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Electrochemical aspects insome of the hydrometallurgicalprocessesR. K. Paramguru aa Regional Research Laboratory (Council of Scientificand Industrial Research), Bhubaneswar, IndiaVersion of record first published: 26 Oct 2010.

To cite this article: R. K. Paramguru (2002): Electrochemical aspects in some of thehydrometallurgical processes, Mineral Processing and Extractive Metallurgy Review:An International Journal, 23:2, 65-100

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ELECTROCHEMICAL ASPECTS INSOMEOF THEHYDROMETALLURGICALPROCESSES

R.K. PARAMGURU

Regional Research Laboratory (Council of Scientific andIndustrial Research) Bhubaneswar, India

Hydrometallurgy deals with extraction of metal values from ore minerals

through aqueous processing by solubilizing the specific metal value into the

aqueous phase and subsequently taking it out of the solution. These steps may

be chemical or electrochemical in nature. The electrochemical phenomena

depends on the electrical properties of the solid material and the redox

characteristics of the solution. The present paper discusses electrochemical

aspects of some hydrometallurgical operations involving the following phe-

nomena: (i) corrosion coupling, (ii) galvanic coupling, (iii) dissolution via

cyclic action of a redox couple, (iv) displacement or cementation reaction, (v)

precipitation under reducing conditions, (vi) dissolution under high pressure,

and (vii) dissolution influenced by hole transfer.

Keywords: electrochemical aspects, hydrometallurgical operations, electron

transfer, hole transfer

Extractive metallurgy involves efficiently taking the valuable metal, present

as a compound, out of the ore body. Broadly, two types of operations

(i.e., pyro- and hydrometallurgy) are in practice. Although pyro-

metallurgical treatment is common, hydrometallurgical techniques are

generally preferable because of environmental and many other advantages.

The author wishes to thank Dr. V.N. Misra, Director, Regional Research Laboratory,

Bhubaneswar, for his keen interest in this paper and permission to publish this work.

Address correspondence to R. K. Paramguru, Regional Research Laboratory, Council

of Scientific and Industrial Research, Bhubaneswar, Orissa 751013, India.

E-mail: paramguru@rrlbhu.res.in

Mineral Processing and Extractive Metallurgy Review, 23: 65�100, 2002Copyright# 2002 Taylor & Francis

0882-7508/02 $12.00+ .00

DOI: 10.1080=08827500290110179

65

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Generally, hydrometallurgy operates in two distinct steps: (i) a

leaching step involving dissolution of the desired mineral(s) or a

processed mineral from the ore into solution, and (ii) the winning step for

the separation of metal values from the solution. Invariably, some other

operations, such as solid-liquid separation, solution purification, and

thickening, are also incorporated between these two steps.

Leaching is the process of selectively extracting a soluble con-

stituent from a solid into a solvent. In extractive metallurgy, this solid

can be a mineral or minerals present in an ore or concentrate, or a

metallurgical product, such as calcine, matte, scrap alloy, anodic slime,

etc. The solvent may range from simple water to strong acids, alkalis,

or aqueous salt solutions to provide a strong oxidizing or reducing

atmosphere during leaching or act as a complexing agent. The leaching

operation may be performed in situ, on a heap or dump, in a vat or

column via percolation or in a stirred tank reactor at ambient or

elevated temperatures. In case of operations at temperatures above the

boiling point of the solvent, a high-pressure reactor is used. At pres-

ent, use of bacteria to assist leaching is also in practice. Once the

metal value is solubilized, the leach liquor is filtered, purified, thick-

ened, and processed for winning the metal. The residue, if any, is

either rejected or processed further for recovering additional values.

The normal metal-winning operation involves any one or more of the

following processes: crystallization, adsorption, ionic precipitation, pre-

cipitation by either metals or gases, ion exchange, solvent extraction, or

electrolysis onto either solid cathode or mercury. Recent developments in

corrosion-resistant materials and the purification processes, like ion

exchange and solvent extraction, have made hydrometallurgy more ver-

satile and competitive. Table 1 lists some of the hydrometallurgical

processes, many of which are of commercial significance in different parts

of the world.

Most of the operations involved in the process routes listed in Table 1

are heterogenous in nature, involving at least solid and liquid phases.

Quite often the gaseous phase is also involved. The process steps may be

chemical or electrochemical in nature, which differs from each other both

from the thermodynamic as well as the kinetic considerations; however,

the present paper concentrates on the electrochemical aspects in hydro-

metallurgical operations. Many of such operations are electrochemical in

nature as indicated in the examples of Table 1, where the electrochemical

nature of some or all the process steps was already established. The

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Table 1. Hydrometallurgy used in the metal extraction processes

Serial

number

Metal

produced

Raw

material Operations

Whether

electrochemistry

involved in

some steps

1 Gold Native ore Leaching, precipitation Yes

2 Silver Native ore -do- Yes

3 Platinum Native ore,

anode slime

-do- and calcination —

4 Copper Native ore,

oxide ore,

sulphide ore

Leaching, SX-EW Yes

5 Nickel Native ore,

laterite

(oxide)

sulphide

Leaching, SX-EW, direct

electrowinning of metal

Yes

6 Aluminium Oxide ore Leaching-fused

salt electrolysis

Yes

7 Tin Oxide ore Pressure leaching,

amalgamation

Yes

8 Iron Low grade

oxide and

sulphide ore

Acid leaching (þO2)

also bacterial leaching

Yes

9 Manganese Oxide Leaching (reductive)

precipitation=EW

Yes

10 Zinc Calcine

sulphide

Acid leaching-EW

Pressure leaching-EW

Yes

11 Titanium Complex

oxide

Acid leaching

(purification)

—

12 Tungsten -do- Acid leaching=alkali

leaching=roast leaching

—

13 Uranium Oxide Acid leaching with oxide=

alkali carbonate leaching,

bacterial leaching.

Yes

14 Molybdenum Sulphide NAOCl leaching Yes

15 Lead Sulphide Leaching-EW Yes

‘—’ implies ‘not yet established’.

HYDROMETALLURGICAL PROCESSES 67

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following sections of this paper present theoretical considerations

involved in these steps.

THEORETICAL CONSIDERATIONS

The electrochemical nature of a process is determined by the character of

the solid and the liquid associated with it, and involves a solid solution

interface, as is well understood in electrochemical literature. Some solid

ore bodies (e.g., native ores of gold, silver, platinum, and copper) are

good conductors. Most sulphides, which constitute primary sources for

many metals (e.g., chalcopyrite, chalcocite, and covellite for copper;

galena for lead; sphalerite for zinc), are semiconductors (Table 2)

(Habashi 1970; Koch 1975; Osseo-Asare 1992). Some oxides are also

semiconductors (e.g., nickel oxide, cassiterite [SnO2], hematite [Fe2O3],

magnetite [Fe3O4], and pyrolusite [MnO2]). They provide the necessary

solid surface for electrochemical interaction during leaching. The leach-

ant, which usually is either a reductant or an oxidant, acts as the elec-

trolyte. Table 3 provides the list of oxidants used in hydrometallurgical

operations (Peters 1992). In the winning step, the emerging product

usually is a metal, which generally is a good conductor. Thus, hydro-

metallurgy provides suitable conditions for the operation of electro-

chemical systems. Sometimes external electric current is supplied, as in

the case of direct electrowinning of metals (Serial 5, Table 1), electro-

winning or electrorefining of metals (Serials 4�6, 9, 10, 15, Table 1), and

amalgamation (Serial 7, Table 1); however, this paper is confined to those

cases in Table 1 where electrochemical phenomena operate in situ with-

out any external electric current. These systems operate in a way similar

to metal corrosion systems familiar to any electrochemist. As will be

shown, many of these processes are examples of corrosion coupling,

galvanic coupling, etc., and follow the principles of metal corrosion.

Leaching as a Process Analogous to Corrosion

Leaching of conducting or semiconducting metals and minerals follows

electrochemical principles and can, therefore, be explained in terms of

electron transfer. The conducting solid surface in contact with an elec-

trolyte provides the solid=solution interface for electron transfer to take

place. In certain cases, presence of more than one such solid in contact

with the electrolyte may present a pathway for electron transfer through

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the solid-solid interface while the possibility of surface passivation or

catalytic activation in some cases is also not ruled out. All of these

processes are analogous to the well-established corrosion reaction for

metal systems. A corroding metal electrode dipped in an electrolyte

provides two half-cell reactions, namely, the anodic and cathodic, which

balance each other electronically so that there is no net charge transfer;

however, there is net mass transfer since the reactions, although different,

are not exactly the reverse of each other. Similarly, the leaching of

semiconducting minerals can be explained by applying the established

principles for metal corrosion involving charge transfer, though the

resistivity of the semiconductors is a critical factor. The resistivity values,

Table 2. List of Important Sulphide Minerals (Habayashi 1970; Koch 1975; Osseo-Asare

1992)

Metal Mineral Formula

Resistivity,

O m

Type of

semi-conductor

Energy

Gap, eV

Antimony Stibnite Sb2S3 — — —

Arsenic Realgar As4S4 — — —

Orpiment As2S3 — — —

Arsenopyrite FeAsS 3-570 — —

Bismuth Bismuthinite Bi2S3 — — —

Cadmium Greenockite=

Hewleyite

CdS — n 2.42

Cobalt Linnaeite Co3S4 — — —

Copper Chalcocite Cu2S 1072�1075 p 1.10

Covellite CuS 0.3-83� 1076 Metallic, p —

Digenite Cu9S5 — — —

Bornite Cu5FeS4 1.6-6000� 1076 — —

Chalcopyrite CuFeS2 150-9000� 1076 n 0.6*

Iron Pyrite FeS2 1.2-600� 1073 p, n, & p-n junction 1.2

Pyrrhotite FeS 2-160� 1076 p Very low

Lead Galena PbS 6.8� 1076 n,p 0.37

Mercury Cinnabar HgS — n 2.00

Manganese Haverite MnS2 10�20 — —

Molybdenum Molybdenite MoS2 — — —

Nickel Millerite NiS 2-4� 1077 — —

Pentlandite (Fe,Ni)S 1-11� 1076 — —

Silver Argentite Ag2S 1.5-2� 1073 n �1

Tin Harzenbergite SnS 3 p 1.08

Zinc Sphalerite ZnS 2.7� 1073�1.2� 104 n 3.67

*(Crundwell 1988).

HYDROMETALLURGICAL PROCESSES 69

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semiconducting type, and energy gap of some of the sulphide minerals are

given in Table 2. It can be seen that these values vary widely even for each

of the sulphides. This may be due to minor differences in their compo-

sition and the presence of hairline cracks. Koch (1975) pointed out that

some of the sulphides, like nickel, copper, iron, and lead, have resistivities

comparable to those of metals, and therefore can be used as massive

electrodes so that the rates of dissolution at current densities up to about

100 mA=cm2 can be measured by the standard electrochemical tech-

niques. Cominco demonstrated the commercial applications of this

aspect by directly electrowinning nickel using massive nickel matte elec-

trodes (Habashi 1971). In this case, the oxidant is the direct current, and

the direct electrowinning process can be explained in terms of electron

transfer; however, in many of the leaching processes involving semi-

conducting minerals also, the reactions can be explained in terms of

electron transfer. In the case of zinc, cadmium, manganese, and bismuth

sulphides, their conductivity is good enough to use the techniques

developed in semiconductor electrochemistry. Recently, Osseo-Assare

(1992) discussed the importance of semiconductor electrochemistry in

mineral leaching highlighting the significance of hole transfer, and this

aspect will be briefly discussed at a later stage. The following section deals

with only the electron transfer phenomena.

Mixed Potential Concept

When a metal or mineral with sufficient conductivity is dipped in an

electrolyte, a steady-state (corrosion) potential is developed due to the

Table 3. Redox potentials of hydrometallurgical oxidants (Peters 1992) on the hydrogen

scale (V)

Oxidant Redox equation Eho (pH¼ 0) Eh

o (pH¼ 10)

Fe3þ Fe3þþ e7¼Fe2þ 0.77 —

Fe(CN)637þ e7¼Fe(CN)6

47 — 0.46

O2(g) O2þ 4Hþþ 4e7¼ 2H2O 1.23 0.64

HNO3 NO37þ 4Hþþ 3e7¼NO(g)þ 2H2O 0.957 (0.17)*

(HNO2) NO27þ 2Hþþ e7¼NO(g)þH2O 1.202 (0.02)*

(NOþ) NOþþ e7¼NO(g) 1.45 —

Cl2(g) Cl2(g)þ 2e7¼ 2Cl7 1.358 (1.36)*

(ClO7) ClO7þ 2Hþþ 2e¼Cl7þHO 1.63 1.126

*Calculated for standard conditions. These potentials are not realizable because of side

reactions.

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two partial half-cell reactions and is termed as the mixed potential, Em.

This system is called a corrosion cell, and the probable reactions may be

represented in a simple form as follows:

Anodic:

Me ¼ Me2þ þ 2e� ð1Þ

MeS ¼ Me2þ þ S þ 2 e�; ð2Þ

Cathodic (in acid):

O2 þ 4 Hþ þ 4 e� ¼ 2H2O; ð3Þ

MeOn þ 2nHþ þ 2 e� ¼ Me2þ þ nH2O; ð4Þ

Anodic and cathodic (in presence of a redox couple Rn=Rn�1):

Rðn�1Þþ ¼ Rnþ þ e�; ð5Þ

Rnþ þ e� ¼ Rðn�1Þþ: ð6Þ

A convenient way to deal with such a system is to plot the current-

potential curves for the respective half-cells involved in the leaching

reactions. Superposition of these two plots on one another gives the

mixed potential, Em, and the mixed current, im. A conceptual diagram is

provided in Figure 1.

As seen in the figure, each polarization plot has three distinct seg-

ments: 1, 2, and 3 (A: anodic, C: cathodic). Segment 1 represents the

equilibrium (steady state) region where the particular electrode reaction is

mostly reversible and the potential obeys Nernst’s relation. Kinetics, in

this region, depend on conduction through the electrode and the elec-

trolyte. Stage 2 is known as the Tafel region, where the potential is

directly dependent on the logarithm of the current and the process is

under activation control. Stage 3 represents the limiting current region

where diffusion of either the reactant or the product species controls the

reaction. Partial and general kinetic expressions can be derived from the

position of the point of intersection with respect to the stage of each half

reaction. Earlier workers (Warren, Kim, and Heinen 1987; Rath,

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Paramguru, and Jena 1988) assumed Stage 2 for each half of the reaction

to derive the kinetic expressions, but Wadsworth (1984) derived separate

equations for all three situations. The present authors (Paramguru and

Ray 1996) derived equations with nine different combinations. Table 4

presents expressions for mixed potential, Em, and mixed current, im, at

different situations for reactions (2) and (6), and Table 5 provides

expressions for reactions (4) (Me is Mn, n¼ 2) and (5) (n¼ 3).

The reliability of quantitative application of polarization data may

sometimes be questioned, especially when the purity and surface area of

the semiconducting electrodes are uncertain. Paramguru and Ray (1996)

have suggested a practical approach for studying the kinetics and

mechanism of the process using the polarization data. It involves the

approximate identification of the stages (1, 2, or 3) of each half plot at

the point of intersection, implying that only qualitative significance is

attached to the polarization data. Expressions are then derived theoreti-

cally for Em and im as presented in the Tables 4 and 5. These expressions

are validated using the data from leaching studies. The methodology has

been successful in the study of reactions falling under situations 2�8 of

Figure 1. Hypothetical polarization curves for anodic and cathodic reactions superimposed

on one another.

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Tables 4 and 5 (Paramguru and Ray 1996). These include reactions

under mixed control (Tafel-Tafel interaction) (Paramguru and Ray

1995), dissolution by diffusion-controlled half-cell reactions (Paramguru

and Ray 1995; Dutrizac 1985), leaching by activation-controlled half-cell

reactions (Warren et al. 1987; Paramguru 1995; Jin, Warren, and Heinen

1993), and significantly, the dissolution by cyclic action of a redox couple

(Paramguru and Kanungo 1995; Nayak, Parida, Rao, and Paramguru

1995). A similar analogy can also be extended to galvanic coupling,

cyclic action of a redox couple to dissolve two semiconducting minerals,

and other hydrometallurgical operations, such as cementation and

precipitation under reducing conditions, including high-pressure opera-

tions. Some of these aspects will now be discussed under the following

headings:

(i) Corrosion coupling

(ii) Galvanic coupling

(iii) Electrochemical aspects in high pressure operation.

In addition, semiconducting aspects will also be discussed briefly at

the end.

Table 4. Expressions for mixed potential, Em, and mixed current, im, at different

situations for reactions (2) and (6)

Sl. no. Situation Expression for Em Expression for im

1 A1-Cl (2.3RT=F) log {(K2[Me2þ]

þK4[Rnþ])=(K1þK3[R

(n71)þ])}

K7[Rnþ]={(K1þK3[R(n71)þ])

(K2[Me2þ]þK4[Rnþ1])}1=2

2 A2-C1 (2.3RT=F) log {K4[Rnþ])=

(K1þK3[R(n71)þ])}

FK1{K4[Rnþ]=

(K1þK3[R(n71)þ])}1=2

3 A3-Cl �Ec¼Ecoþ(2.3RT=

ZcF) log {[Rnþ]=[R(n71)þ]}

FK5[Me2þ]

4 A1-C2 (2.3RT=F) log {(K2[Me2þ]

þ K4[Rnþ])=K1}

FK4[Rnþ](K1)

1=2=

(K2[Me2þ]þK4[Rnþ1])}1=2

5 A2-C2 (2.3RT=F) log {K4[Rnþ]=K1} FK1{K4[R

nþ]=K1}1=2

6 A3-C2 (2� 2.3RT=F) log {K4[Rnþ])=

K5[Me2þ]}

FK5[Me2þ]

7 Al-C3 �Ea¼Eaoþ(2.3RT=ZaF)

log [Me2þ]}

FK6[R*]

8 A2-C3 (2� 2.3RT=F) log {K6[R*]=K1} FK6[R*]

9 A3-C3 — —

K1K7 are constants, and *may be nþ or (n71)þ.

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CORROSIONCOUPLING

In hydrometallurgy there are several examples of corrosion coupling.

Fathi Habashi (1970) has chosen the 18th century process of gold

leaching in cyanide bath to demonstrate the corrosion coupling. Li et al.

(1992) have cited three classic examples, such as copper=iron cementa-

tion, silver dissolution in cyanide, and leaching of chalcocite in ferric

chloride solution involving corrosion coupling. The present author has

chosen the dissolution of base metal sulphides in ferric chloride, cyclic

action of a redox couple to dissolve two minerals, cementation of copper

by zinc, and electroless deposition of copper as examples for a detailed

discussion on this aspect.

Dissolution of BaseMetal Sulphides in Ferric Chloride

Galena, sphalerite, and chalcopyrite are a few base metal sulphide

minerals that have commercial significance. The CYMET Process

Table 5. Expressions for mixed potential, Em and mixed current, im, at different

situations for reactions (4) and (5)

Sl.

no. Situation Expression for Em Expression for im

1 A01-C0l (2.3RT=F) log{(C2[Fe3þ]

þC4[Hþ])=(C1[Fe2þ]

þC3[Me2þ])}

C7[Hþ][Fe2þ]={(C1[Fe2þ]{(C1[Fe2þ]

þC3[Me2þ])1=2(C2[Fe3þ]þC4[Hþ])}1=2

2 A02-C01 (2.3RT=F)log{C4[Hþ])=

(C1[Fe2þ]þC3[Me2þ])}

FC1{C4[Hþ]=(C1[Fe2þ]

þC3[Me2þ])}1=2

3 A03-C0l �Ec¼Ecoþ(2.3RT=ZcF)

log{[Hþ]2=[Me2þ]}

FC5[Fe*]

4 A01-C02 (2.3RT=F)log{(C2[Fe3þ]

þC4[Hþ])=C1[Fe2þ]

FC4[Hþ]{(C1[Fe2þ])=(C2[Fe3þ]

þC4[Hþ])}1=2

5 A02-C02 (2.3RT=F)log {C4[Hþ]=

K1[Fe2þ]}

F{C1C4[Fe2þ][Hþ]}1=2

6 A03-C02 (2� 2.3RT=F) log{C4[Hþ])=

C5[Fe*]}

FC5[Fe*]

7 A0l-C03 �Ea¼Eaoþ(2.3RT=ZaF)

{[Fe3þ]=[Fe2þ]}

FC6[Me2þ]

8 A02-C03 (2� 2.3RT=F) log{C6[Me2þ]=

C1[Fe2þ]

FC6[Me2þ]

9 A03-C03 — —

C1�C7 are constants corresponding to reactions (4)�(5), identical to K1�K7 in Table 4;

*may be nþ or (n71)þ.

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(McNamara, Ahrens, and Franek 1980), the CLEAR-Duval Process

(Atwood and Livinston 1980), the ZINCEX Process (Nogueira,

Regife, and Vigas 1982), the ELKEM Process (Barbery, Fletcher, and

Sirois 1980), the MINEMET Recherches Process (Demarthe, Gandon,

and Georgeaux 1976), and the USBM Process (Wong, Haver, and

Sandberg 1980) are a few processes that attempt to solubilize the

metal values using a redox system like Fe3þ=Fe2þ or Cu2þ=Cuþ or

both, and remain the classic examples of use of corrosion coupling

principles. Many investigations have been carried out on the elec-

trochemical character of these leaching systems (Wadsworth 1984;

Dutrizac 1985; Warren et al. 1987; Rath et al. 1988; Jin et al. 1993;

Paramguru and Ray 1995; Paramguru 1995; Paramguru and Kanungo

1995; Nayak et al. 1995; Paramguru and Ray 1996). Figure 2 pre-

sents a conceptual diagram of such corrosion coupling, and Figures 3

and 4 present polarization plots for galena (Paramguru and Ray

1995) and sphalerite (Paramguru 1995) leaching in ferric chloride.

Leaching of galena in ferric chloride has been studied in detail, and

Kobayashi, Dutrizac, and Toguri (1990) have published an extensive

review on this subject. The leaching mechanism in dilute ferric

chloride solutions is different from that in more concentrated media.

The former conforms to linear kinetics, the latter to diffusion

kinetics. The polarization data (Paramguru and Ray 1995) presented

in Figure 3 provides explanation for this difference. Plots (b) show

cathodic polarization plots for a range of ferric ion concentrations

while plot (a) presents the anodic one for the galena electrode. In

relatively dilute solutions, Em lies on the Tafel regions of both the

anodic and cathodic half reactions, conforming to linear kinetics as

described by case 5 of Table 4. Replacing Rnþ with Fe3þ, the fol-

lowing equations for Em and Im are obtained:

Em ¼ ð2:3 RT=FÞ log fK4½Fe3þ�=K1g; ð7Þ

im ¼ FK1fK4½Fe3þ�=K1g1=2: ð8Þ

These equations have been verified through polarization measure-

ments (Paramguru and Ray 1995) for lower ferric ion concentrations.

The leaching rate has also been found to follow equation (8) (Rath et al.

1988).

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At higher ferric ion concentrations, the cathodic plots (Fig. 3)

intersect the anodic curve on its limiting current region, signifying dif-

fusion kinetics. Situation 6 of Table 4 should describe Em and im, which

indicates a slope of (4.6 RT=F) or 0.118 V at 298 K when Em is plotted

Figure 2. Conceptual diagram for a mineral sulphide solution corrosion coupling.

Figure 3. Concurrent polarization curves at a ramp rate of 1 mA=s. (a) anodic galena in

(1 M NaCl þ 0.1 N HCl); (b) cathodic platinum in different fenic ion concentrations in

(1 M NaCl þ 0.1 N HCl þ 0.001 M Fe2þ); (1) 1.2, (2) 0.4, (3) 0.12, (4) 0.04, (5) 0.012,

(6) 0.004, (7) 0.0012, and (8) 0.00012 M Fe3þ.

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against log [Fe3þ], and im should be independent of [Fe3þ] but should

depend on [Pb2þ]. This dependence of Em and im on [Fe3þ] has been

verified from polarization measurements (Paramguru and Ray 1995).

Leaching studies reported by Dutrizac (1985) also support these

observations (i.e., the rate follows parabolic kinetics and depends on

[Pb2þ] and [Cl7], but independent of [Fe3þ] and [Fe2þ]). These results

stand as illustrative examples of a leaching reaction following corrosion

coupling principles.

Leaching of ZnS in FeCl3 is another attempt of this kind. Earlier

studies (Warren et al. 1987, Rath et al. 1988) indicated that the dis-

solution reaction was controlled by electrochemical surface reaction.

The leaching rate was proportional to [Fe3þ]1=2 and [Cl7] at lower

concentrations, but the rate became insensitive to the concentration of

these ions at higher concentrations. Addition of large amounts of Fe2þ

ion rather retarded the reaction. Warren et al. proposed an electro-

chemical model incorporating a charge transfer process for each ion

and an adsorption step for Fe3þ and Cl7 ions to explain these results.

Polarization measurements reported recently (Jin et al. 1993, Param-

Figure 4. Concurrent polarization curves at a ramp rate of 1 mA=s. (a) anodic sphalerite in

(1 M NaCl þ 0.1 N HCl); (b) cathodic platinum in different fenic ion concentrations in (1 M

NaCl þ 0.1 N HCl þ 0.001 M Fe2þ); (1) 1.2, (2) 0.4, (3) 0.12, (4) 0.04, (5) 0.012, (6) 0.004, (7)

0.0012, and (8) 0.00012 M Fe2þ.

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guru 1995) indicate that situation 2 of Table 4 is followed in this case,

and the expressions for Em and im are given under (Fe2þ in place of

R(n71)þ):

Em ¼ ð2:3 RT=FÞ log fK4½Fe3þ�=ðK1 þ K3½Fe2þ�g; ð9Þ

im ¼ FK1 fK4½Fe3þ�=ðK1 þ K3½Fe2þ�Þg1=2: ð10Þ

These expressions are well supported by the experimental results. If

[Fe2þ] is much less than K1, then im is directly dependent on [Fe3þ]1=2;

when [Fe2þ] is much larger, then it affects im. As concentration of [Fe3þ]

decreases, the process deviates from situation 2 to 5 of Table 4 and the

[Fe2þ] term vanishes from equations (9) and (10).

Cyclic Action of a Redox Couple

Earlier studies (Nayak et al. 1995; Paramguru and Kanungo 1995;

Paramguru and Ray 1996) revealed that reduction leaching of MnO2 in

the presence of FeS2 in acidic solution is accomplished by the simulta-

neous action of the two corrosion couples, FeS2=Fe3þ and MnO2=Fe2þ.

The slower of the two controls the overall rate, and the other is main-

tained at that rate by adjustment of [Fe3þ] or [Fe2þ], whichever is

involved in that couple.

Recent studies (Paramguru and Kanungo 1998a; Paramguru and

Kanungo 1998b; Paramguru, Mishra, and Kanungo 1998) provide more

insight into this mechanism. The dissolution mechanism may not always

fall under the nine situations described in Tables 4 and 5. The mixed

potential Em, depending at times on the solution conditions, may lie on

linear-Tafel transition region of one of the polarization plots or below the

self-corrosion point on the current axis. In both cases the kinetic equa-

tions given in Tables 4 and 5 do not hold good. For dissolution of

manganese nodule in the presence of FeS2 in dilute HCl medium

(Paramguru and Kanungo 1998b), the following situations have been

observed:

(i) In presence of sufficient acid, dissolution of FeS2=Fe3þ couple runs at

a slower rate. The following two kinetic equations give close fit to the

experimental leaching data:

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(a) Em on anodic linear-Tafel transition and cathodic Tafel:

½Fe3þ�=ðRateÞ2 ¼ K8ð½Fe2þ�½Hþ�½SO42��=½Fe3þ�Þ þ K9: ð11Þ

(b) Em on anodic Tafel and cathodic linear-Tafel transition region:

½Fe3þ�=ðRateÞ2 ¼ K10½Fe2þ� þ K11: ð12Þ

When [Fe3þ]=[Fe2þ] is greater than 50, the first situation may exist

(Paramguru and Kanungo 1995, Paramguru and Kanungo 1998b),

otherwise, the second one is valid.

(ii) In presence of insufficient acid, the following two kinetic equations

hold good:

ðRateÞ ¼ K12ð½Fe2þ�½Hþ�½SO42��Þ; ð13Þ

ðRateÞ ¼ K13ð½Hþ�½Fe2þ�Þ1=2: ð14Þ

These equations refer to self-corrosion of FeS2 and Tafel-Tafel interac-

tion of MnO2=Fe2þ couple, respectively. Owing to the low acid con-

centrations, the MnO2=Fe2þ rate came down to a level of rate-controlling

stage. Depletion of Fe3þ through precipitation as hydroxide or basic

sulphate due to low acid may push the im of FeS2=Fe3þ couple to a value

lower than the self-corrosion current of this couple. There is a possibility

that both couples operate at the same rate. Figure 5 provides the proof

for these findings (i.e., ba (¼ im of FeS2=Fe3þ) is less than bc (¼ im of

MnO2=Fe2þ) at higher acid level, and ac is smaller than aa (ac and aa refer

to insufficient acid).

Similar studies have also been made on the MnO2-FeS2 system in

H2SO4 (Nayk, Mishra, and Paramguru 1999).

Dissolution of ZnS in presence of MnO2 has also been explained

recently on these lines (Rao and Paramguru 1996). Pande, Gupta, and

Altekar (1982) reported an attractive method of acid processing of

sphalerite concentrates in the presence of manganese dioxide with the

dissolution reaction:

ZnS þ MnO2 þ 2H2SO4 ¼ ZnSO4 þ MnSO4 þ S0 þ 2H2O: ð15Þ

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Neither ZnS nor MnO2 dissolves as such in dilute H2SO4; however,

they dissolve when placed together in the acid by the action of two

corrosion couples, namely, ZnS=Fe3þ and MnO2=Fe2þ. Equations (2)

and (6) constitute the first couple by replacing Me, R, and n, respectively,

with Zn, Fe, and 3. MnO2=Fe2þ couple is represented by the reactions (4)

replacing Me with Mn and n with 2, and (5) replacing R with Fe and

n with 3.

These two couples run simultaneously to cause the dissolution of

both ZnS and MnO2 (Rao and Paramguru 1996, Rao and Paramguru

1998). The couples balance each other through adjustment of Fe3þ and

Fe2þ concentrations. The couples may operate as per the situation 5 of

Figure 5. Concurrent polarization curves for anodic FeS2 and cathodic MnO2 at a ramp

rate of 1 mA=s at different HCl concentrations. (a) 0.08 M, (b) 0.25 M, and (c) 1.0 M. Also

present for platinum electrode in electrolyte: Fe3þ: 0.05 M, Fe2þ: 0.001 M.

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Tables 4 and 5, and both play a role in the rate-controlling process.

Equation (16), indicating a constant ratio between [Fe3þ] and ([Fe2þ]

[Hþ]), describes the dissolution process:

ð½Fe3þ�=½Fe2þ�½Hþ�Þ ¼ K14: ð16Þ

Cementation of Cu2+ on Zinc

Zinc dust cementation for impurity (Cu2þ, Co2þ, and Cd2þ) removal

from the zinc electrolyte is one classic example of the cementation reac-

tion. It has been established in Power and Ritchie (1976) that for

cementation of Cu2þ, the mixed potential lies on the Tafel region of the

anodic (Zn) and the diffusion segment of the cathodic Cu2þ plot.

Therefore, the reaction is diffusion controlled with limited dependence on

temperature. The rate follows a first-order process with respect to [Cu2þ].

The initial rate is followed by a second stage with an enhanced rate. Zinc

dissolution via hydrogen reduction also takes place simultaneously.

Recently, studies by Mishra and Paramguru (2000), dealing indepen-

dently with the anodic (Zn=Zn2þ) and cathodic (Cu2þ=Cu) half-cell

reactions in a dual cell have indicated that the protonation reaction is

catalyzed by the deposited copper on the zinc surface. Expressions for Em

and im were derived for the following reactions:

ZnðSÞ ¼ Zn2þ þ 2e�; ð17aÞ

Cu2þ þ 2e� ¼ CuðSÞ; ð17cÞ

ZnðSÞ þ Cu2þ ¼ Zn2þ þ CuðSÞ; ð17Þ

Em ¼ fð2:303 � 2RTÞ=Fg logðK15½Cu2þ�Þ; ð18Þ

im ¼ K16½Cu2þ�; ð19Þ

ZnðSÞ ¼ Zn2þ þ 2e�; ð17aÞ

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2½Hþ þ e� ¼ 1=2 H2�; ð20cÞ

Zn þ 2Hþ ¼ Zn2þ þ H2; ð20Þ

Em ¼ fð2:303 RTÞ=Fg logðK17½Hþ�Þ; ð21Þ

im ¼ K18½Hþ�1=2: ð22Þ

In case of reaction (20), Em was found to lie on the Tafel regions of

both the anodic and cathodic half polarization plots. The practical

cementation data was found to be the result of these two reactions (17

and 20), i.e., the experimental slope (of Em versus log [Hþ]) of 0.105 V per

decade is in between 0.059 (eqn. 21) and 0.118 V per decade (eqn. 18).

Electroless Copper Deposition

Electroless metal deposition onto a substrate is a good example of pre-

cipitation under reducing conditions. This system can be well understood

using polarization studies, as the solid involved is a good conductor and

does not pose measurement problems. Recently, electroless deposition of

copper onto a copper substrate was explained through polarization stu-

dies (Mishra and Paramguru 1996, 1997a, 1997b, 1999). The following

reactions were mentioned:

Anodic: 2HCHO þ 4OH� ¼ 2HCOO� þ H2 þ 2H2O þ 2e�; ð23aÞ

Cathodic: CuðLÞ2 þ 2e� ¼ Cu0 þ 2L�; ð23cÞ

Overall: CuðLÞ2 þ 2HCHO þ 4OH�

¼ Cu0 þ 2HCOO� þ H2 þ 2H2O þ 2L�: ð23Þ

Here L represents the complexing agent.

Kinetic experessions can be derived on the basis of the position of Em

on the individual polarization plots, and these can be validated with

reference to actual deposition rate. Table 6 gives four such cases for

which Em and im have been derived and experimentally verified (Mishra

and Paramguru 1996, 1997a, 1997b, 1999). Deposition rates reported by

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Ohno (1991) under different conditions were also found to support this

view (Fig. 6). Many other precipitation processes used in hydro-

metallurgy are of this nature.

GALVANICCOUPLING

As galvanic interaction is known to take place during leaching operations

involving metals and semiconductors, leaching of sulphide minerals

always involves some galvanic coupling since these minerals are semi-

conductors and contain other sulphide minerals as impurities. Replacing

Me in reaction (2) with any two metals provides the two reactions of the

galvanic couple, and the one with higher potential acts as the cathode. On

galvanic contact, dissolution of the cathodic mineral gets retarded, while

that of the anodic one is enhanced. Wadsworth (1984) has brought out the

importance of the concept of mixed potential to describe the trend of

galvanic interactions. Using the Butler-Volmer equation, he has expressed

the rate of anodic dissolution of the active electrode as a function of

various parameters. Yelloji Rao and Natarajan (1986) measured the

combination potentials and galvanic currents for different sulphide

mineral couples and established the order of galvanic activity. Nowak,

Krause, and Pomianowski (1984) used a small-scale amplitude cyclic

voltametry to study the galvanic behavior of minerals and developed

equations in terms of the rest and mixed potentials along with Tafel slopes

to describe galvanic interaction. Holmes and Crundwell (1995) used a

voltage balance over the galvanic couple as basis for a mathematical

Table 6. Different mechanisms of electroless copper deposition based on position of Em

on specific regions of polarization plots of anodic and cathodic half reactions

Polarization regions

in which Em lies

Solution condition

Mechanism [HCHO] [Cu2þ]

(1) Cathodic linear and

anodic limiting current

Anodic diffusion control Very low —

(2) Cathodic Tafel and

anodic limiting current

Anodic diffusion Low Moderate

(3) Cathodic and anodic

Tafel

Activation control Low to moderate Modrate to

high

(4) Cathodic-limiting current

and anodic Tafel

Cathodic diffusion

control

High Low to

moderate

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description of the magnitude of the galvanic interaction. The present

author used compact mineral electrodes in a dual cell to generate polari-

zation plots for the partial anodic and cathodic parts of the galvanic

couple, which not only measured galvanic interaction quantitatively, but

also provided the mechanism of interaction (Rao, Paramguru, Das, and

Ray 1992; Paramguru 1992). One interesting finding (Nayak, Parida, Rao,

Sahoo, and Paramguru 1994; Rath and Paramguru 1994; Paramguru and

Nayak 1996) has been the galvanic interaction between sulphide minerals

with MnO2. Here both the minerals corrode, MnO2 as cathode and MeS

as anode, and the potential difference between the electrodes is much

larger than that of the sulphide-sulphide galvanic couples and results in

larger galvanic currents. The couple may be better understood following

an approach similar to that of corrosion coupling. The two partial

polarization plots can be obtained independently to get Em and im (in this

case, Eg, galvanic potential, and ig, galvanic current) on superimposition.

Equations (4) (replacing Me with Mn and n with 2) and (2) provide the

cathodic and anodic reactions, respectively. Figure 7 presents the Evans

diagram for MeS-MnO2 galvanic couples and shows the interaction in the

Tafel region of both the anodic and cathodic plots for FeS2-MnO2 and

Figure 6. Plots of log iplating versus log [Cu2þ] and log [HCHO] for results obtained from

Ref. (Ohno 1991). Hetched areas indicate regions of different mechanisms.

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CuFeS2-MnO2. For PbS-MnO2, the interaction is on the anodic limiting

current region. The ZnS-MnO2 couple is in between these two. Situation 5

of Table 4 should apply for the FeS2-MnO2 couple replacing equation (6)

with equation (4) and equation (2) with the following:

FeS2 þ 8H2O ¼ Fe2þ þ 2SO42� þ 16Hþ þ 14e�: ð24Þ

The relationships for Eg and ig are as given below:

Eg ¼ 2:3 RT/F log fK19½Hþ�g; ð25Þ

Figure 7. Concurrent anodic polarization curves of (1) PbS, (2) ZnS, (3) CuFeS2, and (4)

FeS2 against (5) cathodic plot of MnO2 in 0.1 M HCl at room temperature. Ramp rate: 1

mA=s. 1M NaCl incorporated in the bath for PbS electrode.

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ig ¼ K20½H�1=2: ð26Þ

These equations have been verified experimentally (Paramguru and

Nayak 1996). The same equations with different constants are valid for

Eg and ig of the other galvanic couple MnO2-CuFeS2, where the inter-

action takes place in the Tafel regions of both the plots. This has been

observed in case of this couple also (Madhuchhanda, Devi, Rao, Rath,

and Paramguru 2000). When the interaction takes place in the limiting

current region of the anodic curve, as in the case of the MnO2-PbS

couple, the expression for ig would not contain the [Hþ] term because the

anodic reaction of PbS, which is expected to be rate controlling, does

not contain Hþ term and hence no diffusion of Hþ. Thus, ig, instead of

depending on Hþ, may depend on Pb2þ ion concentration. Since the

anodic plot of zinc appears to be somewhere in between the Tafel and

limiting current regions at the point of intersection, the dependence of igon [Hþ] should be in between. In fact, the experimental dependence of igon [Hþ] for PbS-MnO2 and ZnS-MnO2 couples have been found to be

zero and around 0.2, respectively (Madhuchhanda et al. 2000).

It is interesting that pyrite-chalcopyrite-sphalerite and galena are

in this order in their declining noble character. Hence, their potential

difference in the galvanic cells with MnO2 increases in this order and also

the observed ig values. Thus, the polarization plots in the case of

the MnO2-PbS couple, with the highest potential difference, traverse

the longest path in the current axis and the plot for galena (with lower

corrosion current density) is already in the limiting current region by the

time it intersects the MnO2 plot. Eg being the farthest from EPbS (in

comparison to EMnO2), the process is under anodic control; however,

the case is different for the MnO2-FeS2 couple with lower potential

difference. The intersection takes place in the Tafel-Tafel regions and

the process is controlled by the cathodic reaction. CuFeS2-MnO2

and ZnS-MnO2 couples appear in between these two extremes

(Madhuchhanda et al. 2000).

ELECTROCHEMICAL ASPECTS INHIGH-PRESSUREOPERATIONS

High-pressure operations are a part of hydrometallurgical practices in

some cases and involve high-temperature and high-pressure aqueous

solutions. Thus, many of the electrochemical aspects, such as corrosion

coupling and galvanic coupling, are also involved in these operations.

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Though the same principles are extended to understand the phenomena,

with the advent of facilities to monitor electrochemical parameters of

aqueous solutions under high temperature and pressure electrochemical

aspects are now being studied in detail for high-pressure systems. The

present study considers the following cases for discussion:

(i) Aqueous oxidation of sulphide minerals with oxygen pressure: (a)

Ammoniacal pressure oxidative leaching; (b) acid pressure oxidative

leaching.

(ii) Pressure leaching of reduced ilmenite.

(iii) Pressure reduction of aqueous metals by hydrogen.

Aqueous Oxidation of SulphideMinerals with OxygenUnder Pressure

Gaseous oxygen can also act as an oxidant to solubilize sulphide

minerals. Reaction (3) describes the cathodic half of the reaction in acidic

conditions, and the following equation holds in alkaline medium:

O2 þ 2H2O þ 4e� ¼ 4OH�: ð27Þ

Unfortunately, the oxidation reaction of sulphide minerals with

oxygen or air proceeds too slowly to be of commercial significance;

however, with the advent of pressure technology, the aqueous oxidation

of sulphide minerals could be affected at a rate suitable for commercial

application. Pressure operation improves the kinetics of aqueous reaction

dramatically by the following means:

(i) Raising the reaction temperature above the boiling point of the sol-

vent or leachant

(ii) Increasing the partial pressure of the gaseous oxidant

Thus, pressure leaching has already been commercially practiced for

treating sulphide minerals in ammoniacal (Rosenzweig 1969; Kerfoot

1989; Veltman and Weir 1981) as well as acidic medium (Parker and

Romanchuk 1979; Parker 1981). Recently, Deng (1995) has published an

excellent review describing the development and application of pressure

oxidation technology, including the chemistry involved. Some aspects of

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the process may now be discussed with emphasis on electrochemical

operations.

Ammoniacal Pressure Oxidative Leaching. Sherrit Gorden was the first

commercial operator to employ the ammoniacal pressure leaching proc-

ess for extracting nickel directly from sulphide ores (Deng 1995). The

large commercial scale operation at Fort Saskatchewan, Alberta, has

successfully treated nickel-copper-cobalt sulphide minerals. The chem-

istry of this process has been described by Forward and Mackiw (1955).

Anacondas’ Arbitor Process (Kuhn, Arbiter, and Kling 1974) avoided

high partial pressure of oxygen by intense mixing of the slurry. Recently,

the Regional Research Laboratory, Bhubaneswar, India, reported

another approach to treat complex copper-zinc-lead sulphide mineral in

ammoniacal medium under oxygen pressure (Rao et al. 1984). A number

of studies have also been made to understand the kinetics (Beckstead and

Miller 1977; Das, Anand, and Rao 1984; Warren and Wadsworth 1984;

Anand, Rao, and Das 1985a, 1985b; Rao et al. 1992). Interesting studies

conducted on chalcopyrite indicate that the dissolution process is iden-

tical to that of the corrosion of metals. Beckstead and Miller have shown

that the ammonia oxidation leaching of chalcopyrite is controlled by a

catalytic electrochemical surface reaction. The anodic and cathodic half-

cell reactions are as follows:

Anodic:

CuFeS2 þ 19OH� ¼ Cu2þ þ 1=2 Fe2O3 þ 2SO42� þ 19=2 H2O þ 17e�;

ð28Þ

Cathodic:

O2 þ 2H2O þ Cu2þ þ 4e� ¼ 4OH� þ Cu2þ: ð29Þ

Equation (27) also runs along with equation (29) as the cathodic

reaction.

The authors have analyzed this system like a corrosion couple and

used the Butler-Volmer equation to derive the kinetic expression. Sub-

sequently, Warren and Wadsworth (1984) investigated the anodic oxi-

dation of CuFeS2 in ammoniacal solutions using potentiodynamic

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polarization and constant potential techniques to give further support to

this theory. They observed the formation of a copper-deficient sulphide

layer during the initial dissolution at low overpotentials according to the

following reaction:

ð1=xÞCuFeS2 þ 4NH3 ¼ ð1=xÞCu1�xFeS2 þ CuðNH3Þ42þ þ 2e�: ð30Þ

The catalytic effect of copper ion is on the cathodic half reaction. The

reaction model proposed by them was typical of the corrosion of some

metals. This reaction is a charge transfer reaction of first order with

respect to the [OH7] and is in agreement with that proposed by Beckstead

and Miller (1977).

Acid Pressure Oxidative Leaching. Acid pressure oxidative leaching is

generally used more extensively in commercial operations of (i) extraction

of nickel, cobalt, and copper from matte or sulphide minerals; (ii) direct

leaching of zinc sulphides; and (iii) pretreatment of refractory gold ores.

For Ni3S2 dissolution, the following reaction scheme has been reported

(Hofirek and Kerfoot 1992).

3Ni3S2 þ 4Hþ þ O2 ¼ Ni7S6 þ 2Ni2þ þ 2H2O; ð31Þ

Ni7S6 þ 2Hþ þ 0:5O2 ¼ 6NiS þ Ni2þ þ H2O: ð32Þ

During this process, a substantial quantity of ferrous iron is also

released into the solution. It is therefore assumed that the dissolved iron

acts as an electron carrier and enhances the leaching rate as per the

following equations:

Ni3S2 þ 2Fe3þ ¼ 2Fe2þ þ 2NiS þ Ni2þ; ð33Þ

2Fe2þ þ 2Hþ þ 0:5O2 ¼ H2O þ 2Fe3þ: ð34Þ

Thus, in the corrosion coupling type of reactions (2) and (3), pressure

operation provides a higher partial pressure of oxygen, which enhances

not only the dissolution reactions (31) and (32), but also the Fe2þ to Fe3þ

oxidation reaction (34). As a result, reaction (33) is also enhanced.

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Direct acid pressure leaching of ZnS is another recent development

(Deng 1995). At elevated temperature and oxygen pressure, sphalerite

dissolves according to the following reactions:

ZnS þ 2O2 ¼ ZnSO4; ð35Þ

2ZnS þ O2 þ 2H2SO4 ¼ 2ZnSO4 þ 2S0 þ 2H2O: ð36Þ

Both reactions proceed very slowly with chemically pure zinc sul-

phide (Bjorliy 1954), but the presence of dissolved iron can accelerate the

reaction significantly with the following leaching mechanism:

Fe2ðSO4Þ3 þ ZnS ¼ 2FeSO4 þ ZnSO4 þ S0; ð37Þ

4FeSO4 þ O2 þ 2H2SO4 ¼ 2Fe2ðSO4Þ3 þ 2H2O: ð38Þ

This is similar to the nickel sulphide dissolution (33) and (34) by the

action of Fe3þ=Fe2þ couple. Further, hole conduction might play a

prominent role in this dissolution process. This will be discussed in the

next section.

The pre-treatment of refractory gold ores mostly aims at oxidation of

pyritic sulphur to sulphate or elemental sulphur and release of ferrous ion

which oxidizes to ferric for hydrolysis to take place. The gold part is left

in the residue for cyanidation. The pyrite dissolution mechanism is

identical to that of the other sulphides discussed above.

Pressure Leaching of Reduced Ilmenite

One of the processes to treat ilmenite is to reduce it in a rotary kiln to

produce reduced ilmenite followed by the removal of metallic iron by

leaching in aerated ammonium chloride solution (Behr, Canning,

Goodheart, and Uusna 1965). The second step in this process is essen-

tially a redox reaction represented by equation (3) as the cathodic and the

following reaction as the anodic half:

Fe ¼ Fe2þ þ 2e�: ð39Þ

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Ammonium chloride aids this corrosion process by preventing the

precipitation of Fe2þ as hydroxide within the pores of the rutile grain

through the following equilibrium:

NH4þ þ OH� , NH3 þ H2O: ð40Þ

Recent studies such as Jayasekera, Marinovich, Avraamides, and

Bailey (1995), have considered electrochemical measurements to gain

understanding of the process. Evans diagrams based on the individual

anodic (eqn 39) and cathodic (eqn 3) half processes were constructed for

0.2 M NH4Cl solution in (a) saturated air at 80 �C and (b) with 300 kPa

po2 at 150 �C. The intersection occurred in the limiting current region of

the cathodic curve, implying that the reaction rate is controlled by the

diffusion of O2 to the iron surface. It was observed that the reaction

rate at 300 kPa po2 and 150�C was about 50 times faster than with air

under atmospheric pressure at 80 �C. At a constant temperature, the

corrosion rate, icorr, increased with increasing oxygen partial pressure,

and at a constant oxygen partial pressure, it increased with increasing

temperature. Preliminary pressure leaching trials indicated that the rate

of reaction increased significantly at elevated temperature and pressure.

Pressure Reduction of AqueousMetals by Hydrogen

Reduction of aqueous metal ions with hydrogen at elevated temperature

and pressure is described by the following reaction:

Me2þ þ H2 ¼ Me þ 2Hþ: ð41Þ

This electrochemical reaction proceeds with the following half-cell reac-

tions:

Anodic: H2 ¼ 2Hþ þ 2e�; ð41aÞ

Cathodic: Me2þ þ 2e� ¼ Me: ð41cÞ

Recently, Nagai and Sato (1978) conducted electrochemical investi-

gations on nickel in sulphate-acetate aqueous solutions treated with

hydrogen under pressures ranging from 0.5 to 2.0 Mpa. Individual anodic

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polarization curves for Ni2þ reduction were plotted along with those for

anodic oxidation of hydrogen. The intersection point provided the

reduction rate (exchange current of reaction). Wajszezuk and Charewicz

(1994a) carried out similar studies for the reduction of aqueous cobalt

and observed that the kinetic dependencies on [Co2þ], pH2, and tem-

perature obtained from the polarization measurements were identical to

those found from kinetic measurements (Wajszezuk and Charewicz

1994b). They also observed that the exchange current (reduction rate)

increased with [Co2þ] at a constant pH2 and temperature and increased

with temperature at a constant [Co2þ] and pH2.

MINERALDISSOLUTION INFLUENCEDBYHOLE TRANSFER

Hitherto, electrochemical aspects mostly based on electron transfer

principles have been covered; however, another important facet of

semiconductors is that they are characterized by a limited concentration

of charge carriers. Thus, there are interfacial reactions through hole

transfer. This aspect has been analyzed by Vaughan (1984); Crundwell

(1988); and (Osseo-Asare) 1992. Some salient features are briefly dis-

cussed here.

Semiconductors are marked by the presence of energy gaps between

the valence and conduction bands. When this gap is low, it is possible for

thermal agitation to promote an electron from the valence band to the

conduction band, thereby producing an electron hole in the valence band

(Osseo-Asare 1992). Such materials are known as intrinsic semi-

conductors. When a material acquires conduction band electrons and

valence band holes by receiving impurity atoms, it is called an extrinsic

semiconductor. It may be termed as n-type or p-type, depending on

whether the donor atom possesses extra or less electron in comparison to

those of the material. The importance of the semiconductor solid-state

properties in the electrochemistry of semiconductors is well known. In the

case of a metal, the application of the mixed potential theory to leaching

kinetics assumes that the potential across the solid side of the interface

(known as the space-charge region) remains constant so that the applied

potential difference appears across the Helmholtz layer. The dissolution

reaction is predicted to have a charge transfer coefficient of 0.5 and a

Tafel slope of 0.118 V=log(mA) for an ideal single-electron reaction;

however, in the case of a semiconductor-solution interface charge con-

centrations can be found in the space-charge region, the Helmholtz plane,

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and the Gouy layer. The net effect in the case of an ideal semiconductor is

that the current-voltage curve for a single-electron reaction results in a

charge-transfer coefficient of 1 and a Tafel slope of 0.059 V=log(mA).

Crundwell (1988) observed that sulphide electrodes do not display

classical semiconductor-solution behavior, owing to the formation of

surface states that aid the transfer of charge. The presence of surface

states, especially those due to zero valent sulphur, may give rise to beha-

vior which is intermediate between that of a metal and a semiconductor.

The bonding structure, nonstoichiometry present in the sulphides also

affect the dissolution process. These aspects are discussed below.

Holes and Sulphide Dissolution

In general, the position of the Fermi level with respect to the band

structure of semiconductor in the solid side and the energy level of the

redox couple in the solution side determine whether electron or hole

transfer would result. It has also been observed that, for a given polar-

ization, the higher the energy gap, the higher is the hole pathway relative

to the electron pathway. In semiconductors with a band gap greater than

1 eV, anodic dissolution occurs almost entirely as a result of hole injection

into the valence band (Gerischer and Mindt 1968). The oxidative

decomposition of a sulphide semiconductor can occur in the following

steps (Crundwell 1988).

ðMeSÞs þ X ¼ ðMeXÞþ þ �Ss þ e�; ð42aÞ

ðMeSÞs þ X þ hþ ¼ ðMeXÞþ þ �Ss; ð42bÞ

�Ss þ Y ¼ ðSYÞþ þ e�; ð42cÞ

�Ss þ Y þ hþ ¼ ðSYÞþ; ð42dÞ

�Ss þ �Ss ¼ (S-S)s; ð42eÞ

Sn þ �Ss ¼ Snþ1: ð42fÞ

Here, X, Y represents ionic species in the solution.

Dissolution of ZnS in ferric ion-containing solution and in acid

under pressure have been discussed in previous sections. In the case of the

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former, a half-order dependence of the rate on the concentration of

[Fe3þ] has been explained in terms of electron transfer as happens in

metal-like conduction. In the case of the latter, it has been indicated that

presence of lattice Fe in ZnS structure enhances dissolution. Crundwell

(1988) has explained these facts in terms of hole transfer process. Zinc

sulphide is a semiconductor, with a band gap of 3.6 eV and a high

resistivity; however, presence of 12.4 atomic % Fe showed a band gap of

only 0.49 eV with little change in the resistivity, as the substitution of iron

in sphalerite provides d-band energy levels within the band gap. Another

consequence of the iron impurity is that the Fermi level is pinned at the

energy level of the d-orbital. Hence, all the applied potential appears

across the Helmhoz layer. From fundamental principles based on hole

transfer, Crundwell derived an equation which indicated a first-order

dependence of rate on the occupied states in the d-orbital band (i.e., the

concentration of Fe impurity in the [Zn, Fe] S), and half order in the

concentration of oxidant. A stronger oxidant and ultraviolet irradiation

should enhance the dissolution rate. The latter has been proved by Exner,

Gerlach, and Pawlek (1969).

The dissolution of ZnS in the presence of Fe3þ ion has been

explained in terms of surface states created by the Fe3þ ion. It has been

proposed that, simultaneous with the Fe3þ reaction, a nonoxidative

reaction occurs, causing partial disruption of surface chemical bonds at

the ZnS surface, resulting in surface states with energy levels in the for-

bidden gap. These surface states are believed to be -SHd- groups formed

by the protonation of the ZnS surface (Tributsch and Bennett 1981;

Osseo-Asare 1992).

The nonstoichiometry of galena influences the semiconduction type.

Galena is a good conductor displaying n- and p-type semiconduction,

depending on whether sulphur is deficient or in excess. The band gap is

about 0.37 eV. A Tafel slope consistent with the semiconductor model is

observed, though this is influenced by the equilibrium with lead ions in

solution, and the applied potential is distributed across both the space-

charge layer and the Helmholtz layer (Crundwell 1988); however, when

the electrode potential is more in the anodic direction, the potential

difference appears across the Helmholtz layer and metal-like Tafel

behavior occurs (Richardson and O’Dell 1984).

The kinetics of chalcopyrite dissolution is complicated by the for-

mation of a surface film. The influence of redox couples is in accordance

with the semiconductor model, and a dissolution mechanism involving

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both electrons and holes is proposed (Crundwell 1988). In the case of

pyrite and molybdenite, however, the valence band is narrow and of

nonbonding character because of crystal-field splitting. Therefore, elec-

tron transfer with this band (hole injection) has no effect on the bonding

of the solid. Thus, these two minerals do not get oxidized easily.

Holes and Oxide Dissolution

Dissolution of oxides may be oxidative and reductive, as well as non-

oxidative. Irrespective of the type of reaction path, reaction rates are

influenced due to stoichiometry and thermal history of the solid, storage

environment, and nature and concentration of solid-state impurities

(Osseo-Asare 1992). Oxide dissolution may be best illustrated through an

example. Studies have been undertaken to understand the acidic dis-

solution of NiO (Nii 1970; Jones, Segall, Smarth, and Turner 1977, 1978;

Lussiez, Osseo-Asare, and Simkovich 1981; Furrer and Stumm 1986;

Stumm, Suzberger, and Sinniger 1990), which expectedly follows the

following nonoxidative reaction:

NiO þ 2Hþ ¼ Ni2þ þ H2O: ð43Þ

It has been observed that this reaction is affected by oxygen pressure

(Nii 1970), dopant type (Lussiez et al. 1981), annealing temperature

(Jones et al. 1977), and Co3þ (Jones et al. 1978). The first three factors,

in some way, influence the concentration of the holes. Nickel oxide is a

p-type semiconductor, where the defects are holes and cation vacancies.

The nonstoichiometry is a function of oxygen pressure; the higher the

oxygen pressure, the higher the amount of oxygen in the crystal lattice,

and therefore the higher the concentration of holes and cation vacancies

(Nii). Doping with lower valence cation, Liþ, increases the concentra-

tion of holes, whereas the higher valence cation, Cr3þ, has the opposite

effect (Lussiez et al.). Heat treatment like annealing is known to affect

the hole concentration (Jones et al. 1977). Stumm et al. (1986, 1990)

analyzed the adsorption phenomena in this dissolution process and

proposed the ligand-promoted dissolution mechanism. In light of this

mechanism, the effect of Co3þ on NiO can be explained in terms of hole

injection (Osseo-Asare 1992; Jones et al. 1978; Furrer and Stumm 1986;

Stumm et al. 1990).

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Thus it can be said that hole transfer plays an important role in

semiconducting mineral dissolution, which is gaining importance.

CONCLUSION

Many hydrometallurgical processes are electrochemical in nature. There

is an increasing recognition of application of electrochemical approaches

to study and understand them. As of today, the interfacial electro-

chemistry of mineral dissolution is treated in terms of electron transfer.

The hole transfer phenomenon, which is an inherent part of semi-

conductor mineral leaching, is now gaining its deserving importance.

More understanding of these aspects may pave the way for further

improvements in hydrometallurgical operations.

LISTOFSYMBOLSUSED

e7 : The electron

h : The electron hole

i : The current

ig : The galvanic current

im : The mixed current

n : No. of oxygen atoms

E : Electrode potential

Ea : Rest potential of the anode

Ec : Rest potential of the cathode

Eao : Standard equilibrium potential of the anodic reaction

Eco : Standard equilibrium potential of the cathodic reaction

Eg : Galvanic potential

Em : The mixed potential

EMnO2: The rest potential of MnO2 electrode

EPbS : The rest potential of PbS electrode

F : Faraday’s number

L : The complexing agent

Me : The metal

R : Universal gas constant

T : Absolute temperature

X,Y : The ionic species

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