly zero, indicating a complete reaction between 0.1 M iodide and 50 m M hypochlorite to form mostly solid I2 with a stoichiometry of 2:1 in this case. The initial concentration of iodide (0.1 M) is reduced by the same stoichiometry as the formation of IO~ at a pH higher than 10 in all cases.

Oxidation capacity of iodide-oxidant solutions

Scanning voltammetry was used to determine the reduction rate of different mixtures of iodide-iodine (or hypochlorite). The current-potential curves show a clear diffusion zone at potentials lower than 0.0 V versus SCE. Figure

223 I ~ 293t 8XZDANT CONCENTRATION

1 0 0.5 mM I 2 ]80 ,5. 0 , 5 mM OCJ /+" + +

mO 189l [] 10 mM 12 ,¢' / ~ 10 mM OCI /

F: 14s~ +~0~M 123 ~ o 50 mH OC I "~ oO []

I00 i o s o - o a o . . . . . . . ' % o o o o o o-

} 8 0 / /~9~ o o - o

/ + + + + . . + ,

/ u I 43 / D d

so # 0 0

I 2 3 4 5 6 7 8 g 10 1! 12 13 14 p H

Fig. 3. Comparison of free iodide concentration profiles in 0.1 M KI using different oxidants, pH 0-14.


P@TENT IRL (V) -I .0 -0.8 0.6 0.4 -0.2 O 0.2

(a) . . . . . ' . . . . . , , j [

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3000

Fig. 4. ( a ) C u r r e n t - p o t e n t i a l curves (a t scan rate 50 m V s -1 ) fo r the Reduc t i on o f iod ine on gold disc electrode (surface area 0.869 mm 2) in 0.1 M KI, 0.5 mM total 12 and pH 6.5. (b) Current-potential curves (at scan rate 50 mV s -1) for the Reduction of iodine on gold disc electrode (surface area 0.869 mm 2) in 0.1 M KI, 0.5 mM total I2 and pH 12.3.

4 shows the difference of the two typical scans at pH 6.5 and pH 12.3. The first steady state current believed to be due to the diffusion of I~- species is not obvious in the high pH range. Also, the solutions o f p H 6.5 possess a much higher oxidation capacity (higher reduction currents) in the potential range of - 0 . 4 - 0 . 2 V versus SCE. The limiting current densities of the 0.5 m M iodine-0.1 M iodide at different pH values follow Levich's equation [ 13 ], as shown in Fig. 5. (Data points for pH 6.54, 9.53 and 10.50 are also plotted, although hidden behind other points). The reduction rates of the iodide- iodine and iodide-hypochlor i te solutions at different pH values, as measured by the slopes of the Levich's lines of Fig. 5, are shown in Fig. 6. Compared with an iodide- iodine system of the same concentrations the reduction of the iodide-hypochlor i te solution seems to pass through two inflection points. In the more acidic zone, the reduction rate of the iodide-hypochlori te mixture is also about 30% higher than for the iodide- iodine system. The same behav-


~I. I.J pH "TA[ LJES

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c~: 1.6'- • 8 . 4 / ~ 0 . 5 3 o 1 / . 5 0

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LU / O . 8.8= ~ . - ~

0 .6 ! o ~ ~ .e - - 1

0 2 4 6 8 10 17/2 114 2 16 18 20 ~FRoTRTION SPEED ( t a d / s e c )

Fig. 5. Levich's relationship for the reduction of 0.5 mM iodine in 0.1 M KI at different pH values.

O' ; I f ~ 0.10

• F C,; " 0.08 J

E 0.07~ ()

o,oe!

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0.02~

O.Ol

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o o ~ \o \

[] a q

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Fig. 6. Effect ofpH on the reduction rate of different oxidants (0.5 mM OC1- or I2)-iodide (0. l M KI) mixtures as measured by the slopes of the limiting current density versus 091/2 lines.

iour was also observed for the solutions o f 5 m M hypochlorite in 0.1 M KI. The higher reduction rate of the iodide-hypochlorite mixture in the more acidic zone explains its higher rate o f gold dissolution for low additions of the oxidant (as shown in Fig. 1 1 ).

These kinetic results at pH < 6 seem to contradict with the thermodynamic modell ing data. This is because, at equilibrium, both systems should attain the same level o f I ; . The only explanation found at this stage is that the solutions prepared for the kinetic studies did not reach equilibrium completely. Therefore, the speciation o f OC1-/C12 species still exists, which tends to fa- vour the formation o f chlorine at a low pH according to the reaction:

152 A. DAVIS ET AL.

OC1- + H + + C1- ~ C12 + O H - (11)

Gold dissolution in iodide solutions

The Au-I--Ie-HeO system Figures 7a-c present the equilibrium concentrations of predominant spe-

cies versus pH for the Au -I - - I2 -H20 system at 25 °C. The initial input conditions are solid gold at 0.1 mol/1 solution, [ I - ] =0.1 M and [12] =0.01 M, 0.035 M and 0.05 M. The species considered were Iy, I2, IO- , HIO, IOy, AuI~- and AuI4 as aqueous and AuI as solid. In Fig. 7a and 7b, the predominant equilibrium species are I j-, AuI~-, 12 and IOy. The concentration profile of AuI£ species is influenced by the concentration of I~- which dominates in the solution below pH 9.5. Above pH 9.5 IOn- becomes the dominant species in the solution and the gold iodide complex, AuI~- becomes unstable above pH 10.

These calculations indicate that although there exist in the solution a sur- plus of gold, tri-iodide (Fig. 7a-c) and free iodide (Fig. 8 ) at the final equilibrium in all cases with initially 10, 35 m M iodine and 0.1 M iodide, the soluble gold species AuI~- reaches a maximum concentration of 4.8 mM and 5.8 m M respectively. This means that the electrode potentials of the two half- cell reactions ( 1 ) and (2) involved in the gold dissolution are equal under these conditions. For the solution of initially 0.1 M iodide and 10 m M iodine, substituting the equilibrium concentrations of iodide, tri-iodide and soluble gold iodide AuIy species as 90, 7.25 and 4.8 m M respectively to the Nernst equations, eqs. (3) and (4), yield both El and E2 values of 0.565 V versus SHE.

At 50 m M 12 initially (Fig. 7c), the predominant species are I~-, AuI~-, 12

and IO~- as aqueous and AuI as solid. The formation of solid AuI species in the solution below pH 8 could passivate the gold surface and inhibit AuI~- formation even though a higher solution potential exists due to the higher concentration of I~- and a lower free iodide concentration in the solution. Above pH 9.5, IO~- becomes the predominant species in the solution and AuI£ complex is unstable above pH 10.

AuI£ becomes unstable at pH > 10 and will decompose to precipitate gold. This observation on thermodynamic modelling is confirmed in a gold recovery process where Au is removed from gold iodide solution simply by adjust- ing the solution pH [ 14 ]. Naturally, kinetic considerations have to be given as this can only be applied commercially with solutions of high soluble gold concentrations (e.g. liquors from leaching of electronic scraps ).


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i 13 14

Fig. 7. (a) Concentration profiles of different species in solution with initially solid gold at 0.1 mol/I solution, 0.1 M 1% 0.01 M total I2 and pH 0-14. (b) Concentration profiles of different species in solution with initially solid gold at 0.1 mol / l solution, 0.1 M I- , 0.035 M total I2 and pH 0-14. (c) Concentration profiles of different species in solution with initially solid gold at 0.1 mol / l solution, 0.1 M I - , 0.05 M total I2 and pH 0-14.


220 1

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[11

OXIDANT CONCENTRATION

o 10 mR 12

A 10 mM OCI- .+--+ + ÷

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Fig. 8. Comparison of free iodide concentration profiles in solutions using different oxidants with initially solid gold at 0.1 mol/1 solution, 0.1 M KI and pH 0-14.

The Au-I--0Cl--1-120 system Figures 9a-c show the equilibrium concentrations of predominant species

versus pH for the Au-I - - O C 1 - - H 2 0 system at 25 °C. The initial input conditions are solid gold at 0.1 mol/1 solution, [ I - ] =0.1 M a n d [OC1- ] =0.01 M, 0.025 M and 0.05 M. The species considered were similar to those for the A u - I - - I 2 - H 2 0 system. It seems that the behaviour of hypochlorite is similar to that of iodine, except for the formation of solid gold iodide species, AuI (S), below pH 8 at a lower hypochlorite concentration of 0.025 M compared to the iodide-iodine system where solid AuI exists at a significant level only at iodine concentrations of 50 m M or higher. This means that passivation of the gold surface below pH 8 could occur at a lower concentration of hypochlorite in an iodide-hypochlorite solution compared to the iodide-iodine solution.

At 50 m M hypochlorite (Fig. 9c), AuI(S) is the only gold product and the formation of the AuI~- complex in the solution is completely inhibited due to the depletion of iodide at pH values < 7. Aqueous Ij- and AuI£ species are only present in the solution within a small range of pH 7-9. At a pH higher than 10 no gold can be dissolved and only iodate and iodide exist in the solution.

Related kinetic aspects of the dissolution of gold

The dissolution of gold in the iodide-iodine system is dependent on the solution pH and concentrations of iodide and iodine. The rate of the gold oxidation half-cell reaction increases at an increasing iodide concentration and is affected by the pH. To optimize the iodide-iodine mole ratio of a fixed initial iodide concentration, a series of gold dissolution tests was carried out

SOLUTION CHEMISTRY OF IODIDE LEACHING OF GOLD 15 5

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Z 5 C3 4

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pH

Fig. 9. (a) Concentration profiles of different species in solution with initially solid gold at 0.1 mol / l solution, 0.1 M I - , 0.01 M OC1- and pH 0-14. (b) Concentration profiles of different species in solution with initially solid gold at 0.1 mol/1 solution, 0.1 M 1-, 0.025 M OCI- and pH 0-14. (c) Concentration profiles of different species in solution with initially solid gold at 0.1 tool/1 solution, 0.05 M O C I - and pH 0-14.

156 A. DAVIS ET AL.

at three different pH values (2.7, 7.0 and 11.5). At 0.1 M K I the gold dissolution rate increases as more iodine is added into the solution. Depending on the solution pH, this rate reaches a peak at an addition of 35-40 m M iodine (Fig. 10). The lower rate of gold dissolution at pH 11.5 is due to the disap- pearance of the more reactive tri-iodide at this pH (Fig. 7). The maxima observed in Fig. 10 confirm the passivation of the gold surface due to the formation of solid gold-iodide species, AuI (S) at iodine concentrations at 50 m M or higher. Under the conditions tested at ambient temperatures, the op- t imum iodine/ iodide ratio is within 0.35-0.4.

Although hypochlorite is a much stronger oxidant, its gold dissolution per- formance in an iodide solution is not as good as with iodine. The gold dissolution rate in the hypochlorite-iodide system is greatly dependent on the solution pH as predicted from reaction ( 10 ). As shown in Fig. 1 l, for 0.1 M KI at pH 2.7, a maximum is observed at an addition of 25 m M hypochlorite, corresponding to the formation of AuI (S) at and above this level. The higher gold dissolution rate observed at pH 2.7 compared with the more neutral pH 7.2 is due to the stronger oxidation capacity of iodide-hypochlorite solution in acidic zone as shown in Fig. 6. For mixtures of 0.1 M KI and > 25 mM hypochlorite the gold dissolution rate at pH 2.7, however, is lower than in solutions at pH 7.2. This dramatic effect is again due to the passivation of the gold surface by the formation of the solid gold-iodide species, AuI (S), below pH 7, as predicted from the thermodynamic equilibria (Fig. 9b and c). At a higher hypochlorite concentration (50 mM), the aqueous gold iodide complex, AuI2, is only stable in the solution within a pH range 7-9 (Fig. 9c), where gold dissolution is possible.

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,.i ! 0

0

pH VRLUES [ K I ] = 0 . 1 i1 o 2 . 7 A Ternp.=20OC (+1 o )

5_ 7 , 0 / ~ d ~ ~. ,/ A

[] I 1 . 5 / / O ~ \ / o \ O /o

/ /

5 ~

/ O

5 10 15 20 25 30 35 40 45 50 65 60 65 70 75 80

1 2 C O N C E N T R R T I O N ( m M )

Fig. 10. Variation of gold dissolution rate with iodine additions, 0.1 M KI, pH 2.7, 7.0 and 11.5, 20°C.

SOLUTION CHEMISTRY OF IODIDE LEACHING OF GOLD 1 5 7

10 F

D: [ K I ] 0.1 M pH VALUES

o.," ! Temp.=2@°C (+ I ° ) o 2 . 7 E 8! z~ 7 .8 ;

'~) 7b [] 1 ! . 5 ,~

0 5 10 15 20 25 30 35 40 45 50 55 68 6.6 70 OCI CONCENTRRT I ON ( m M }

75 80

Fig. 11. Variation of gold dissolution rate with hypochlorite additions, 0.1 M KI, pH 2.7, 7.2 and 11.5, 20°C.

CONCLUSIONS

The variation of gold dissolution rate under various iodide solution conditions can be explained from the concentrations olin- and I - as predicted from CHEMIX. The stability diagrams developed show that the predominant aqueous species in the solution are I~-, AuI~-, I2 and IOn- at low concentrations of the oxidants (I2 and OCl- ). The concentration profiles of I~- and AuI7 species are similar, indicating that gold dissolution is influenced by the concentration of I~- which dominates in the solution below pH 9.5. Above pH 9.5, IOn- is the dominant species in the solution and AuI~- becomes unstable. The gold dissolution rate is very low above pH 10.

At high oxidant concentrations the predominant species in the solution are I~-, AuIF, I~- and IOn- as aqueous and AuI as solid. The formation ofAuI (S) occurs below pH 8 and results in a shortage of free iodide species, I - in the solution. The passivation of the gold surface due to the formation of AuI (S) and the shortage of I - could account for the significant drop in gold dissolution rate below pH 8 at high oxidant concentrations.

ACKNOWLEDGEMENTS

We are grateful for a research scholarship from the University of New South Wales and thank Dr. D. McDonald of BHP Minerals for his comments on the manuscripts.

158 A. DAVIS ET AL

APPENDIX 1

T h e r m o d y n a m i c D a t a on go ld- iod ide sys tem at 25 °C and 0.1 M Pa

F o r m u l a State AG ° k J m o l - I F o r m u l a State A G ° k J m o l -~

I2 g 19.339 KI s - 3 2 3 . 1 3 7 HI g 1.7 K + aq - 2 8 3 . 2 8 7 0 2 g 0 C1- aq - 131.269 N2 g 0 C12 aq 6.94 H2 g 0 Cly aq - 120.4

C12 g 0 C 1 0 - aq - 36.306 HC1 g - 95.267 C102 aq 120.180 HzO g - 2 2 8 . 5 3 8 C l O y aq 17.195

I2 S 0 C10~- aq - 7.95 I2 aq 16.4 C10~- aq - 8.534

I - aq - 5 1 . 5 7 HC1 ai - 131.269 I~- aq - 51.4 H C 1 0 aq - 79.9

I O - aq - 38.5 HC102 aq 5.9 IO3 aq - 128.0 HCIO3 aq - 7 . 9 5 IOn- aq - 58.5 HC104 aq - 8.52

I202- aq - 8 2 . 4 NaC1 ai - 3 9 3 . 1 8 5 HI ai - 5 1 . 5 7 N a C 1 0 ai - 2 9 8 . 0 1 2 H I O aq - 9 9 . 1 N a + aq - 2 6 1 . 9 2 4 HIO3 aq - 132.6 C11~- aq - 116.320 H2IO + aq - 106.7 KC1 ai - 4 1 4 . 5 4 7 I 2 O H - aq - 230.1 KC1 I 2 ai - 399.695 H20 aq - 2 3 7 . 1 4 1 K C 1 0 ai - 3 1 9 . 9 6 0 O H - aq - 157.245 AuC1 s - 15.035 H + aq 0 AuC13 s - 4 5 . 3 2 7 Au s 0 A u C I 2 aq - 151.12

AuI s - 0 . 5 AuC1z aq - 2 3 5 . 1 4 8 A u ( O H ) 3 s - 3 1 6 . 9 2 A u C 1 3 O H - aq - 3 0 6 . 0 AuO33- aq - 5 1 . 8 AuCIz (OH)~- aq - 3 7 1 . 2 8 HAuO3 z - aq - 124.2 A u C I ( O H ) ~ - aq - 4 3 1 . 5 6

HzAuO 3 aq - 2 1 8 . 3 HAuC14 ai - 2 3 5 . 1 4 8

A u ( O H ) 3 aq - 2 8 3 . 3 7 NaAuC14 ai - 4 9 7 . 0 5 2 AuI~- aq - -47 .37 HzSO 4 aq - 169.147 A u I z aq -- 44.18 SO42- aq - 744.495 Au + aq 176.0 HSO~- aq - 7 5 5 . 8 6 6 Au 3 + aq 440.0


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