2.4. Hematite Surface Speciation
2.4.3. Adsorption of Anions on Hematite
The adsorption of inorganic and organic ions alters the hematite surface charge and potential characteristics. Consequently, the stability of hematite is expected to depend on these adsorption processes. In this study, the following anions are investigated: phosphate , sulfate, chloride, and anions of carboxylic groups associated with naturally occurring organic matter. The heterogeneous properties of natural organic matter make it difficult to interpret and model the adsorption on hematite.
As a first approximation, it is assumed that the adsorption characteristics of low molecular weight carboxylic acids resemble those of natural organic matter. Results from the analysis of carboxylic acid adsorption are applied to interpret the extent of adsorption of natural organic matter on hematite particles.
The complex formation constants of different anions with surface groups are found to follow the same trend as those for the formation of anion-metal complexes in solution (Sigg and Stumm, 1980). Assuming that a similar analogy applies to hematite surfaces, the equilibrium constants for an anion surface complex can be estimated
Figure 2.4: The surface charge, σo, on hematite particles plotted as a function of pH for differing ionic strengths.
Figure 2.5: Hematite surface density as a function of pH as determined by a surface equilibrium model. The model parameters are given in the text.
from the equilibrium constants for solution complexes. If the equilibrium reaction for solution complexes is written as
FeOH2+ + HL ≠ FeL2+ + H2O (2.38)
The corresponding surface complex formation is represented by
≡FeOH + HL ≠= ≡FeL + H2O (2.39)
Table 2-1 is a list of published equilibrium constants for iron complexes of various anions in solution. The reactions are of the following form
Fe3+ + HmLm~n FeHmL3+m~n (2.40)
This equation can be transformed to
FeOH2+ + HnL FeHmL3+m-n + (n - m)H+ + H2O (2.41) For example, complex formation of iron and H2POJ^ has an equilibrium constant ιo3∙θ5,
Fe3+ + H2PO4 ≠= FeH2PO42+ Log∕< = 3.65 (2.42) To obtain the equilibrium constant for reactions of the form of Eq. (2.39) the following equilibrium reactions are used,
Fe3+ + H2PO4 r=÷ FeH2PO42+ LogX = 3.65
FeOH2+÷=i Fe3++ OH~ -11.81
(2.43)
H3PO4 ≡= H2PO4-+ H+ -2.15
H+ + OH^^ H2O 14.0
Summing the above equations yields
FeOH2+ + H3PO4 ^=i FeH2PO42+ + H2O Log∕< = 3.69 (2.44) The corresponding surface complex reaction is:
≡FeOH + H3PO4 ≡FePO4H2 + H2O (2.45)
(a) Iron Complexes.
Anions (L) Equilibrium Constant, K Log K
Nitrate (NO3) FeL/Fe ∙ L 0.76
Perchlorate (CIO4 ) FeL/Fe ∙L 1.15
Chloride (Cl-) FeL/Fe ∙ L 1.48
Acetate (CH3COO~) FeL/Fe ∙ L 4.00*
Sulfate (SO^-) FeL/Fe ∙L 4.04
Propionate (CH3CH2COO-) FeL/Fe∙ L 4.29*
Oxalate (C2O4~) FeHL/Fe· HL
FeL/Fe∙L
4.35*
7.74*
Phthalate (CsH6θ4^^) FeL/Fe ∙L N/A
Salicylate (C7H6θ3-) FeHL/Fe ∙HL
FeL/Fe ∙ L
5.06*
17.57
Fluoride (F^) FeL/Fe ∙ L 6.00*
Phosphate (PO4-) FeH2L∕Fe∙H2L
FeHL/Fe ∙ HL
3.65*
10.15*
(b) Other Metals and Hydroxide Ions (OH ).
Sodium (Na+) NaOH∕Na∙OH -0.20
Potassium (K+) KOH∕K∙OH -0.50
Calcium (Ca2+) Ca(OH)∕Ca∙OH 1.30
Magnesium (Mg2+) Mg(OH)∕Mg∙OH 2.58
Iron (Fe3+) Fe(OH)∕Fe ∙ OH 11.81
(c) Acidity Constants of Organic Ligands.
Acetic acid pKα = 4.75
Propionic acid pKa = 4.87
Caprylic acid pKa = 4.89
Capric acid pKa = 4.85
Aspartic acid pK1=2.0, pK2=3.9, pK3=10.0
Oxalicacid pKι=1.25, pK2=4.27
Phthalic acid pKι=2.95, pK2=5.41
Salicylic acid pK1=2.97,pK2=13.74
* Constants are extrapolated to 1=0 M usingthe Davies equation.
Table ‘2.1: Equilibrium constants for the formation of aqueous complexes from metals and ligands (from Smith and Martel (1976), and Perrin (1979)).
Assuming the surface complex formation constant for =FePO4H2 has a similar order of magnitude as the aqueous formation constant (Log∕<=3.69) for FeH2PO42+, the equilibrium constant in terms of Log/i is of the order of 4. According to Table 2-1, the strength of complex formation between iron(III) and negatively-charged ions is ordered as follows: HPO2- > F-^ > Salicylate > H2PO^" > Propionate ~ Acetate >
SO^" > Cl“ ~ C1O^ > NO^.
Equilibrium constant information is not available for phthalate-iron complexes.
To estimate the complex formation constant, it is assumed that the strength of the iron-phthalate complex lies between those of iron with oxalate and salicylate ions.
The estimated value for the iron-phthalate complex is thus approximately 106.
Chloride ions form a very weak aqueous complex with iron(III) (Logl< ~ 1.5). It is possible that chloride ions have a weak specific interaction with hematite surface hydroxyl groups. Experimental results by Breeuwsma (1973), however, did not show specific interaction of chloride on hematite surfaces. Data in Fig. 2.4 of the current research also indicate that the chloride ion is non-specific, as the three data sets all intersect at the same point, and no shift of pHzpc occurs with a change in electrolyte concentration of near three orders of magnitude. Hence, it is expected that chloride is not adsorbed on hematite through specific chemical interaction.
Phosphate forms the strongest complex among all the anions considered here.
Electrophoretic mobility data show that concentrations of phosphate as small as 10μM are capable of shifting the isoelectric pH point of a hematite suspension (~30 mg∕l) from 8.5 to 6.5 (see Fig. 4.20).
The equilibrium adsorption of phosphate was calculated with a diffuse layer model using the SURFEQL computer code. Fig. 4.26 illustrates the predicted total adsorp
tion of phosphate in the pH range of 4 to 10; the total adsorption of phosphate is compared with the experimental results.
In the model calculations, the surface acid-base equilibrium constants, Kla⅛t and Kza2t, are taken to be the same as in Fig. 2.5, without phosphate present. Additional constants for phosphate complexation with the surface are estimated based on the values given in Sigg (1979) and constants for the corresponding solution complexes.
The reactions considered are
≡FeOH + H3PO4 ≡FeH2PO4 + H2O (2.46)
≡FeOH + H3PO4 ≡FeHP07 + H+ + H2O (2.47)
≡FeOH + H3PO4 ≠= ≡FePO4- + 2H+ + H2O (2.48)
The equilibrium constants for acetate-iron and propionate-iron solution complexes indicate that the aqueous carboxylate ions specifically interact with iron and form strong complexes in solution. Complex formation constants are not available for fatty acids whose carbon chains contain more than three carbon atoms. Since the -CH2 group has hydrophobic characteristics, and the interaction of each -CH2 group involves an energy of about RT (~2500 Joules/mole or ~600 cal/mole), it is expected that an increase in the number of carbons has a direct influence on complex formation.
The surface complex formation constants for these fatty acids are estimated from those of the simple acids (namely, acetic and propionic acids), accounting for the additional contribution of the hydrophobic tails.