~ Pergamon

Canadian Metallurgical Quarterly, Vol. 35, No.4, pp. 363-373, 1996Copyright «;;1996Canadian Institute of Mining and Metallurgy. Published by Elsevier Science LId

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THERMODYNAMIC ANALYSIS OF THE REACTIONS OFIRON OXIDES: DISSOLUTION IN OXALIC ACID

D. PANIAS, M. TAXIARCHOU, I. DOUNI, I. PASPALIARIS andA. KONTOPOULOS

Laboratory of Metallurgy, National Technical University of Athens, GR-15? 80 Zografos, Greece

(Received 6 July 1995; in revisedfarm 4 March 1996)

Abstract-The mechanism of dissolution of pure iron oxides by means of organic acids comprises twodifferent chemical pathways: (a) non-reductive dissolution and (b) reductive dissolution. In this paper, thethermodynamic analysis of these pathways for the iron oxides-oxalic acid system is presented. In low acidsolutions (pH higher than 3) the only thermodynamically stable complex ions of bivalent and trivalent ironare [Fe2+(C20.)2)2- and [Fe3+(C20.)3)3-. Uncomplexed Fe2+ ion can be identified only in high acidsolutions, while uncomplexed Fe3+ ion is not likely to build-up in oxalic acid solutions. In the pH range1-2 the [Fe3+(C20.h)- and [Fe3+C20.t ions are stable, while at pH less than I, the [Fe3+HC20.)2+ is theonly ion existing. Copyright ~ 1996 Canadian Institute of Mining and Metallurgy. Published by ElsevierScience Ltd

Resume-La dissolution d'oxides de fer purs par les acides organiques peut etre realisee selon deuxmechanismes chimiques differents: (a) une dissolution non reductive et (b) une dissolution reductive.L'analyse thermodynamique de ces deux traitements chimiques pour Ie systeme oxide de fer-acide oxaliqueest presente dans cet article. En solutions faiblement acides (pH> 3), les seuls complexes ioniques du ferbivalent et trivalent thermodynamiquement stables sont [Fe2+(C20.h)2- et [Fe3+(C20.)3f-. Les ions Fe2+non complexes ne peuvent etre identifies qu'en solution fortement acides, tandis que la presence d'ionsFe3+ non complexes restent peu probable en solution d'acide oxalique. Dans une gamme de pH de 1-2, lesions [Fe3+(C20.h)- et [Fe3+C20.)+ sont stables, tandis qu'a pH < I, la seule forme ionique existante est[Fe3+HC20.f+ .

NOTATION

[H+), [C20/-), [HC20.), [H2C20.): Concentrations of each species inequilibrium

Cox: Analytical concentration of oxalic acid in molarity units(mol/I)

CF/+: Total concentration of trivalent iron in solution (grion/I)CF/+: Total concentration of bivalent iron in solution (grion/l)

I'1Go: Standard free energy of reaction (kJ)K,: Formation constant of complex ionsKd: Dissociation constant of complex ionsR: Ideal gas constant, 8.314J/K,molT: Absolute temperature, KK: Equilibrium constant

ao, a" a2: Fractions of non-dissociated H2C20., HC20.and C20/-ions, respectively.

Bo, B" B2'B3: Fractions of the total trivalent iron in solution, presentas Fe3+, FeC20:' Fe(C20.)2- and Fe(C20.)/-, respectively.

Ao, A" A2' A3' A.: Fractions of the total trivalent iron in solution,present as Fe3+, FeC20:' Fe(C20.)2' Fe(C20.)/- and[FeHC20.)2+, respectively.

Do, D" D2: Fractions of the total bivalent iron in solution, present asFe2+, Fe(C20.)/- and Fe(C20.)3.-, respectively.

INTRODUCTION

Industrial minerals such as kaolin, quartz and feldspar are oftencontaminated with impurities, particularly in the form of ironoxides, In order to reduce the iron content of industrial mineralshydrometallurgical processes are used. Dissolution of metal

363

oxides, especially iron oxides, is also an important process inseveral fields other than hydrometallurgy, such as cleaning ofmetal surfaces, passivity of metals, etc. The chemical treatmentin all these processes is practised with aqueous solutions ofmineral or organic acids.

The dissolution of pure iron oxides by means of organicacids has been extensively studied [1-6]. The mechanism of irondissolution comprises two different chemical pathways [1,2, 7-I I]: (a) non-reductive dissolution and (b) reductive dissolution.

Among the different organic acids, oxalic acid is one of themost widely used for the dissolution of iron owing to its effec-tiveness as solvent reagent [2, 3, 6, 9].

In the present paper, a thermodynamic analysis of the abovepathways for the iron oxides-oxalic acid system is presented.For this purpose, the speciation of oxalic acid and the probableiron oxalato complexes as a function of pH and the standardfree energy change of some representative chemical reactionshave been evaluated.

SPECIATION OF OXALIC ACID AS A FUNCTION OFpH AT 25°C

Oxalic acid as a diprotic acid has two ionisation constants,given in Table I.

In oxalic acid solution, molecules of oxalic acid and its ions,

--

364

Table I. Ionisation constants of oxalic acid at 25°C [12-15]

D. PANIAS et af..: REACTIONS OF IRON OXIDES

H,C,04+-+H+ + HC,O..HC,O..+-+H+ +C,O~-

Kal = 5.6 X 10-'Ka' = 6.2 X 10-5

HC,Oi and C,O/-, coexist in equilibrium. These ions havebeen produced by the following ionisation reactions:

H,C,04+-+H+ + HC,Oi (2.1)

Kal_[H+][HC,Oi][H,C,04]

K _ [H+][C,O~-]a' [HC,Oi]

If Coxis the initial concentration of H,C,04 in the solution,then the fractions of non-dissociated H,C,04' HC,Oi andC,O/- are defined by the following equations, respectively:

[H,C,04]ao- Cox

and:

Cox= [H,C,04] + [HC,Oi] + [C,O~-] (2.8)Combining eqns (2.2) and (2.4)--{2.8) the fractions ao, a" a,

as a function of pH can be derived:

[H+]'

ao [H+]'+Ka.[H+] + KalKa, (2.9)

Kat[H+]a. [H+]'+Kal[H+]+ KalKa, (2.10)

KalKa,

a, [H+]'+Kal[H+]+Ka.Ka, (2.11)In Fig. I, the variations of the undissociatedH,C,04 and

(2.2)

dissociated fractions of HC,Oiand C,O~- as a function of pHare shown.

It is obvious from Fig. I that the C,O~- concentration is highin alkaline solutions and low in acidic solutions. At pH lessthan 2, the C,O~- concentration is negligible. In such solutions,the active species is HC,O~ - rather than C,O/- , indicating thatin the presence of ferric ions in the solution the [FeC,04H]'+complex ion is probably formed. The dissociation constant ofthat complexion at 25°Cis:kd = 2.95 X 10-10[13].

At pH higher than 3, almost complete ionisation of oxalicacid is observed. In this region, the active species are bothHC,Oi and C,O~-. Above pH 6 the concentration of HC,Oibecomes negligible.(2.3)

(2.4) IRON OXALATO COMPLEXES

The well-known phenomenon of ring formation by a ligandin a complex is called chelation and the ring formed is calleda chelate ring. Oxalate ion (C,O~-), as demonstrated in thefollowing formula:

(2.5)

[

:O-f=0

]

'-

:O-C=O(2.6)

(2.7)

has two oxygen atoms with unshared pairs of electrons. Theseoxygen atoms have the ability to co-ordinate to the same metalatom or ion and form a ring. In the presence of ferric (Fe3+)and ferrous (Fe'+) ions, oxalate ions have the ability to generatefive-membered rings and form complex ions. In Table 2, the

Table 2. Probable iron oxalato complexes and their overall dissociationconstants (kd) at 25°C

Complex ion Name

[Fe'+ (C,04h]4-[Fe'+(C,04),]'-[Fe3+(C,04)3f-[Fe3+(C,04),]-[Fe3+C,04]+[Fe3+C,04H]'+

6 X 10-6 [12-14]

2 x 10-8 [12-14]

3 x 10-'1 [12-14]

6.31 x 10-17 [13, 14]

3.98 x 10-10 [13, 14]

2.95 x 10-10 [13]

trioxalatoferrate(II) iondioxalatoferrate(II) iontrioxalatoferrate(III) iondioxalatoferrate(III) ionoxalatoiron(III) ionbioxalatoiron(III) ion

pHFig. I. Speciation in oxalic acid solution.

--

0.9

.... 0.8q. 0.7

::I: 0.6.....0 0.5'"c:: 0.40.....

0.3ucoI-<

..... 0.2

0.1

D. PANIAS el af..: REACTIONS OF IRON OXIDES

probable iron oxalato complexes and their overall dissociationconstants (kd) are given.

From Table 2 the following conclusions can be deduced:

(a) [Fe3+(C204)3P- is the most stable iron oxalato complex.(b) All iron(III) oxalato complexes are more stable than the

iron(II) oxalato complexes.

In the presence of oxalate ions, ferrous ion can form an insol-uble compound, FeC204, with solubility product value Ksp= 2X 10-7 [15]:

The solubility of ferrous ions in oxalate solutions is higherthan the theoretical value, which is calculated by the above Kspvalue, as a result of the formation of soluble triox-alatoferrate(H) and dioxalatoferrate(II) ions.

STANDARD FREE ENERGY, A(?, CALCULATIONS

Hematite-oxalic acid systemThe dissolution of iron from hematite in oxalic acid can

be described in two different pathways, reductive and non-reductive [8].

Non-reductive pathway

The non-reductive pathway involves the presence of iron inthe form of iron(III) oxalato complexes in the solution. Asreported in the previous section, the trioxalatoferrate(III) ion[Fe3+(C204)3P- is the most stable iron(III) oxalato complex.Thus, the following representative reaction can be assumed forthe non-reductive pathway:

-+2[Fe(C204)3]la~)+3H20(I) (4.1.1)

This reaction predominates at pH higher than 3.5. However,in high acid solutions, the active species are the HC20i ions, aspresented in Fig. I. Therefore, the complex ion [Fe3+C204Hf+is probably formed in the solution. Thus, in high acid solutionsthe following representative reaction can be assumed:

Fe203(S) + 2HC204(aq) + 6H(~q)

-+2[Fe3+HC204]la~)+3H20(l) (4.1.2)

The calculation of I:iG' is based on Hess's law. The values of

I:iGo for an appropriate series of reactions are given in Table 3.The I:iGoof reactions (4.1.1) and (4.1.2) are:

Table 3. l:1G' values for some reactions at 25°C

365

(3. I)

I:iG298(4.1.I)=-212.75kJI:iG298(4.1.2) = - 87.35kJ

The negative calculated value of I:iGois an indicator of aspontaneous chemical reaction, but is not in any way an indi-cator of kinetic stability or instability.

The I:iGovalues, given in Table 3, for reactions (4.1.4), (4.1.6)and (4.1.7) are taken directly from thermodynamic data tables[16, 17]. The I:iGovalue for reaction (4.1.5) is chosen by con-vention to be zero [16, 17].The I:iGovalues for reactions (4.1.3)and (4.1.8) have been calculated according to the followingequation:

I:iG298= -2.303RTlogKr

using the values of Krgiven in Table 2.

Reductive pathway

The reductive dissolution pathway can be characterised byan induction period describing the build-up of iron(II) oxalatocomplexes in the solution. This is a consequence of a redoxreaction taking place on hematite surface and can be describedby the following representative reaction:

Fe203(s) + 5C20~;q) + 6H(~q)-+2[Fe(C204)2]la~)

+2C02(g)+3H20(l) (4.1.9)

The [Fe(C204)2f- complex ion has been selected because itis the most stable iron(II) oxalato complex according to thedata presented in Table 2.

The I:iGOvalues for an appropriate series of reactions aregiven in Table 4.

Based on Hess's law, I:iGOof reaction (4.1.9) can be cal-culated:

I:iG298(4.1.9)= - 152.29 kJ

The value of I:iG298for reaction (4.1.9) indicates that thereaction is thermodynamically favourable.

The I:iGo values, given in Table 4, for reactions (4.1.11),(4.1.12), (4.1.14) and (4.1.15) are taken directly from ther-modynamic data tables [16, 17]. The I:iGovalue for reaction(4.1.13) is chosen by convention to be zero [16,17]. The I:iG'value for reaction (4.1.16) is taken by Outokumpu ther-modynamic data bank [17].The I:iGovalue for reaction (4.1.10)has been calculated according to the following equation:

I:iG298= -2.303RTlogKr

Table 4. l:1Govalues for some reactions at 25°C

Feta~) + 2C20~- [Fe(C20.)~ta~)

Fe203(.)--+ 2 Fe(.) + 3/202(1)

Fe(.)--+ Feta~)+ 2e-

H(~q) + e - --+ 1/2 H 2(1)

C(.) + 02(1) --+CO2(1)

H2(1)+ 1/202(1)--+H20(l)

C20~-;q)--+ 2C(.) + 202(1) + 2e-

l:1GO(kl/mol)

-43.93742.20

- 78.90o

- 394.36-237.13

851.28

Equationnumber

ReactionEquation Reaction

l:1Go(kl/mol) number

Fela~)+ 3C20~-;q) [Fe(C20')3]la~)2Fe(.) + 3/202(1)--+Fe203(.)1/2H2(g)--+ H(~q) + e-F e(s)--+F ela~)+ 3e-

H2~1)+ 1/202(1)--+H20(I)Feca:)+ HC204iaq) [Fe3+C20.H]2+

-117.08- 742.20

o-4.70

-237.13- 54.38

~I.~~I~~I.~~I~~I.~~I.~

(4.1.10)(4.1.11)(4.1.12)(4.1.13)(4.1.14)(4.1.15)(4.1.16)

_ .~ ~.UI..:REACTIONS OF IRON OXIDES

.. ~ I-Huposed chemical pathways, reductive and non-reductive, are thermodynamically spontaneous. Therefore, kin-etic parameters determine which mechanism predominates.

Goethite-oxalic acid system

As in the case of hematite-oxalic acid system, the dissolutionof iron can be described by two different chemical pathways,reductive and non-reductive [8].

Non-reductive pathway

In weak acidic solutions, the non-reductive dissolution isdescribed by the following reaction:

Fe203.H20(s) + 6C20~-;;q)+ 6H(~q)

->2[Fe(C204)3]la~)+4H20(I) (4.2.1)

In high acid solutions, the following reaction must beassumed as representative:

Fe203.H20(s)2HC20.j(aq) + 6H(~q)

->2[FeHC204]ta~)+4H20(I) (4.2.2)

Based on l:1G' values presented in Tables 3 and 5, l:1Govaluesof reactions (4.2.1) and (4.2.2) are calculated:

l:1G~98(4.2.1) = - 215.4 kJ

l:1G~98(4.2.2) = - 89.0 kJ

Reductive pathway

The reductive pathway of geotite dissolution with oxalic acidcan be described by the following reaction:

Combining the l:1Govalues presented in Tables 4 and 5, l:1Go)f reaction (4.2.4) is calculated:

l:1G~98(4.2.4) = -153.9 kJ

<.:QUILIBRIUM DIAGRAMS OF IRON(III) OX ALA TO

COMPLEXES (AT 25°C)

1function ofC20~- concentration

he following species are present in an aqueous solution of-C20/- system: Fe3+, [FeC204]+, [Fe(C204)2]- and::204)3f- .pH less than 3, the formation of hydroxoiron(III) com-; (Fex(OH)Yx-y can be excluded. Chemical equilibrium

Table 5.!1G' values at 25'C [17]

Equation!1G0(kJjmol) number

- 740.0 (4.2.3)

Table 6. Chemical reactions and equilibrium constants descr.Fe3+-C20l- system at 25' C

Reactions Equilibrium co,

Fe'+ +C20~ [FeC20.]+[FeC20.]+ +C20~- [Fe(C20.)2]-[Fe(C20')2]- + C20~- [Fe(C20.hP-

K,=2.51 x IK2=6.31 x \IK,=2.1 X 10'

in this system can be described by the three reactions giveTable 6.

The equilibrium constants Kb K2 and K3 of these reactiare:

K _ [FeC20:]I - [Fe3+][C20~-]

[Fe(C204)i]K2=

[FeC20:][C20~-]

K3_ [Fe(C204)~-][Fe(C204)i][C20~-]

The total concentrationof trivalent iron in solution,CFe3+,is:

(5.1

(5.L

(5.1.3,

CFe3+= [Fe3+]+ [FeC20:] + [Fe(C204)i] + [Fe(C204)~-]

(5.1.4)

The fraction of the total trivalent iron in solution, present asFe3+, FeC20:' Fe(C204)2- and Fe(C204)/-, respectively, canbe defined as:

(5.1.5)

BI = [FeC20:]CFe,+

(5.1.6)

B2 = [Fe(C204)2]CFe,+

(5.1.7)

B [Fe(C204)~-]3 CFe,+

Combining eqns (5.1.1)-(5.1.8) the fractions Bo, Bb B2 andB3 can be calculated as a function of C20/- concentration inthe solution:

(5.1.8)

IBo

I + kl[C20~-] + k1k2[C20~-f + k1k2k3[C20~-f

BI=Bokl[C20~-]

B2= Boklk2[C20~-f

B3= Bok,k2k3[C20~-f

The fractions of total iron (Bo, Bb B2, B3) as a function ofC20/- concentration in the solution are plotted in Fig. 2.

As seen from Fig. 2, free trivalent iron Fe3+ ions are stable inoxalate solutions only at extremely low oxalate concentrations(below 10-7 M). Therefore, it is highly unlikely to find freetrivalent iron ions in common oxalate solutions. Above oxalateconcentration 10-6 M, the complex ion [Fe(C204)3f- is formed

D. PANIAS et al..: REACTIONS OF IRON OXIDES 367

2 4 6 8 10 12 14

-log[C2042-]Fig. 2. Fractions of total iron as a function of C20~- concentration in the solution.

and is the only significant complex ion existing in the solutionat concentrations above 10-2 M.

As afunction of pH and oxalic acid concentration

In the Fe3+-oxalic acid system, the assumed equilibrium reac-tions can be described by equations in Tables 1 and 6 togetherwith the following equation:

Fe3++ HC20i -[FeC204Hf+

K4=3.39 X 109

where: Kal and Ka2 are defined in eqns (2.2) and (2.4), respec-tively; KI> K2 and K3 are defined in eqns (5.1.1), (5.1.2) and(5.1.3), respectively; and

6.E-8+

~ 5.E-8

2.5

2.5

pH

Fig. 3. The fraction of total iron present as Fe3+ as a function of pH atoxalic acid concentration 1,0.1 and 0.01 M, respectively.

(5.2.1)

The analytical concentration of oxalic acid, CoX>in the solu-tion is:

Cox = [H2C204]+ [HC20i] + [C20~-]

+ [Fe(C204)+] + 2[Fe(C204)2] + 3[Fe(C204H-] (5.2.2)

Assuming that the oxalic acid is in large excess in the solution,which is a valid assumption in every experimental work, thespeciation of oxalic acid is not affected by the presence of ferricion. Thus, the concentration of iron(III) oxalato complexes isnegligible in proportion to the concentrations of oxalic aciddissociation products. As a result, eqn (5.2.2) is simplified tothe following:

(5.2.3)

The total concentration of ferric ion in the solution is equalto:

CFeH = [Fe3+]+ [FeC20.t] + [Fe(C204)2]

+ [Fe(C204)~-] + [FeHC20~+] (5.2.4)

Combination of eqns (2.2), (2.4) and (5.2.3) results in eqn(5.2.5) from which the concentration of C20/- as a functionof Coxand pH can be calculated:

(5.2.5)

As in the previous paragraph, the fractions of the differentiron ions in the solution can be expressed by the followingequations:

(5.2.6)

A._[FeC20tJC FeH

(5.2.7)

0.9

0.8

Q) 0.7

4-0 0.60en

0.5c::0......

0.4....utI:I

0.3I-<4-0

0.2

0.1

00

368 D. PANIAS et at..: REACTIONS OF IRON OXIDES

3 3.5 4

pHFig. 4. The fraction of total iron present as [FeCzO.]+ as a function of pH at oxalic acid concentration ],

0.1 and 0.0] M, respectively.

(5.2.8)

(5.2.9)

(5.2.10)

Combining eqns (2.4), (5.1.1)-(5.1.3), (5.2.1) and (5.204), thefollowing equations are derived:

low, below 6 x 10-7 M, in all cases. There are no Fe3+ions insolutions with pH above 2. The variation of the [FeCzO.]+complex fraction in the solution is shown in Fig. 4. Its con-centration is very low at all oxalic acid concentrations. It raisessteeply until pH I and falls down sharply until pH 3. AbovepH 3, its concentration is essentially zero.

In Fig. 5, the fraction of [Fe(CZ04)z]-, [Fe(CZ04)3P- and[FeHCz04]z+ in solution as a function of pH is presented. Atall oxalic acid concentrations, the fraction of [Fe(CzO.h]- inthe solution is below 004 and gradually decreases as the con-centration of oxalic acid is increased. The pH region in which[Fe(CzO.)z]- is stable strongly depends on the concentration ofoxalic acid. At 0.01 M, [Fe(CzO.h]- is stable between pH 1-4,

(5.2.12)

(5.2.13)

(5.2.14)

(5.2.15)

Equations (5.2.5) and (5.2.11)-(5.2.15) describe math-ematically the influence of Coxand pH on the composition ofthe Fe3+-oxalic acid system. The equilibrium composition ofthe Fe3+-oxalic acid system is a function of both pH and oxalicacid concentration. The above set of equations was used tostudy the effect of pH on the composition of the solution at I,0.1 and 0.01 M oxalic acid concentration. The results are shownin Figs 3-5, respectively.

The concentration of Fe3+ in the solution as a function ofpH is shown in Fig. 3. The Fe3+ concentration is extremely

(5.2.11)

while at I M oxalic acid concentration, it is stable betweenpH 0.5-2. The [FeHCzO.]2+ concentration in solution sharplydecreases as pH increases and it is almost zero at pH higherthan 2.5. Its stability region depends on the oxalic acid con-centration in solution. At very low pH, the [FeHCzO.f+ is thepredominant complex ion present in solution.

EQUILIBRIUM DIAGRAMS OF IRON(II) OXALATOCOMPLEXES (AT 25°C)

As afunction ofC20~- concentration

Chemical equilibrium in the system Fez+-CzO.z- can bedescribed by the chemical reactions presented in Table 7.

Assuming that the solution is not saturated in FeCzO., reac-tion (6.1.1) does not take place. Consequently, the chemicalequilibrium in the system is described by reactions (6.1.2) and

---

0.0016

+ '<t 0.00140

('.

U 0.0012c...

t.l.'" 0.001

V 0.0008t.l.<+-0 0.0006I:

... 0.0004:..J:;j.... 0.0002<+-

D. PANIAS et al..: REACTIONS OF IRON OXIDES 369

1M

2

pH

3 4

pHFig. 5. The fraction of total iron present as [Fe(CP4W-, [Fe(C204hr- and [FeHC204)]2+, as a function of

pH at oxalic acid concentration 1,0.1 and 0.01 M, respectively.

Table 7. Chemical reactions and equilibrium constants describing thesystem Fe2+-C20/- at 25°C

(6.1.3). The equilibrium constants Ks and K6 of these reactionsare given by the following equations:

Equilibrium Equation F C 0 2-Reaction constants number K = [ e( 2 4h ] (6.1.4)

5 [Fe2+][C20~-fFeC204(s) Fe2+ +C20~- Ksp = 2 X 10-7 (6.1.1)

Fe2++2C20~ [Fe(C204hf- K,= 5 x 107 (6.1.2) [Fe(C 0 )4-][Fe(C204)2]2-+C20~ [Fe(C204h]4- K6=3.33 x 10-3 (6.1.3) K6 2 43 (6.1.5)

[Fe(C204)~-][C20~-]

- ---

£ 0.8

.....0 0.6ens::0 0.4..........Uco 0.2s...

.....

0

0

O.1M

d) + "-0.8 FeHC 0 2+

.....-

..... 2 40 0.6ens::0 0.4..........uco 0.2s...

.....

0

0 1 2 3

pH4

O.01M1

£ 0.8 2+FeHC204

.....0 0.6ens::0 0.4

'puCO 0.2s...

.....

0

0 1 2 3 4

370 D. PANIAS et al..: REACTIONS OF IRON OXIDES

2

-log[C20/]Fig. 6. The composition of iron(II) oxalato complexes in an unsaturated Fe2+ solution as a function of

C20/- .

1

0.90.8

0.7

0.6

~ 0.5.8 0.4t) 0.3~

J:: 0.20.1

oo 2

The total iron concentration in the solution, CFeH,is equalto:

Do, DJ and D2 express the fractions of total iron present inthe solution as Fe2+, Fe(C204)/- and Fe(C204)34-, respectively.Combining the eqns (6.1.4)-(6.1.6), Do, DI and D2 can be cal-culated:

I

I + K5[C20~-j2+ K5K6[C20n3(6.1.7)

The composition of iron(II) oxalato complexes in an unsatu-rated Fe2+ solution as a function of (C204)2- is presented inFig. 6. In solutions with very low (C204)2- concentration (less

6 8

3 4

than 10-5 M) it is highly probable to identify uncomplexed Fe2+ions. Their concentration in the solution decreases rapidly as(C204)2- concentration increases from 10-5 to 10-3 M. Below10-3 M, uncomplexed Fe2+ ions are unlikely. In that area, theonly stable complex is the [Fe(C204h2-j. The [Fe(C204)/-jcomplex also builds up in the same area, but not at comparableconcentrations in the solution and thus can be neglected.

As afunction of pH and oxalic acid concentration

The Fe2+-oxalic acid aqueous system unsaturated in Fe2+ isdescribed by reactions (6.1.2) and (6.1.3). Assuming that oxalicacid is in large excess in the solution, the concentration ofC20/- in the solution can be calculated by using only theionisation reactions of oxalic acid. As previously demonstrated,C20/- concentration is given by the following equation:

(5.2.5)

As previously presented, the fraction of total iron present inthe solution as Fe2+, Fe(C204h2- and Fe(C204)/- is given by

-- --

0.0035

0.003+N

Q) 0.0025

'+-<0.0020

'"= 0.0015.8.....u

0.001....

'+-<

0.0005

DJ[Fe(C204H-j

DoK5[C2O-j2 (6.1.8)CFeH

D2 [Fe(C204);-j = DoK5K6[C2O-F (6.1.9)CFe2+

- - ---

D. PANIAS et at..: REACTIONS OF IRON OXIDES 371

pHFig. 7. The fraction of total iron present as Fe2+ and [Fe(C20.hf- as a function of pH.

-- --

+ 0.9N

Q,) 0.8'" 0.7tIS

a5' 0.6

'- 0.50'" 0.4c::

.8 0.3...u 0.2tISI-.'- 0.1

00 1 2 3 4

pH

,1N

,.......,N

C5 0.9U 0.8'-'"Q,) 0.7

.......'" 0.6tIS

+ 0.5NQ,)

0.4'-0 0.3'"c:: 0.20....... 0.1utISI-. 0'-

0 1 2 3 4

372 D. PANIAS et a/..: REACTIONS OF IRON OXIDES

~ 0.00035'",-.,~o 0.0003N

U'Q)' 0.00025~'"-'

~ 0.0002+NQ) 0.00015~

"E 0.0001c::.~ 0.00005....u'"

ct 0o 2 4 6 8

pH

pHFig. 8. The fraction of total iron present as [Fe(C204)34-j as a function of pH.

eqns (6.1.7)-(6.1.9), where Ks and K6 are the chemical equi-librium constants given in Table 7.

The composition of a Fe2+-oxalic acid system unsaturatedin Fe2+ as a function of pH and oxalic acid concentration isshown in Figs 7 and 8. At all oxalic acid concentrations (1, 0.1and 0.01 M) the fraction of Fe(C204)34- in the solution is verylow (less than 0.0035) and gradually decreases as the oxalic acidconcentration increases. The role of that ion in Fe2+ com-plexation in the solution is practically insignificant. In high acidsolutions, uncomplexed Fe2+ is the stable ion in solution. Thestability region is strongly dependent on oxalic acid concen-tration. At I M oxalic acid solution, Fe2+ is stable below pH0.5. As the oxalic acid concentration decreases the stabilityregion is extended to higher pH values. In low acid solutions(pH higher than 2), [Fe(C204h]2- is practically the only ionpresent. The stability region moves to higher pH values as theoxalic acid concentration decreases.

CONCLUSIONS

. Equilibrium diagrams show that in low acid solutions (pHhigher than 3) the only thermodynamically stable complexions of bivalent and trivalent iron are [Fe2+(C204)2]2- and[Fe3+(C204)3P-, As the oxalic acid concentration increasesthe stability region of the above complex ions is extended tomore acidic regions.

. Uncomplexed Fe2+ ion can be identified only in high acidsolutions, while uncomplexed Fe3+ ion is not likely to buildupin oxalic acid solutions.

. [Fe3+(C204)2]- is stable in the pH range 1-2, while[Fe3+C204]+ is stable in the same region but with incom-parable concentrations with the other complex ions.

. At pH less than I, the [Fe3+HC204]2+ complex ion is verystable in comparison with the other ions in the solution. Thus,in such solutions [Fe3+HC204f+ is the only ion existing.

~ --

'" 0.0035?

0.003U'-'

0.0025'"-''" 0.002'"

+0.0015NQ)

<+-0 0.0010c:: 0.0005.....U 0'".... 0 2 4 6 8<+-0

D. PANIAS et al..: REACTIONS OF IRON OXIDES

Acknowledgements-The financial support of the European Com-mission within the framework of the Brite-Euram II Program (ContractNo. BRE2-CT92-0215) is gratefully acknowledged.

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