A comparative study of leaching kinetics of limonitic laterite and

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Hydrometallurgy 72 (2004) 59–72

A comparative study of leaching kinetics of limonitic laterite and

synthetic iron oxides in sulfuric acid containing sulfur dioxide

G. Senanayake*, G.K. Das

A.J. Parker Cooperative Research Center for Hydrometallurgy, Department of Mineral Science and Extractive Metallurgy,

Murdoch University, Perth, WA 6150, Australia

Received 4 January 2002; received in revised form 2 September 2002; accepted 7 June 2003

Abstract

Limonitic laterite ore of particle size 90–125 Am containing goethite, magnetite and hematite was leached for 6 h at a pulp

density of 10% (wt/vol) in sulfuric acid in the absence or presence of sulfur dioxide at atmospheric pressure and 90 jC in a glass

reactor vessel. The sulfur dioxide flow rate was kept at 0.6 L min� 1 L� 1 of slurry to maintain a constant SO2 concentration of

c 0.3 mol L� 1 in solution, and the sulfuric acid concentration was varied between 0 and 0.72 mol L� 1. The relative percentage

extractions of Fe, Ni, Co and Mn indicate that the Fe and Ni extractions are inter-related at a ratio of Ni/Fe = 0.7–0.9 and

suggest the possibility of catalysis of manganese dissolution by solubilized iron(II). This leads to a Mn extraction of over 90%

in less than 30 min compared with 20–40% Fe extraction in the same period, depending on the acid concentration. The initial

rate of leaching of iron shows first-order dependence with respect to H+. Whilst the synthetic iron oxides leach according to the

shrinking particle/sphere kinetic model, the results obtained in the first 4 h of laterite leaching can be described by a shrinking

particle model with an insoluble product layer that retards the diffusion of H+ to the reaction sites at the interface. The

heterogeneous rate constants for both models increase with the increase in H+ concentration. The effective diffusion coefficient

of H+ (DH+ ) through the product layer (0.5� 10� 9 to 4� 10� 9 cm2 s� 1), determined in the present study, is in the magnitude

range of the reported data for DH+ in polycrystalline Fe3O4 and MnO2, but lower than DH

+ in aqueous media, 9� 10� 5 cm2 s� 1.

D 2003 Elsevier B.V. All rights reserved.

Keywords: Limonitic laterite; Iron; Nickel; Cobalt; Manganese; Leaching; Heterogeneous kinetic models; H+ diffusion

1. Introduction

Sulfur dioxide is an efficient leaching agent for

minerals containing oxides of iron, nickel, cobalt and

manganese (Byerley et al., 1979; Miller and Wan,

1983; Abbruzzese, 1990; Grimanelis et al., 1992;

Kumar et al., 1993; Das et al., 1997). It offers the

0304-386X/$ - see front matter D 2003 Elsevier B.V. All rights reserved.

doi:10.1016/S0304-386X(03)00132-4

* Corresponding author. Fax: +61-8-9360-6343.

E-mail address: [email protected]

(G. Senanayake).

choice of acidic leaching of these metal oxides in

nickeliferous limonite under atmospheric conditions

compared to the acid pressure leaching commercially

practiced in Cuba and Western Australia (Chou et al.,

1977; Kyle, 1996). The kinetics of the acid dissolution

(Majima et al., 1985; Cornell et al., 1976) and

reductive dissolution of natural or synthetic iron

oxides in the presence or absence of sulfur dioxide

(Warren and Hay, 1975; Kumar et al., 1993; Byerley

et al., 1979; Chiarizia and Horwitz, 1991) as well as

the reductive dissolution of manganese dioxide by

Table 1

Equilibrium constants (logK) for protonation and sulfate complex-

ation at 25 jCa

Reaction Ionic strength logK

H++ SO42� =HSO4

� 0 1.99 (3.07at 100 jC)1 1.07

H+ +HSO3� =H2O+SO2 0 1.77

1 1.37

Fe2 + + SO42� = FeSO4

0 0 2.3

1 1.0

G. Senanayake, G.K. Das / Hydrometallurgy 72 (2004) 59–7260

sulfur dioxide or by Fe(II) sulfate (Miller and Wan,

1983; Tekin and Byramoglu, 1993) have been studied

in detail. The relevant kinetic data have been corre-

lated with one or more of the physico-chemical factors

that include the ionic activity of H+, complexation

with OH� or SO32�, electrochemical surface reaction

and changes in surface area.

Along with manganese, which may also be present

as pyrolusite, MnO2, nickel and cobalt in laterite

deposits are enriched in weathering products such as

goethite and limonite by ion replacement (Kyle,

1996). A proper understanding of the kinetics and

mechanism of reductive leaching of oxide minerals by

acid and sulfur dioxide is useful in developing meth-

ods for selective leaching of valuable metals from

nickeliferous ores. The reaction mechanisms for pure

components proposed by previous researchers are

important in the discussion of the reductive leaching

of limonitic laterite ore containing multi-valent oxides

of important metals. However, the interpretation of

kinetic behavior of such ores is complicated due to the

fact that they contain more than one component that

can react with reductive leaching agents such as sulfur

dioxide and/or iron(II) sulfate. For example, the

limonitic laterite used in one of the previous inves-

tigations (Das et al., 1997) and the subject of this

study consists of goethite a-FeOOH and magnetite

Fe3O4 as major components and hematite a-Fe2O3 as

the minor component along with quartz. Whilst the

Fe(III) oxides are dissolved by H+ and/or by SO2 to

produce Fe(III) and Fe(II), the SO2 as well as Fe(II)

produced during reduction can subsequently act as

reductive leachants for Mn(IV), Fe(III) and Co(III).

This paper describes the analysis of leaching results of

limonitic laterite in sulfuric acid and in perchloric acid

in the absence or presence of sulfur dioxide on the

basis of surface chemical reaction with H+ and SO2 as

well as Fe(II) and the heterogeneous mechanisms

described by the shrinking particle models with or

without insoluble product layer.

Fe2 + +HSO4� = FeHSO4

� 1.2 0.78

Fe3 + + SO42� = FeSO4

+ 1.2 2.23

Fe3 + + 2SO42� = Fe(SO4)2

� 1.2 4.23

Ni2 + + SO42� =NiSO4

0 0 2.40

1 0.57

Mn2 ++ SO42� =MnSO4

0 0 2.28 (3.0 at 45 jC)Co2 ++ SO4

2� =CoSO40 0 2.36

1 1.06

a Data from Sillen and Martell, 1964.

2. Iron chemistry and kinetic models

2.1. Acid dissolution of iron oxides

The initial rate of dissolution of hematite or mag-

netite (Eqs. (1) and (2)) follows first-order kinetics

with respect to the concentration of H+ (Cornell et al.,

1976) due to the slow desorption of FeOH2 + from the

surface, followed by aqueous reaction as shown by

Eqs. (3)–(5) (Warren and Hay, 1975):

FeOOHðsÞ or 0:5Fe2O3H2OðsÞþ3Hþ ¼ Fe3þþ 2H2O

ð1Þ

Fe3O4ðsÞ or Fe2O3 � FeOðsÞ þ 8Hþ

¼ 2Fe3þ þ Fe2þ þ 4H2O ð2Þ

pOFe� OHðsÞ þ Hþ

¼ pOFeþðsÞ þ H2O ðfast equilibrationÞ ð3Þ

pOFeþðsÞ þ Hþ ¼ FeOH2þðaqÞ ðslowÞ ð4Þ

FeOH2þðaqÞ þ Hþ ¼ Fe3þ þ H2O ðfastÞ ð5Þ

Despite the fact that the rate-determining step (4) is

first order with respect to H+, the formal reaction order

of H+ can show different values in the range 0.5–2

depending on the hydration, complexation, redox

reactions and other factors (Azuma and Kametani,

1964; Gorichev and Kipriyanow, 1984; Majima et

al., 1985). For example, Table 1 lists some of the

relevant complexes and their equilibrium constants (K)

showing that K increases with increasing temperature,

but decreases with increasing ionic strength. Potenti-

Fig. 1. Eh–pH diagram for Fe(III)/Fe(II)/SO2/SO42� system at 25

jC, based on thermodynamic data in Table 1 at ionic strength 1 and

equimolar species in solution.

G. Senanayake, G.K. Das / Hydrometallurgy 72 (2004) 59–72 61

ometric measurements and predicted Eh–pH and spe-

cies distribution diagrams have shown that FeSO40 and

Fe(SO4)2�, respectively, are the predominant Fe(II) and

Fe(III)–sulfate complex species in acidic sulfate sol-

utions (Senanayake and Muir, 1988). The Eh–pH

diagram shown in Fig. 1 highlights the need to rewrite

the stoichiometry of acid dissolution shown in Eqs. (1)

and (2) to accommodate actual complex species such

as FeSO40 and Fe(SO4)2

� in solution:

FeOOHðsÞ or 0:5Fe2O3H2OðsÞ þ 2H2SO4

¼ FeðSO4Þ�2 þ Hþ þ 2H2O ð6Þ

Fe3O4ðsÞ or Fe2O3FeOðsÞ þ 5H2SO4

¼ 2FeðSO4Þ�2 þ 2Hþ þ FeSO04 þ 4H2O ð7Þ

2.2. Effect of SO2

The catalytic effect of a reducing agent is the result

of a requirement that the surface of the mineral be

composed of a stoichiometric compound that will

dissolve into the solution (Nicol, 1983). For example,

the proposed mechanism for the dissolution of mag-

netite and goethite in SO2 involves slow desorption of

FeHSO3+ from the oxide surface (Byerley et al., 1979).

The presence of reducing agents capable of lowering

the [Fe(III)]/[Fe(II)] concentration ratio will accelerate

iron oxide dissolution (Gorichev and Kipriyanow,

1984). This shows the importance of considering the

reaction between Fe(III) and SO2 in solution.

Solution chemistry studies of Fe(III) and SO2 have

shown a stoichiometry of 1:1.2 for the reversible

reaction between Fe(III) and SO2 to form Fe(II) and

the monomer HSO3 free radical or the dimer H2S2O6

of dithionic acid at pH 0.5 (Higginson and Marshall,

1957). Despite the experimental evidence for the

formation of both S2O62� and SO4

2� at pH 0–3 (Sato

et al., 1978), IR spectroscopic studies have indicated

that SO42� is the major species formed during the

leaching of goethite with SO2 (Kumar et al., 1993).

Considering the formation of both dithionate and

sulfate and the sulfate complex of Fe(II), the reductive

dissolution of goethite by SO2 may be represented as:

FeOOHðsÞ or 0:5Fe2O3H2OðsÞ þ SO2 þ H2SO4

¼ FeSO04 þ 0:5H2S2O6 þ H2O ð8Þ

Fe3O4ðsÞ or Fe2O3 � FeOðsÞ þ 3H2SO4 þ 2SO2

¼ 3FeSO04 þ H2S2O6 þ 2H2O ð9Þ

H2S2O6 ¼ H2SO4 þ SO2 ð10Þ

FeOOHðsÞ or 0:5Fe2O3�H2OðsÞþ0:5SO2þ0:5H2SO4

¼ FeSO04 þ H2O ð11Þ

Fe3O4ðsÞ or Fe2O3 � FeOðsÞ þ 2H2SO4 þ SO2

¼ 3FeSO04 þ 2H2O ð12Þ

2.3. Kinetic models

The initial rates of leaching have been used to

model the kinetics of the dissolution of hematite on

Table 2

Experimental conditions, apparent rate constants (kss and kpl) and

proton diffusion coefficients (DH+) for iron dissolution from

synthetic iron oxides and limonitic laterite ore

Material Particle

size

(�m

mean)

Pulp

density

(%,

wt/vol)

T

(jC)SO2

(M)

H2SO4

(M)

106

kss(s� 1)

106

kpla

(s� 1)

109

DH+ b

(cm2

s� 1)

FeOOHc 17 2.5 80 nil 3 111

0.1 3 157

Fe3O4d 75 0.4 50 0.54 nil 1.7

Lateritee

Test 1 (T1) 106 10 90 0.3 nilf 0.06 0.12

Test 2 (T2) nil 0.72 3.3 0.51

Test 3 (T3) 0.3 0.18 1.5 0.87

Test 4 (T4) 0.3 0.36 3.3 1.0

Test 5 (T5) 0.3 0.54 30g 10 2.0

Test 6 (T6) 0.3 0.72 43g 27 4.2

Test 7 (T7) 0.3 nilh 18 2.2

a Rate constant for shrinking particle model with solid product

layer for 4 h (Fig. 10).b Proton diffusion coefficient based on Eqs. (14) and (15):

r = 5.3� 10� 3 cm, qm= 3.8 g cm� 3 (Berkman, 1995) and %Fe = 39.c Synthetic goethite (Chiarizia and Horwitz, 1991).d Synthetic magnetite,1 g solid in 250 mL liquid (Byerley et al.,

1979).e This work.f No H2SO4 added, natural [H+] = 0.062 M, based on mass

balance for Eq. (16).g Rate constant for shrinking particle model for the first 60 min

(Das et al., 1997).h No H2SO4 added, 1 M HClO4.


the basis of surface chemical reaction with H+ in the

absence of SO2 (Cornell et al., 1976; Majima et al.,

1985) or surface electrochemical reaction in the pres-

ence of SO2 (Kumar et al., 1993). For example, the

initial rate of dissolution of a-FeOOH, h-MnO2 and

g-MnO2 has been found to be proportional to [SO2]0.5

indicating a half order reaction with respect to SO2

(Miller and Wan, 1983; Kumar et al., 1993). In the

case of relatively fast surface reactions, it is essential

to consider the decrease in surface area in a shrinking

sphere model with no product layer formation (Leven-

spiel, 1972; Ray, 1993). For the first-order 1:1 molar

heterogeneous rate-controlling step for the dissolution

of FeOOH (Eq. (4)), this model can be represented by

the mathematical relationship:

1� ð1� X Þ1=3 ¼ ½Hþbulkkr�1q�1t ¼ ksst ð13Þ

where, [H+]bulk (mol cm� 3) = concentration in the

bulk solution, k (cm s� 1) = intrinsic rate constant of

the surface chemical reaction having the units of a

mass transfer coefficient, r (cm) = initial particle radi-

us of solid, q (mol cm� 3) =molar concentration of Fe

in the solid, kss (s� 1) = apparent rate constant in the

shrinking sphere kinetic model and X = fraction of Fe

reacted in time t. The relationship between q and wt.%

Fe in the material and its density (qm, g cm� 3) is:

q¼ qm�%Fe in material=ð100� molar mass of FeÞð14Þ

If a porous solid product layer is formed on the

surface during the reaction, the slow diffusion of H+

through the product layer becomes the rate-controlling

step. The shrinking core with product layer model in

such cases is represented by the equation:

1� 3ð1� X Þ2=3 þ 2ð1� X Þ

¼ 6½HþbulkDþHr

�2q�1t ¼ kplt ð15Þ

where DH+ (cm2 s� 1) = diffusion coefficient of H+

through the product layer and kpl (s� 1) = apparent rate

constant. Eq. (13) assumes that the decrease in rate

with time is due to the decrease in particle size and

thus the surface area. Eq. (15) assumes that the

diffusion of H+ through a solid product layer is the

rate-determining step and the decrease in rate is due to

the increase in thickness of the solid product layer

with time, whilst the spherical nature and the radius of

the particle remains unchanged.

The leaching of pure MnO2 has been studied in

SO2/Na2SO3/pH 1–2 (Miller and Wan, 1983) and in

FeSO4/H2SO4 (Tekin and Byramoglu, 1993) at tem-

peratures up to 50 jC under controlled hydrodynam-

ic conditions. The stirring speed was sufficiently

large to confirm that the rate of leaching was not

governed by the transport process. The decrease in

the rate of leaching with time has been related to the

shrinking particle/sphere kinetic model with a value

of kss = 4.4� 10� 3 s� 1 and k = 1.2� 10� 3 cm s� 1

for the leaching of MnO2 in 0.5 M SO2 at pH 2 at a

stirring speed of 650 rpm. The value of k was less

than the predicted mass transfer coefficient for

suspended particles, 2.4� 10� 2 cm s� 1 in the leach-

ing reactor, showing that the leaching rate was con-


trolled by the surface reaction rate (Miller and Wan,

1983).

3. Experimental

Awater-jacketed 1-L Pyrex glass vessel fitted with

a Teflon-coated stirrer, baffles, gas inlet tube, sampling

tube, thermometer and water condenser was used as

the reactor for leaching; other details of dimensions

were similar to those reported by Tekin and Byramoglu

(1993). The limonitic nickel laterite of particle size

range 90–125 Am from Bulong, Western Australia,

was used in all experiments. The standard leaching

conditions were: pulp density = 10% (wt/vol), temper-

ature = 90 jC, agitation speed = 650 rpm, leaching time

6 h. The sulfur dioxide flow rate was maintained at 0.6

L min� 1 L� 1 of slurry and the acid concentration was

varied between 0 and 0.72 M H2SO4 or 1 M HClO4 in

seven tests T1–T7 (Table 2). These conditions were

consistent with the previous investigation (Das et al.,

1997). Samples were collected at 0.5, 1, 2, 4 and 6 h,

and the solutions were analysed for Ni, Co, Mn and Fe

by atomic absorption spectrophotometry.

Fig. 2. Kinetic plots for iron extraction from limonitic laterite. T1, T2 and T

2 for conditions.

4. Results and discussion

4.1. Time and acid dependence of Fe and Ni

extraction

Fig. 2 plots %Fe extractions from limonitic laterite

at time intervals 0.5, 1, 2, 4 and 6 h and compares

with some of the data reported previously (Das et al.,

1997). The leaching with SO2 in the absence of added

acid (Test 1) shows the lowest iron extraction but a

linear increase in % extraction with time. The %Ni

extraction presented in Fig. 3 also shows the lowest in

Test 1, but %Ni remains at 20% and independent of

time. The highest Fe and Ni extraction at 6 h is shown

by SO2 + 0.72 M H2SO4 (Test 6) but the rate decreases

with time. In Test 3, Ni extraction reaches only 25%

but Co and Mn reach 100% in 30 min.

The %Fe and %Ni extraction after a given

leaching time generally increases in the order of

acid concentration: Test 1 < Test 3 < Test 4 < Test

5 < Test 6. This can be mainly attributed to the

initial rate of leaching with SO2, represented by

the initial slopes in Figs. 2 and 3, which also

increase in the same order. However, the rate

6 from Das et al. (1997); other data from the present study, see Table

Fig. 3. Kinetic plots for nickel, cobalt and manganese extraction from limonitic laterite. Solid lines represent Ni extraction in T1 to T6 (Table 2),

dashed line represents Co and Mn extraction in T3.


decreases with time, except in Tests 2 and 3 carried

out with no SO2 or low acid concentrations (0.18

M), respectively, where the rate after 2 h remains

fairly constant. In general, the decrease in rate with

increasing leaching time can be related to one or

more of the following reasons: (i) the decrease in

acid concentration with the consumption of H+ in

Eqs. (6)–(12), (ii) the change in equilibrium SO2

concentration in Eq. (16) caused by the decreasing

acid concentration and changing ionic strength (Ta-

ble 1), (iii) the decrease in surface area of the

particles or (iv) the increase in thickness of a solid

layer, which retards the diffusion of reactants or

products.

HSO�3 þ Hþ X SO2 þ H2O ð16Þ

Unlike in the case of Fe and Ni, the % extraction of

Co and Mn rapidly increases and reaches 85–100%

during the first 30 min in the presence of SO2 due to

the reactivity of SO2 and the occurrence of Co with

Mn (Das et al., 1997). Consequently, despite the

high acid concentration of 0.72 M in Test 2, the

extraction of these metals remains low at 12–28%

Co and 20–40% Mn in 0.5–6 h of extraction time

due to the absence of SO2 (see later in Fig. 5). This

behavior can be further examined by considering the

relative extractions of the four metals.

4.2. Relative extraction of Fe, Ni, Co and Mn

4.2.1. Ni vs. Fe

Despite the higher iron content compared to nickel

in the starting material, 39.05% Fe, 1.17% Ni, 0.124%

Co, 1.25% Mn (Das et al., 1997), Fig. 4 shows linear

correlations between the %Ni and %Fe extractions at

different time intervals up to 6 h. These linear rela-

tionships of slopes ranging from 0.7 to 0.9 in the

range 20–85% Ni extractions can be related to the

incorporation of nickel in the iron oxide fraction of

laterite and thus the release of nickel(II) into the

solution is associated with the leaching of iron.

However, the nickel and iron extraction by sulfur

dioxide in the absence of added acid (Test 1) was

low, < 20% Ni and < 10% Fe, and the results do not

show a correlation with the other data in Fig. 3. In

contrast, 0.72 M H2SO4 alone with no SO2 (Test 2)

extracts c 45% Fe and Ni; SO2 + 0.72 M H2SO4

(Test 6) extracts c 85% Fe and Ni in 6 h. These

results also show the necessity of breaking up of the

goethite structure by SO2 and/or H+ to release both

iron and nickel into solution.

Fig. 4. Correlation between nickel and iron extraction at different time intervals in Tests T1–T7.


4.2.2. Ni, Co, Mn vs. Fe

Fig. 5 plots the % extraction of Ni, Co and Mn

against Fe in selected Tests 2, 6 and 7. The lower

extraction of 45% Fe in Test 2 in 6 h is clearly due to

the fact that the dissolution of iron oxide is caused

only by the acid attack according to Eqs. (6) and (7) in

the absence of SO2. This is further supported by the

measured Eh of the leach liquors in Test 2, which

were cooled to 25 jC. The measured Eh in the range

0.60–0.68 Vat pH c 1.5 in Test 2 corresponds to the

stability region of Fe(SO4)2�, indicating a ratio of

[Fe(III)]/[Fe(II)] greater than unity (Fig. 1). Although

the %Ni vs. %Fe extraction corresponds to the line of

slope 0.9, Co and Mn extraction corresponds to lower

slopes of 0.5–0.8 due to the lower extraction in the

absence of SO2 in Test 2.

Fig. 5 also shows the % extraction of Ni, Co and

Mn against Fe in the three tests to compare the effect of

SO2 in 0.72 M H2SO4 (Tests 2 and 6) and the effect of

changing 0.72 MH2SO4 to 1 MHClO4 (Tests 6 and 7).

The measured Eh of the leach liquors produced in the

presence of SO2 in Tests 4–6, which were cooled to 25

jC, were lower (0.56 V) than in the case of Test 2. ThisEh falls in the stability region of FeSO4

0 in Fig. 1 and

shows the reductive dissolution of ore according to

Eqs. (8)–(12) forming FeSO40, causing the [Fe(III)]/

[Fe(II)] ratio to drop to values below unity. The main

feature in Fig. 5 is that the Co and Mn extractions

remain 85–100% even at lower Fe extractions; this is

much higher than the Co and Mn extractions ( < 40%)

in the absence of SO2 in Test 2. Higher and rapid

extraction of Mn and Co in the presence of SO2 is a

result of the reductive leaching of the type:

MnO2ðsÞ þ SO2 ¼ MnSO04 ð17Þ

The mechanism of this reaction has been well docu-

mented (Miller and Wan, 1983; Abbruzzese, 1990;

Grimanelis et al., 1992). High leaching rates at pH 1–2

have been related to the formation of complexes of

Mn(II/III) of the type Mn(SO3)22� and Mn(SO3)2

�,

respectively.

Fig. 5. Correlation between nickel, cobalt, manganese and iron extraction at different time intervals. Comparison between the effect of SO2 (T2

and T6), 0.72 M H2SO4 and 1 M HClO4 (T6 and T7).


Fig. 5 shows that the Ni vs. Fe extraction in Tests 6

and 7 corresponds to a line of slope 0.7, lower than

the slope 0.9 for data in Test 2, irrespective of the

different acids, 0.72 M H2SO4 (in Test 6) and 1 M

HClO4 (in Test 7), but remains above the line of slope

0.9. This indicates a higher Ni extraction, compared to

Fe, in the presence of SO2. However, the Ni extraction

in Tests 6 and 7 starts to drop back to the line of slope

0.9 after the complete extraction of Co and Mn. It is

possible that the decrease in Fe extraction is due to the

involvement of the dissolved Fe(II) in the reductive

dissolution of MnO2, and the cobalt associated with it,

in acidic media causing Fe(II) to re-precipitate as

Fe(III) species

4.3. Catalytic effect of Fe(II)

The catalysis of reductive leaching of MnO2 by

H+/Fe(II) takes place via the intermediate MnOOH

(Tekin and Byramoglu, 1993) leading to the partial

precipitation of iron as the basic sulphate FeSO4OH or

the oxides Fe2O3, Fe3O4 or FeOOH due to the

hydrolysis and/or the increase in Fe(II) concentration

in solution:

MnO2ðsÞ þ 2FeSO04 þ H2O

¼ MnSO04 þ Fe2O3ðsÞ þ H2SO4 ð18Þ

MnO2H2OðsÞþFeSO04 ¼ MnOOHðsÞ þ FeSO4OHðsÞ

ð19Þ

MnOOHðsÞ þ 0:5SO2 þ 0:5H2SO4 ¼ MnSO04 þ H2O

ð20Þ

MnOOHðsÞ þ FeSO04 ¼ MnSO0

4 þ FeOOHðsÞ ð21Þ

FeSO4OHðsÞ þ H2O ¼ FeOOHðsÞ þ H2SO4 ð22Þ

2FeSO4OHðsÞ þ H2O ¼ Fe2O3ðsÞ þ 2H2SO4 ð23Þ

2FeSO4OHðsÞ þ FeSO04 þ 2H2O

¼ Fe3O4ðsÞ þ 3H2SO4 ð24Þ

G. Senanayake, G.K. Das / Hydro

Some evidence for these conclusions and chemical

reactions also comes from the previous studies (Gri-

manelis et al., 1992) with the sulfur dioxide leaching

of a manganese rich pyrolusite ore containing 40.74%

MnO2 and 0.84% Fe2O3 at ambient temperatures. The

extraction of Mn and Fe at 30 jC in the first 2 min

was 75% Mn and 20% Fe, respectively. In 20 min, Mn

extraction increased to 85%, which remained constant

over time, whilst the Fe extraction decreased to about

10% during the same time interval due to the re-

precipitation of hydrated Fe(III) oxide or an insoluble

basic sulfate. Moreover, the ratio of %Mn/%Fe in-

creased with the decrease in pH indicating that high

acidity led to more Mn dissolved relative to Fe, due to

the catalytic effect of Fe(II).

It is also of interest to note that the XRD patterns

showed peaks indicating the presence of FeOOH,

Fe3O4, Fe2O3 and SiO2 in the starting material used

in the present study, but the presence of only Fe3O4 and

SiO2 in the leach residue (Das et al., 1997). This

observation supports the formation of Fe3O4 Eq. (24),

in contrast to the hydrothermal conversion of goethite

to hematite in the acid pressure leaching of limonitic

laterite (Briceno and Osseo-Asare, 1995) where there

were no peaks corresponding to Fe3O4 in the leach

residue. The formation of FeSO4OH(s) as a solid

species in sulfuric acid pressure leaching of laterite

has not been confirmed due to the extremely fast ki-

netics of Eq. (23) (Rubisov and Papangelakis, 2000).

4.4. Effect of changing H2SO4 to HClO4

When the acid was changed from 0.72 M H2SO4 to

1 M HClO4 in the present study (Test 6 to Test 7), the

Mn and Co extraction decreased from 100% to 85–

95% in Fig. 5, although the Ni extraction was unaf-

fected. This shows the influence of background sulfate

on Mn and Co extraction, compared to background

perchlorate. It is possible that the formation of stable

sulfate complexes of MnSO40 and CoSO4

0 (Table 1) as

well as the intermediates such as FeSO4OH(s) and

FeSO40 formed in sulfate solutions as indicated in Eqs.

(18)–(24) would favor the thermodynamics and kinet-

ics of leaching of these metals. Thus, the results

summarized in Figs. 2–5 clearly indicate that the

overall leaching behavior of laterite is largely con-

trolled by the kinetics and mechanism of the dissolu-

tion of iron.

4.5. Kinetic models

4.5.1. Initial rate of iron leaching

Results from the previous studies on initial rate of

dissolution of goethite failed to agree with the rate

equation: � {dnFeOOH/dt} = constant [H+]n[SO2]m

with the proposed values of n =m = 0.5 based on an

electrochemical reaction model (Kumar et al., 1993).

The experimental values for m for leaching in 0.3–

0.5 M SO2 at pH 1.1–2.2 varied between 0.50 and

0.67; but n showed a much larger variation from 0.17

to 0.40. This leads to the suggestion that leaching

may also be taking place via acid attack in addition to

the reductive leaching (Kumar et al., 1993). However,

the literature data for the dissolution of hematite

(Majima et al., 1985) in sulfuric acid in the absence

of sulfur dioxide at 50–55 jC plotted in Fig. 6 show

a slope of 0.67 with respect to acid concentration.

Therefore, it is important to establish the order of the

initial leaching reaction with respect to H+, before the

effect of change in surface area and/or the formation

of insoluble solids at the reaction interface can be

considered.

The solubility of SO2 in water and in 1 mol L� 1

H2SO4 at 1 atmospheric pressure of SO2 is c 0.3 mol

L� 1 and fairly independent of the acid concentration

(Linke and Seidell, 1958). Therefore, it is reasonable

to assume that the SO2 concentration remains constant

in the tests carried out in the present study under a

constant and continuous flow of SO2. Moreover the

Eh–pH diagram predicts that SO2 is the predominant

species (Fig. 1) and that the equilibrium concentration

of HSO3� in Eq. (16) is negligibly small at pH < 2

(Abbruzzese, 1990). Thus, the slope is close to 1 for

the plot of log{d[Fe]/dt} against log[H+] in Fig. 6 for

the initial (t = 0–0.5 h) leaching in Tests 2–7, reflect-

ing first-order kinetics with respect to H+. This is in

agreement with the previous studies for the dissolution

of a-FeOOH in perchloric acid (Cornell et al., 1976).

4.5.2. Shrinking particle (sphere) model for iron

leaching

The kinetic studies for the dissolution of synthetic

iron oxides in H2SO4 in the absence or presence of

SO2 carried out by previous researchers have provided

evidence for a shrinking particle model (Chiarizia and

Horwitz, 1991). Some of the literature data for goe-

thite and magnetite dissolution are summarized in Fig.

metallurgy 72 (2004) 59–72 67

Fig. 6. Effect of acid concentration on initial rate of iron leaching. Laterite: Das et al. (1997) and this work (Table 2). Hematite: Majima et al.

(1985) (74–104 Am 95.4% Fe2O3).


7 as plots of 1� (1�X)1/3 against time and the rate

constants kss are summarized in Table 2. The value of

kss for the dissolution of synthetic FeOOH in 3 M

H2SO4 at 80 jC increases by a factor of 1.4 with the

addition of 0.1 M SO2. The dissolution of synthetic

Fe3O4 in 0.54 M SO2 at 50 jC corresponds to a kssvalue that is 1/100 times less than that for FeOOH at

80 jC with 3 M H2SO4, due to the lower temperature

and absence of added acid (Table 2).

Fig. 7. Comparison of kinetic plots for iron dissolution from goethite, magn

References and conditions described in Table 2.

On the basis of the dissolution behavior of the

synthetic goethite and magnetite summarized in Fig.

7, it was expected that the dissolution of iron in the

forms of goethite, magnetite and hematite in the

limonitic laterite ore investigated in the present study

would follow the same trend. However, as reported

previously (Das et al., 1997), the iron dissolution

obeys the shrinking particle kinetic model only during

the first hour of leaching and the relevant kss data are

etite and limonitic laterite: shrinking sphere (particle) kinetic model.


listed in Table 2. Limited data (Tests 2 and 6) shown

in Fig. 7 confirm this behavior.

As a matter of interest, the value of kss for the dis-

solution of h-MnO2 in 0.5 M SO2 (Tekin and Byra-

moglu, 1993) and g-MnO2 in FeSO4 (Miller and Wan,

1983) at pH 1–2 seem to be 1000–5000 times greater

than the values for iron reported in this work leading

to faster leaching of Mn compared to Fe. This

explains why the %Mn and %Co extraction reached

100% in Tests 3 in 30 min compared to 7% Fe and

25% Ni (Figs. 2 and 3).

4.5.3. Shrinking core model with product layer for

iron leaching

The results from the two tests T2 and T6 in 0.75 M

H2SO4 in the presence or absence of SO2, plotted in

Fig. 7 do not show a linear correlation for the shrinking

particle model over the longer leaching period of 6 h.

Figs. 8 and 9 compare the plots of 1� (1�X)1/3 and

1� 3(1�X)2/3 + 2(1�X) against t, respectively, for

the dissolution of Fe, Ni, Co and Mn in 0.72 M H2SO4

in the absence of SO2 (Test 2) during the first 2 h.

Clearly, the Fe and Ni extraction in Fig. 9 show an

excellent linear relationship with correlation coeffi-

cients of R2c 0.99 and the rate constant kpl = 4.2�10� 6 s� 1 for iron and 6.1�10� 6 s� 1 for nickel. The

fact that manganese and cobalt do not behave in the

same way as iron and nickel supports the view that the

dissolution of manganese (and cobalt) is a result of the

Fig. 8. Two-hour kinetic data: shrinking sphere (particle) kinetic model for

H2SO4 in the absence of SO2 (Test 2).

reactions with SO2 and Fe(II) described previously.

Fig. 10 shows that all the results obtained for the

dissolution of iron during the first 4 h in the present

investigation fit with a shrinking particle model with a

solid product layer, with R2>0.99 for Tests 3 and 6 and

R2c 0.98 for the others. The kpl values obtained from

the slopes are listed in Table 2. The product layer may

consist of basic iron sulfate and/or iron oxides (Eqs.

(18)–(24)) and quartz, in addition to any other hydro-

lysis products from the gangue.

4.6. Effect of H+ concentration on rate constants

Both H+ and SO2 should be available for the

reductive leaching reactions represented by Eqs. (8)–

(12). The value of the heterogeneous reaction rate

constant, kpl increases eightfold from 3.3� 10� 6 s� 1

in 0.72 MH2SO4 to 27� 10� 6 in SO2 + 0.72MH2SO4

(Table 2), mainly due to the change from acid leaching

(Eqs. (6) and (7)) to reductive leaching (Eqs. (8)–

(12)). Since the SO2 concentration was maintained

constant in the present study, it is reasonable to assume

that the increase in kpl with increasing concentration of

H2SO4 is mainly due to the change in H+ concentration

in the bulk (Eq. (15)). The value of pKa for Eq. (26),

H2SO4 ! Hþ þ HSO�4 ð25Þ

HSO�4 X Hþ þ SO2�

4 ð26Þ

dissolution of Fe, Ni, Co and Mn from limonitic laterite with 0.72 M

Fig. 9. Two-hour kinetic data: shrinking core model with a solid product layer for dissolution of Fe, Ni, Co and Mn from limonitic laterite with

0.72 M H2SO4 in the absence of SO2 (Test 2).


depends on the ionic strength and temperature

(Table 1). The ionic strength of leach liquor changes

with time due to the formation of sulfur species

Fig. 10. Four-hour kinetic data: shrinking core model with a

and dissolved metal ions and their complexes. If it

is assumed that the ionic strength remains close to

1, the mass balance for equilibrium in Eq. (26)

solid product layer for dissolution of Fe in Tests 1–7.

Table 3

Effect of medium and material on proton diffusion coefficient DH+

Medium/material DH+(cm2 s� 1) Reference

Aqueous solutiona 9� 10� 5 Bockris and Reddy, 1977

Single-crystal ironb 8� 10� 5 Bockris and Reddy, 1977

g-MnO2 8� 10� 8 Allen et al., 1979

Product layer in

laterite leaching

0.5� 10� 9 to

4� 10� 9

This work (Table 2)

Fe3O4 8� 10� 10c Allen et al., 1979

NiO 10� 10–10� 12 Allen et al., 1979

a At infinite dilution and 25 jC.b Diffusion of eloctronated H+ (atomic hydrogen) into the metal.c Minimum possible value based on electrochemical studies

with polycrystalline magnetite.


based on the pKa value at I = 1 leads to the linear

relationship:

½Hþ ¼ 1:085½H2SO4 þ 0:0178 ð27Þ

This can be used to examine the effect of H2SO4

concentration on the rate constants. A higher con-

centration of [H+]bulk would favor the adsorption

equilibrium: Solid +H+ X Solid�H+(ads) and thus the

rate controlling surface chemical reaction (Eq. (4);

Cornell et al., 1976) occurs to improve the rate of

dissolution. This would increase kss according to

Eq. (13), as reflected in the data listed in Table 2

for Tests 5–6. Likewise, the increase in [H+]bulkwould increase the rate of diffusion of H+ through

the product layer to the reaction interface and hence

increase kpl according to Eq. (15).

The calculated values of DH+ , based on Eq. (15)

listed in Table 2 range from 0.5� 10� 9 to 4.2� 10� 9

cm2 s� 1. The variation of DH+ with the change in acid

concentration largely reflects the non-validity of

the assumptions of constant values of r and [H+]bulkused in the calculation of DH

+ using Eq. (15). Never-

theless, the increasing order of DH+ shown in Table 3

indicates that DH+ calculated in the present study is

c 1/10000 times smaller than DH+ in aqueous solutions

but within the magnitude range reported for the diffu-

sion of H+ through solids such as Fe3O4 and g-MnO2

based on electrochemical methods (Allen et al., 1979).

5. Summary and conclusions

� At constant [SO2], the rate of leaching of iron from

limonitic laterite in the first 30–60 min is

approximately first order with respect H+ but

appears to obey the shrinking particle model due

to the decrease in surface area.� This is consistent with the leaching behavior of

synthetic iron oxides in the absence or presence of

SO2, which also follows first-order kinetics and the

shrinking particle kinetic model.� Subsequently, the rate of leaching appears to be

controlled by the slow diffusion of H+ through an

insoluble solid layer produced during leaching. The

magnitude of the diffusion coefficient of H+

through the insoluble solid product (0.5�10�9 to

4�10�9 cm2 s�1) is much less than that in aqueous

media, but in the same order as that in solids such

as Fe3O4 and g-MnO2.� The dissolution of nickel follows the same trend as

iron with an extraction ratio of Ni/Fe = 0.7–0.9 and

obeys the shrinking core model with a solid product

layer. However, the dissolution of manganese (and

cobalt), which reaches 90–100% in the first 30 min

before the product layer sets in, does not follow the

same trend as iron and nickel due to the direct

reaction with SO2 and the catalytic effect of Fe(II)

produced by the reductive leaching of iron oxide.

Acknowledgements

The authors thank Profs. David Muir and Pritam

Singh for valuable discussion during the Targeted

Institutional Research Links Program.

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A comparative study of leaching kinetics of limonitic laterite and

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