Solutions of Hematite

Suspensions of hematite have also been used and studied for other aims than photooxidation of water, e.g. catalytic oxidation of sulphur dioxide in aqueous solutions [52]. Aqueous dispersion of semiconductor particles could be an easy and attractive way to photooxidise water, but they have the drawback that dihydrogen and dioxygen are produced simultaneously in the same suspension. Apart from the separation problem the two produced gases may create a pathway for back reactions that reduces the yield of the overall photo-oxidation process. The latter obstacle can partly be avoided by addition of Na2C03, which was successfully shown by Arakawa et al [115]. 

---

To produce such aggregates, a solution of hematite was prepared (with a well ultrasonified colloid sol), the solution chemistry adjusted to the desired values and salt added. This solution was left to aggregate overnight in a 2 L container and stirred at 270 rpm. Most experiments were carried out at pH 3 at various KCl concentrations. 

Iron Oxide Reds. From a chemical point of view, red iron oxides are based on the stmcture of hematite, a-Fe202, and can be prepared in various shades, from orange through pure red to violet. Different shades are controlled primarily by the oxide s particle si2e, shape, and surface properties. Production. Four methods are commercially used in the preparation of iron oxide reds two-stage calcination of FeS047H2 O precipitation from an aqueous solution thermal dehydration of yellow goethite, a-FeO(OH) and oxidation of synthetic black oxide, Fe O. ... 

Transparent red iron oxide is composed mainly of hematite, a-Ee202, having primary particles about 10 nm. It is prepared by a precipitation reaction from a dilute solution of an iron salt at a temperature around 30°C, foUowed by a complete oxidation in the presence of some seeding additives,... 

Figure 7. U(VI) sorption onto muscovite (7a, Schmeide et al. 2000) and hematite (7b, Lenhart and Honeyman 1999) in the absence (U) and in the presence of humic acid (U+HA). 7a [U02 ] = 1 pmol/L, [HA] = 5 mg/L, muscovite content of about 1.2g/L. Complexation of U with HA in solution and onto mineral surface may influence U sorption. For instance, U sorption onto muscovite is enhanced in presence of HA at low pH. 7b [U] = 1 jamol/L, [HA] = 10 mg/L. Hematite content in solution = 0.9 and 9g/L. Uptake of U increases with increasing hematite content. In presence of hematite, an increase of U sorption onto hematite is observed at low pH, especially at low hematite content.

To see how we can use the surface complexation model to trace the kinetics of this reaction, we simulate an experiment conducted at pH 7.5 (Liger et al, 1999, their Fig. 6). They started with a solution containing 100 mmolar NaNC>3, 0.16 mmolar FeS04, and 0.53 g l-1 of hematite nanoparticles. At t = 0, they added enough uranyl to give an initial concentration of 5 x 10-7 molar, almost all of which sorbed to the nanoparticles. They then observed how the mass of sorbed uranyl, which they recovered by NaHCC>3 extraction, varied with time. 

Certain acids with hydroxylic and carboxylic groups have been shown (Schwert-mann and Cornell, 1991) to induce in Fe(HI) solutions the formation of hematite because these acids may act as templates for the nucleation of hematite. These examples illustrate that a complete understanding and quantitative description of the rate of heterogeneous nucleation will have to include surface complexation and other adsorption processes. 

The surface charge of metal oxides (due to surface protonation) as a function of pH can be predicted if their pHpzc are known with the help of the relationship given in Fig. 3.4. Fig. 7.6 exemplifies the effect of various solutes on the colloid stability of hematite at pH around 6.5 (pH = 10.5 for Ca2+ and Na+) (Liang and Morgan, 1990). 

The Rate of reductive Dissolution of Hematite by H2S as observed between pH 4 and 7 is given in Fig. 9.6 (dos Santos Afonso and Stumm, in preparation). The HS" is oxidized to SO. The experiments were carried out at different pH values (pH-stat) and using constant PH2s- 1.8 - 2.0 H+ ions are consumed per Fe(II) released into solution, as long as the solubility product of FeS is not exceeded, the product of the reaction is Fe2+. The reaction proceeds through the formation of inner-sphere =Fe-S. The dissolution rate, R, is given by... 

Bruno, J., W. Stumm, P. Wersin, and F. Brandberg (1991), The Influence of Carbonate in Mineral Dissolution. Part 1, The Thermodynamics and Kinetics of Hematite Dissolution in Bicarbonate Solution at T = 25° C", in preparation. 

The procedure for separating Sb-119 from an alpha-irradiated tin target has been described elsewhere (10,11). The amounts of cobalt and antimony coexisting with the nuclides are estimated to have been about 400 ng/mCi and 300 ng/mCi, respectively, i.e., to have been much smaller than that required for monolayer coverage of 30 mg of the hematite sample. About 10 cm3 of an aqueous solution containing 1 - 2 mCi of divalent Co-57 or 0.1 - 1 mCi of pentavalent Sb-119 was adjusted to an appropriate pH value in a Teflon vessel with a 0.5 mm-thick Teflon window at the bottom, and about 30 mg of hematite powder was added to the solution. The suspension was shaken for 30 min at room temperature. After settling of the powder at the bottom of the vessel, the pH was remeasured. 

In Situ Mossbauer Measurement on Hematite/Divalent Co-57. The adsorption behavior of cobaltous ions on hematite surfaces was essentially the same as that on silica reported by James and Healy (12). Appreciable adsorption begins at about pH 4 followed by an abrupt increase in adsorption between pH 6 and 8. Beyond pH 9, adsorption is practically complete. Emission Mossbauer spectra of Fe-57 arising from the divalent Co-57 ions at the interface between hematite particles and the 0.1 mol/dm3 NaCl solutions of different pH at room temperature are shown in Figure 3 The emission spectra show a marked dependence on the pH of the aqueous phase. No emission lines ascribable to paramagnetic iron species are recognized in... 

Adsorption of Pentavalent Sb Ions on Hematite. So far as we know, there are no experimental data on the adsorption equilibrium of dilute pentavalent Sb ions on metal oxides. Therefore, the pH dependence of the adsorption of pentavalent Sb ions on hematite was measured. Carrier-free pentavalent Sb-119 ions were adsorbed on 30 mg of hematite (prefired at 900°C for 2 hours) from 10 cm3 of 0.25 mol/dm3 LiCl solutions at 24 1°C. The amount of antimony employed in each run is estimated to be about 50 ng. The adsorption proceeds with a measurable rate and attains an apparent equilibrium after shaking for several hours. The reaction is second order with respect to the concentration of pentavalent Sb ions in the solution (13) The values given in Figure 4 are those obtained after 22 hours equilibration. As seen in Figure 4, strong adsorption of pentavalent Sb ions is observed below pH 7, while the percent adsorbed diminishes abruptly above that. Most of the Sb ions adsorbed on hematite from solutions of pH 2-5 are not desorbed by subsequent adjustment to alkaline conditions. Results on desorption of Sb ions pre-adsorbed at pH 4 are shown in Figure 4. 

Figure 4. pH dependence of the adsorption and desorption of carrier-free pentavalent Sb—119 on hematite at room temperature (30 mg of hematite prefired at 900°C in 10 cm3 of 0.25 mol/dm3 LiCl solutions). Desorption was measured on pentavalent Sb-119 pre-adsorbed at pH 4. Shaking time was 22 hours for the adsorption and was 5 days for the desorption. 

Sb carrier ions. The Sb-119 ions were adsorbed on 30 mg of hematite from 10 cm3 of a 0.25 mol/dm3 KC1 solution containing about 1 mg of pentavalent Sb ions. About 0.3 mg of Sb was adsorbed at pH 2.5 and 4.0. The amounts of Sb adsorbed are less than that required to cover all the hematite surfaces as a monolayer. The emission Mossbauer spectra obtained are shown in Figure 7. It is seen from Figure 7 that the width of the emission Mossbauer spectrum at pH 2.5 is much smaller than that of the carrier-free one, while essentially no effect of carrier Sb ions is observed at pH 4.0. 

Effects of Pentavalent Sb Ions on the Adsorption of Divalent Co-57 on Hematite. Benjamin and Bloom reported that arsenate ions enhance the adsorption of cobaltous ions on amorphous iron oxyhydroxide (J 6). Similarly, when divalent Co-57 ions were adsorbed on hematite together with pentavalent Sb ions, an increase of adsorption in the weakly acidic region was observed. For example, when 30 mg of hematite was shaken with 10 cm3 of 0.1 mol/dm3 KC1 solution at pH 5.5 containing carrier-free Co-57 and about 1 mg of pentavalent Sb ions, 95 % of Co-57 and about 30 % of Sb ions were adsorbed. The emission spectra of the divalent Co-57.ions adsorbed under these conditions are shown in Figure 8 together with the results obtained under different conditions. As seen in Figure 8, the spectra of divalent Co-57 co-adsorbed with pentavalent Sb ions are much different from those of Co-57 adsorbed alone (Figure 3). These observations show a marked effect of the.co-adsorbed pentavalent Sb ions on the chemical structure of adsorbed Co-57. 

Figure 12. Extent of dissolution and re-precipitation between aqueous Fe(III) and hematite at 98°C calculated using Fe-enriched tracers. A. Percent Fe exchanged (F values) as calculated for the two enriched- Fe tracer experiments in parts B and C. Large diamonds reflect F values calculated from isotopic compositions of the solution. Small circles reflect F values calculated from isotopic compositions of hematite, which have larger errors due to the relatively small shifts in isotopic composition of the solid (see parts B and C). Curves show third-order rate laws that are fit to the data from the solutions. B. Tracer experiment using Fe-enriched hematite, and isotopically normal Fe(lll). C. Identical experiment as in part B, except that solution Fe(lll) is enriched in Te, and initial hematite had normal isotope compositions. Data from Skulan et al. (2002).

Adsorption Methods. Five grams of hematite were first conditioned in 0.001 M NaCl at pH 4.1. After the SDS had been added to the slurry and the pH adjusted as required, the samples were conditioned on a rotating shaker for two hours. The solutions were then centrifuged, and the supernatant liquid analyzed for its SDS content. The amount of SDS adsorbed was calculated as the difference between the initial amount added and the residual amount measured. Experimental results showed that two hours was sufficient time for equilibrium to be reached. Somasundaran ( ) observed similar equilibrium adsorption times for sulfonate adsorption on aluminum oxide. 

The structure derived from a Rietveld fit of a neutron diffraction pattern of a 6-line ferrihydrite which showed more and sharper lines (Fig. 2.9, lower) than an XRD pattern, was in agreement with the structure proposed by Drits et al. (1993) except that it was not necessary to assume the presence of hematite in order to produce a satisfactory fit (Jansen et al. 2002). The unit cell of the defect free phase had a = 0.29514(9) nm and c = 0.9414(9) nm and the average domain size derived from line broadening was 2.7(0.8) nm. Since forced hydrolysis of an Fe solution at elevated temperatures will ultimately lead to hematite, it is likely that incipient hematite formation may occur under certain synthesis conditions. Neither these studies nor Mbssbauer spectroscopy, which showed only a singular isomer shift at 4.2 K characteristic of Fe, supported the presence of " Fe (Cardile, 1988 Pankhurst Pollard, 1992). However, the presence, at the surface, of some Fe with lower (<6) coordination, perhaps as tetrahedra (Eggleton and Fitzpatrick, 1988) which may have become unsaturated on heating, has been suggested on the basis of XAFS results (Zhao et al. 1994). 

The value of log Kg4 = -19.64 + 0.06 at 60 °C and -14.82 + 0.14 at 300 °C (Diako-nov et al., 1999). Solubility rose with rising [OH ]. In very concentrated solutions of NaOH or KOH, ion pairs can form between Fe(OH)4 and the cation and at temperatures >200°C, these ion pairs increase the solubility of hematite in such solutions. In alkaline media, the solubility of hematite depended upon the alkali hydroxide used in the order (Ishikawa et al., 1997) ... 

Desorption of the reduced metal ion is the rate determining step and is assisted by protons and oxalate ions. The reoxidized surface complex also desorbs owing to its altered molecular structure and is thus available for further reaction. The reductive dissolution step is faster than the initial complexation process. Photochemical dissolution of hematite in acidic oxalate solution is faster when air is excluded from the system (by purging with N2) than when air is present (Taxiarchou et al. 1997). 

---