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Journal of the Chilean Chemical Society

versión On-line ISSN 0717-9707

J. Chil. Chem. Soc. vol.58 no.1 Concepción mar. 2013

http://dx.doi.org/10.4067/S0717-97072013000100023 

THERMODYNAMIC STUDY OF IRON IN ALLOYS OF QUASICRYSTALLINE AlCuFe

 

LUCIANO NASCIMENTO1*LOURDES CRISTINA L.AGOSTINHO2, BERNARDETE F.CAVALCANTI3

1* Civil Engineering Department-CED/UNIPÊ,Campus do UNIPÊ/BR-230-Km 22, Bloco-F, ÁguaFrio-CEP: 58.053-000, JoãoPessoa-PB,Brazil. e-mail: luciano.engenhariacivilunipe@gmail.com
2
Chemistry Department-CD/CST-UEPB, Rua Juvêncio Arruda, s/n-Bodocongó-CEP:58.109-790, Campina Grande-PB,Brazil.
3
Civil Engineering Academic Unity-CEAU/CTNR-UFCG, Av. Aprígio Veloso, 882-Bloco CQ-CEP: 58. 429-900, Campina Grande-PB,Brazil.


ABSTRACT

Solid iron can be observed at environmental conditions and also as metallic component of quasicrystalline alloys. Its behavior in aqueous solution is analyzed from some particular thermochemical relationships. The model utilizes free Gibbs energy to obtain pH and ñå parameters which are related to iron redox reactions. The solid iron system, defined as Fe : Fe2+ - Fe3+ - Fe2O3 (hematite) - Fe3O4 (magnetite), was modelled by considering a low molar concentration and environmental temperature. In order to obtain pε values, equations relating Gibbs standard free energy and equilibrium constants were applied. A stability diagram such as pe x pH was chosen to depict these thermochemical relationships. Thus, straight line dpε/dpH slopes were obtained for each process step of the solid iron system. In systems relating solid iron to iron species Fe2+ and Fe3+ the slopes are equal to zero; i.e. horizontal line shows the iron species dominance in some specific pH (hydronium ions) range. The negative values obtained show the pH dependence of the iron species investigated. In natural water environment, it is possible to distinguish hematite formation from magnetite formation. However in quasicrystalline AlCuFe alloys the last iron species formed is hematite which ion pairs to copper. Due to the strong hydronium ions (pH) dependence, it is necessary to measure it both in environmental conditions and quasicrystalline phase's formation. Considering that the temperature adopted was 25°C, the results are valid only for this temperature.

Keywords: Iron thermochemical model; Standard free energy; Solid iron species.


1.       INTRODUCTION

Corrosion is the main problem arising with iron in natural waters. Corrosion involves electrochemical reactions between, for instance, the water and the metallic component of the system. As a result, it can be observed pitting, nodule formation, and red water among other problems. With regard to quasicrystalline alloys, solid iron is used to prepare AlCuFe alloys. These atomic ordered structures exhibit quasiperiodic translational long -range order and were first reported in 1984 by SCHECHTMAN et al. [1]. Quasicrystals, show forbidden crystallographic symmetries and are prepared by fast cooling techniques [2]. Generally, it is used grinding at high energy for syntherizing metastables and stables quasicrystals via reaction of solid state [3]. After this process, chemical treatments are applied during some hours at temperatures above 300°C in environment or in other medium. Samples are then characterized by using X - ray photoelectron spectroscopy (XPS) and X -ray stimulated Auger electron spectroscopy (XAES) [4]. According to Figure 1 and the second PINHEIRO et al.[5], pure metals used to form quasicrystalline alloys suffer oxidation under several conditions; i.e., in vacuum, in normal air and after liquid immersion. By using XPS data and XAES spectra obtained in fixed angle, each element of the AlCuFe alloy was studied. They have observed aluminum enrichment on the quasicrystal surface and a decrease in copper and iron. They also observed that; (i) aluminum ion oxidizes under all conditions of alloys preparation; (ii) copper and iron are protected from oxidation when the alloy is immersed in water and (iii) this protection is due to the formation of an aluminum passivation layer. Also, they have observed that the 2ρ XPS line falls at 707 eV in pure metal indicating the presence of hematite FeOOH.


Figure 1. X - ray photoelectron spectroscopy (XPS) of AlCuFe alloy after several treatments according [5].

NASCIMENTO et al. [6] analyzed the oxidation behavior of AlCuFe alloy by applying different conditions of thermal treatment. They have used X - Ray Diffratometry technique and Scanning Electron Microscopy (SEM) to analyze also quasicrystal morphology and phases evolution. Regarding solid iron behavior in quasicrystal, their work showed similar results as those reported by ROUXEL et al.[7]. At temperatures above 700°C in the crystalline phase of AlCuFe, a nodule formation was observed as well as an association of copper with hematite (a solid named cuprospinel, CuFe2O4), that is the last compound formed until quasicrystal reaches its icosahedral phase. Figure 2 below, respectively, shows the result of sample Scanning Electron Microscopy in the sample quasicrystalline alloy AlCuFe, showing the geometric uniformity in the image of the microstructure with the formation of hematite (Fe2O3) in the alloy quasicrystalline.


Figure 2. Formation of Fe2O3 (hematite) in the alloy quasicrystalline AlCuFe.

Vacuum oxidation is relatively mild, air oxidation is severe and oxidation from liquid immersion is the most severe. Solubility of iron under different conditions at equilibrium can be conveniently calculated with stability diagrams such as Eh-pH diagram [8]. These diagrams utilize Nernst equations besides chemical and thermochemical relationships and standard free energies values from literature. Another way to analyze these interrelationships is calculating pε and pH (hydrogen potential) parameters for a particular aqueous system. In the same way that an acid is able to donate a proton and a base to accept a proton there are substances which have the capacity to donate electrons (reductors) and others to accept electrons (oxidants). For iron in the ferrous state occurs an oxidation reaction and there is formation of ferric iron species. Similarly, ferric iron as oxidant suffers a reduction reaction producing ferrous iron species. Solid iron Fe(5) used to prepare metallic alloys for the formation of quasicrystals is the most reduced state while ferric ionic species is the most oxidized. Briefly, the kinetic of oxygenation of iron (II) can be explained as follows: Iron (II) is oxidized by a rate of first order with respect to iron (II) and oxygen and of second order with respect to hydroxide ion [9]:

where, k isa constant equals to 2. 1013 min 1־.atrrf 1.mol-3 at 25°C; ( ) indicates active concentration and pp refers to partial pressure. Oxidation of Fe2+ is followed by hydrolysis (process that can proceed with water, hydrogen ion and hydroxyl ion) and oxidation to insoluble forms (ferrous oxides and others). This process is, indeed, very fast and is important for several chemical relations. This paper shows the development of a thermochemical model for solid iron system which encompasses the formation of oxidant and reduced iron species besides the formation of hematite and magnetite species. The objective is to show the thermochemical behavior of iron not only at environmental conditions, but also when it is applied in metallic alloys to produce quasicrystal after immersion in water.

1.1.     THEORETICAL CONSIDERATIONS

Chemical reactions spontaneously proceed towards an equilibrium state that is energetically most favorable. The quantitative description of equilibrium conditions is the subject of the field of chemical thermodynamics. In general, the reaction can be written by the following equation [10]:

In this equation, Bj designates the different species that participate in the reaction and Vj their stoichiometric coefficient. The subscripts ox and red indicate whether a species j corresponds to the "oxidized" or to the "reduced" state; in other words, whether it stands to the left or to the right of the equality sign of the reduction reaction (2).

Here E0 is the standard potential of the electrode reaction and the symbol Ð indicates the multiplication product of the activities of participating species to the powerVj; R is the universal gas constant (8,31 J.mol-1 K-1), T is the absolute temperature (K) and F is the Faraday constant (96480 Coulomb mol-1). pH is defined as the disponibility of hydronium ions in solution. In practice, pH is defined according to the equation of Linderstrom - Lang [11].

The above expression is valid for unity activity coefficient of hydrogen ion. The standard free energy for the formation of 1 mol of substance from pure elements is defined as the standard free energy of formulation. Its equation for a particular reaction is given by [12]:

ΔG°f values for almost all chemical substances of interest are well defined in literature at 25°C temperature.

In turn, the relationship between the equilibrium constant k' of, for instance, a monoprotic weak acid/base and the standard free energy change of reaction is given by:

For the values of R (gas constant) and absolute temperature at 25°C (T = 298,13K):

pε is related to the standard free energy by the equation:

pε is also related to the electrode potential E as follows [12]:

where, E° is the standard electrode potential; "red","ox" refer to reductant and oxidant, respectively. Values for standard potentials of redox process are listed in literature. In all thermochemical studies the standard free energy of formation of hydrogen ion from is ΔG°f of hydrogen elements equal to zero.

2.       THESI MODEL DEVELOPMENT

2.1     Initial Considerations

Thesi Model (Thermodynamic on Solid Iron) is based on equations for k', ΔG°r , pH and pε as given above. The model is applied to solid iron which forms the system Fe(5) = Fe3+ - Fe2+ - Fe2O3 - Fe3O4; i.e., iron in ferric and ferrous state as well as hematite and magnetite. In the present study it was considered for total solid iron a molar concentration of 10-6 mol/L which means a solution with 0,06 mg/L of iron. The main dissociation reactions of the system are as follows:

2.2     Analysis of the Iron Redox Systems

Each equation of the solid iron system is analyzed "per se" in order to determine values for the chemical equilibrium constant k' and for pε and pH parameters from Gibbs standard free energy of formation. For constructing a stability diagram such as pε x pH, the slope represented by is obtained from each particular reaction of the system. Table 1 shows the selected values of Gibbs standard free energy of formation for iron species at 25°C and one atmosphere of pressure.

Table 1. Values of Gibbs standard free energy of formation for iron species at 25°C and one atmosphere of pressure [10].

(1) Fe(5): Ferrous Iron System.

Equilibrium Constant for unity activity and pε value is given by:

Gibbs free energy for reaction of equation (11):

Hence:

log k' = -15. For [Fe2+] = 10-6M and k' value, pε = -10, 5.

The slope of the straight line for constructing a stability diagram is obtained from:

; i.e. in the stability diagram this line is horizontal at pε = -10, 5.

(2) Fe : Ferric Iron System.

Equilibrium Constant for unity activity is given by:

Gibbs free energy for reaction of equation (12):

Hence, the equilibrium constant k' is given by:

log k' = 1,85

Slope given by: ; i.e., line representing ferric ion is horizontal at pε = - 2, 6.

(3) Fe(s): Fe2O3 System.

Equilibrium Constant for unity activity and pε value are given by:

Gibbs free energy for reaction of equation (13):

Results obtained are:

log k'= - 6,6 and pε = - pH - 1,1.

Slope given by: . For pH = 0, then pε = -1,1.

(4) Fe : Fe3O4 System.

Equilibrium Constant for unity activity and pε value given by:

Gibbs free energy for reaction of equation (14):

Results obtained are:

log k'= - 11,4; Slope is given by: . For pH = 0 then pε = -1,4.

(5) Fe3+: Fe2+ System.

Initial equation for the system is given by:

Equilibrium constant k' defined in terms of activity is given by:

Gibbs free energy for reaction of equation (14):

Results for pε and slope of the straight line for stability diagram are given by:

(6) Fe2O3: Fe3O4 System

Equilibrium reaction for this system is given by:

Equilibrium Constant for unity activity and pε value are given by:

; i.e., activities of hematite and magnetite are equal to unity. Thus,

Gibbs free energy for reaction of equation (16):

Results given by:

(7) Fe2+: Fe2O3 System.

Equilibrium reaction of the system is defined as follows:

for activities of hematite and ferrous species equal to unity.

Gibbs free energy for reaction of equation (16):

The results are given by:

(8) Fe2+: Fe3O4 System. Equilibrium reaction given by:

k' = 1/(H+)8 (e-)2/(Fe2+)3

Gibbs free energy for reaction of equation (14):

Results for k' and pε :

log k' = - 33,25 ; pε = - 4pH + 25,6.

Slope is given by: . For pH = 0 then pε = 25,6

(9) Fe3+: Fe2O3 System:

This system is not a redox reaction but a weak acid reaction, the chemical equation is defined as follows:

Equilibrium constant for unity activity is given by:

Gibbs free energy for reaction of equation (19):

Results for k' and pH:

log k'= 2,9 and pH equation is now given by:

 indicating a very weak acid solution. In stage (5) of the model it was necessary to verify if Fe(s) is limited by ferrous species or by ferric species and also by Fe2O3 and Fe3O4. Table 2 shows all data determined for pk', pε, pH and slope for all the stages of the model here developed.


Table 2. Results obtained from the study the Fe(s) = Fe2+ - Fe3+ - Fe2O3 -Fe3O4 system with data of Gibbs free energy from literature at temperature of 25°C.


3.       DISCUSSION

This model shows the solubility of solid iron under a particular condition at equilibrium Results are shown graphically by means of pε versus pH stability diagram in Figure 3.


Figure 3. Stability diagram for solid iron in aqueous solution.

This diagram was prepared by using the Nernst equation and some thermodynamically relationships and data. For natural waters, it was not considered the influence of dissolved carbon dioxide, CO2, and species which are generally involved in precipitation of some solid forms of iron. In this figure it can be seen a strong dependence of the oxidation rate on pH (hydronium ions). The hydrolyzed Fe2+ and reacts more rapidly with O2 than nonhydrolized Fe2+. In the stability diagram, it seems that hydrolyzed Fe3+ species are not dominant, only in saline environment. Regarding the calculations, in stage (6) it can be seen that for dps = 0 , Fe3+-Fe(5) line falls away in the diagram. dpH =

In stages (7) to (8) it was studied the lines surrounding Fe2+ others than Fe3+ and Fe(s). In stage (10) it can be seen that the reaction is not a redox reaction but a weak acid reaction. Thus, in the stability diagram it will get a straight line separating Fe3+ and Fe2O3. An intermediate region of high solubility exists between the domains of pyrite and Fe(OH)2. This can be verified by means of Eh-pH diagrams in which Eh refers to the activity of electrons. Concerning solubility of iron species, the well - crystallized form as hematite and FeOOH are less soluble. Reoxidation of Fe2+ produces a solid ferric hydroxide as a colloid or on a surface.

4.       CONCLUSIONS

The following conclusions can be made:

❖      The interrelationship between chemical and redox equations represented by pk', pH and pε reflect the net effect of chemical and thermodynamic processes in oxidation analysis of solid iron;

❖      It was observed a strong pH influence in all equations developed in the study especially in the oxi-reduction equation of solid iron Fe(s) and ferrous ion, Fe2+. In the system defined by ferrous iron and hematite, pH appears to be on the acid side. For pH values above 8,0 iron precipitation is very fast. The dominance of ferrous iron is stronger than that of ferric iron;

❖      The influence of temperature was not analyzed. However it is very important for all stages of the model. Some researches have showed that during formation of AlCuFe metallic alloys at high temperatures and further immersion in water only iron is oxidized. The cooling process and leaching into acid solutions can create a severe oxidation state;

❖      Due to the strong pH (hydronium ions) dependence, it is necessary to measure hydronium ions (pH) and temperature increase and verify iron species formation after liquid immersion in both environmental condition and quasicrystalline phase's formation.

NOTATION

( ), [ ] or C = active and molar concentration (mol. L-1), respectively;
CT = total concentration of species for iron system in solution, mol. L-1;
e- = ε = electron;
E = electrode potential;
F = Faraday constant;
fH+ = activity coefficient for the hydrogen ion, H+ in the molar scale;
ΔG°r = total free standard energy or Gibbs free energy of reaction, kg cal mol-1
ΔG°f = free energy of formation, kg cal.mol-1;
k' = dissociation Constant for a monoprotic weak acid/base;
R = universal gas constant= 1, 987 cal mol-1kg;
T = absolute temperature = T(°C) + 273,15 (K).

ACKNOLEDGEMENTS

The authors thank CNPq (National Counsel of Technological and Scientific Development) for financial support of this work.

 

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(Received: June 11, 2012 - Accepted: October 20, 2012)

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