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The ion exchange properties of crystalline inorganic oxide-hydroxides
The Ion Exchange Properties of Crystalline Inorganic Oxide-Hydroxides Part I. /~FeOOH: A Variable Capacity Anion Exchanger RUSSELL PATERSON AND HABIBUR RAHMAN Department of Chemistry, The University of Glasgow, Glasgow G12 8QQ, Scotland, United Kingdom Received July 6, 1982; accepted November 15, 1982 Microcrystals of/~FeOOH were prepared some 2000 × 400 × 400 A3 in size. These were shown to have a variable pH-dependent anion exchange capacity, using new titration methods. No cationic branch was found in alkaline solutions. Capacities (for chloride) were high ( ~ 1 mmole/g or ~ 3 mmole/ml) at pH 3 and decreased to zero at pH 10. These capacities were also strongly dependent on counterion (chloride) concentrations in external solution. A thermodynamic analysis, based upon a pore model for the anion exchange mechanism, was verified experimentally. This showed the anion capacity to be a single valued of the activity of acid (HCI) in the equilibrium solution, conveniently expressed in terms of pA = pH + pCl.
Insoluble hydrous oxides of tri- and tetravalent metals have been of interest as ion exchangers since the mid 1950s (1), for possible use as ion exchangers in radiochemical separations of fission products. Typical members of the class were the hydrous oxides of Mn n, Fe m, A1m, Zr Iv, Sn Iv, Ti w, Th w, and Siw. These hydrous oxides have no well-defined stiochiometric formulas, since they contain not only oxide bridges, but also hydroxyl groups, bound water, and free water. the relative proportions of each are variable, depending on many factors, which include aging and history of drying or heat treatment, when the elements of water are lost irreversibly. Amphlett (2) has reviewed the literature on their exchange properties up to 1964 and, since that time, other reviews have appeared (3, 4). Most hydrous oxides have pseudo-amphoteric properties, with pH-variable capacity, anion exchange in acid and cation exchange in alkali. Experimental curves relating capacity to pH have been widely reported, but show wide variations, even between samples prepared batch-wise in the same laboratory. Possibly for these reasons, the mech-
anism of ion exchange has not been fully investigated. Selectivities for simple univalent anions are unremarkable (1, 2, 5). In contrast, polyvalent anions are strongly, often irreversibly bound. Phosphate uptake by zirconia is a good example. Indeed, this affinity has been exploited and has led to a parallel and successful development of acid-salt cation exchangers (3, 6), of which zirconium phosphate is the most celebrated. These disadvantages have limited the development and use of hydrous oxides for theoretical, commercial, and analytical purposes. In the present series of papers the thermodynamic and mechanistic aspects of exchange have been investigated using crystalline analogs of the original hydrous oxides. These are the chemically well-defined oxidehydroxides, which, in contrast to the gels, may be prepared reproducibly. Being crystalline, they might be expected to exhibit well-defined selectivity patterns, unattainable in gels, and even ion-sieve properties (7), all related to the inherent regularity of exchange sites in a crystalline matrix. The slower ki60
0021-9797/83 $3.00 Copyright © 1983 by Academic Press, Inc. All rights of reproduction in any form reserved.
Journal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
I O N E X C H A N G E OF O X I D E - H Y D R O X I D E S
netics of exchange, which might also be expected, due to diffusion through regular crystalline pores, would, in part, be off-set by their microcrystalline forms. Single crystals of 2000 A are typical of those prepared for this series. This first paper is concerned with the ion exchange properties of /3Iron(III)oxide-hydroxide (/3FeOOH). It was first prepared by Brhm (8), but not studied in detail until 1960, when a series of papers by Mackay (911) described the preparation, crystal structure, and morphology. It occurs in nature as the mineral akagrnite and can be prepared by various synthetic methods. The most important of these involve hydrolysis of dilute aqueous ferric chloride; chloride, it appears, is essential for the formation of the beta form of the oxide-hydroxide. When prepared by such methods, the/3FeOOH is found to consist of cigarshaped crystals between 1000 and 5000 /k long and 400-600 /~ wide, with square or circular cross-section. (For any particular preparative method, the size distribution is very narrow: almost uniform crystals are formed.) X-Ray powder diffraction studies (9) show the crystals to have the hollandite (BaMn8016) structure, with a tetragonal unit cell (a = 10.48 A and c = 3.02 A) and space group I4/m. The small crystal pore (5 × 5/~2) which, in hollandite, contains barium ions, in flFeOOH contains chloride and water in variable, nonstoichiometric amounts (9). Single crystal electron diffraction patterns showed the long axis of the cigar-like crystal to be the crystallographic c-axis, along which lie the hollandite pores. Chloride is present in large amounts in freshly prepared flFeOOH. It may be washed out by water (as HC1) until there remains some 2% of chloride (9), which, if removed by prolonged treatment with base, causes structural rearrangement, giving the s-form of FeOOH (geothite). The ion exchange properties of c~FeOOH is the subject of a separate paper (12). In this work /3FeOOH is shown to be a variable capacity anion exchanger with a
61
maximum useable capacity for chloride of approximately 1 mmole g - 1 (or some 3 mmole cm -3 of air-dried crystals. Unlike the hydrous (Fe m) gel or crystalline a F e 2 0 3 o r aFeOOH (12), which all have both cationic and anionic exchange properties, it was found that/3FeOOH is purely an anion exchanger with no cation exchange branch, even in alkali. Using a new titration procedure, it was found that the capacity for chloride (counteflon) was not only pH-dependent, but was also equally dependent on the concentration of chloride (or more precisely, its activity) in the external solution phase. This is a major and significant effect, not attributable to ionic strength effects upon activity coefficients. It is shown that this fundamental observation is a consequence of the mechanism of exchange, which involves coupled participation of both proton and chloride. EXPERIMENTAL
Preparation of/3FeOOH. The method of preparation of t3FeOOH microcrystals was based on the method of Matijevic and Scheiner (13). Three litres of a stock solution, 0.09 M in ferric chloride and 0.01 M in hydrochloric acid, was prepared. This solution was passed through a 0.45-~m Millipore filter to remove any colloidal contaminants which might be present before hydrolysis was begun. (The filter was prewashed to remove possible surfactants introduced as wetting agents in their manufacture.) Hydrolyses were performed at 90 and 100°C and maintained to +0.01°C. The solutions were continuously stirred using a Teflon-coated stirring bar. After 24 hr, they were cooled to room temperature and centrifuged. The precipitated microcrystals were redispersed in water, slightly acidified with hydrochloric acid, to prevent possible precipitation of any partly hydrolyzed iron, and centrifuged once more. This procedure was repeated several times to wash the crystals thoroughly. The final dispersions were Journal of Colloid and Interface Science. Vol. 94, No. I, July 1983
62
PATERSON
AND
made using distilled water. The washings showed no detectable, soluble iron species on addition of base or using orthophenanthrolein indicator. The final washings remained acidic due to the slow release of hydrochloric acid absorbed during the hydrolysis by the mechanism which is the subject of this paper. The crystalline precipates, thus obtained, were air dried overnight at 42°C and stored in sealed glass bottled until required. These preparations were examined by electron microscopy for size distribution and morphology and for single crystal electron diffraction. [In all cases these agreed exactly with those expected for/3FeOOH (9).] Samples of the air-dried /3FeOOH were dispersed in water and in solutions of hydrochloric acid of increasing strength to determine the practical range of usable acidities before dissolution was detectable. pH titration. Samples of between 0.3 and 0.5 g were dispersed in 20-30 ml of water. (Suspensions were usually left overnight, with gentle stirring, before use.) Burettes with very fine capillary tips were employed in the titration procedures and these tips were immersed in the suspension during the titration and measurement. Control experiments showed that diffusion oftitrants into the suspension was not detectable when the capillaries were carefully prepared. Full precautions were taken to prevent carbon dioxide uptake by the sodium hydroxide which was dispensed from an automatic burette protected by "Sofnolite" guard tube and the suspension itself was flushed with presaturated nitrogen. The titration vessel was thermostarted to 25°C and the glass electrode system calibrated using a series of five buffers in the range 2-12. pH measurements were made using an Orion Ion Analyzer, type 801, and were reproducible to 0.002 pH units. Chloride release. When chloride release was implied from the interpretation of pH titration curves, it was verified by direct chloride analysis, using a Philips chloride selective electrode. These were used only for confirmation because the accuracy and reproJournal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
RAHMAN
ducibility of these measurements were inferior to the glass electrode system. All titrations were begun with dispersions of/~FeOOH in water. A number of batches were prepared in the course of these studies. They differed slightly from one another in terms of water and chloride content. For comparison therefore, all batches were ignited to constant weight at 1000°C, forming Fe203. The capacities were then expressed per gram of formula weight of FeOOH in the air-dried sample. Isotopic equilibrations. Samples of air-dried /3FeOOH were equilibrated with dilute solutions of sodium chloride (~0.08 M), labeled with chloride-36. Samples sizes and solution volumes were adjusted, such that the amount of chloride in solution was less than or equal to that expected in the exchanger. After overnight equilibration, the suspension was filtered and aqueous samples counted, using standard B-scintillation methods. TITRATION
THEORY
As prepared, /3FeOOH contained leachable hydrochloric acid. When a sample of the crystals was equilibrated with a volume of water, this acid was released in part, reducing the solution pH to approximately 3. In titration this equilibrium was successively displaced by addition of sodium hydroxide, hydrochloric acid, or sodium chloride. Preliminary tests showed that these equilibria were obtained in a few minutes after addition of reagent and so were suitable for a titration technique. Consider initially only the addition of base, In this process acid in solution is neutralized, causing release of more acid from the exchanger, by an amount AX (mmole g-~) from its initial, but unknown capacity of Xo (mmole g-l). Addition of acid or salt was found to increase acid uptake, in which case, &X would be negative. AX is therefore the change in capacity of the sample relative to its initial equilibrium value when dispersed in water. The absolute value of the capacity, X, corresponding to the amount of
ION EXCHANGE OF OXIDE-HYDROXIDES chloride/proton uptake per gram o f F e O O H content at any point in the titration is therefore u n k n o w n unless X0 m a y be found (Eq. Ill). z x x = ( x 0 - Jr).
[1]
F o r / J F e O O H the value of A X increased to a constant value in alkaline solutions above p H 10 (Tables I and II). This was designated AXmax and indicated that all the acid originally present in the exchanger in equilibrium with the water had been released and neutralised. Consequently from Eq. [1], X0 = AXn,ax.
63
was negligible D o n n a n uptake of salt at the ionic strengths used in this work (<0.1). In addition, it was also assumed that solution volumes were additive so that the total volume o f solution, V, is given by Eq. [2]. V = Vo+ Vb+ Va+ Vs.
[2]
Since sodium ion is assumed excluded from the exchanger, its concentration is determined solely by the a m o u n t of base and salt added (Eq. [3]). (Molar concentrations are represented in square brackets and for clarity, charges are omitted.) [Na] = (S + B)/V.
Calculation o f Capacity Consider an initial equilibration o f a weight of air-dried material, corresponding to G g of FeOOH, with a volume, V0 ml of water at 25°C, giving an acidic solution. Subsequent additions of B, A, or S m m o l e of sodium hydroxide, hydrochloric acid, or sodium chloride were each made using V ml of the respective reagent. Subscripts b, a, and s, used below, distinguish base, acid, and salt, respectively. In the calculation of X, several assumptions were made. It was assumed that there
[3]
From the electroneutrality of the solution, [Na] + [H] = [CI] + [OH].
[4]
In consequence, chloride concentration and ionic strength,/, are defined by Eqs. [5] and [6]. [CI] = (S + B + [H]. V - [ O H ] . V ) / V
[51
and
I = (S + B ) / V + [H].
[6]
AX which is the acid release, if positive, or uptake, if negative, is therefore obtained
TABLE I Titration of a Suspension of flFeOOH(100) in Water with Sodium Hydroxide and without Added Salff'b Vb ~
V~ ~
0.00 0.20 0.40 0.60 0.80 1.00 1.20 1.40 1.60 1.80 2.00 2.20
0 0 0 0 0 0 0 0 0 0 0 0
pH
3.332 3.985 4.885 5.707 6.550 7.420 8.445 10.394 11.120 11.370 11.530 11.644
pA
AX ~
6.664 0.000 6.845 0.090 7.485 0.205 8,141 0.327 8.867 0.451 9.647 0.574 10.599 0.697 12.500 0.795 13.228 ~ 0.801 13.480 Av. 0.807 j 0.810 13.643 | 0.813 13.760 L 0.817 /
/
Xc
0.807 0.717 0.602 0.480 0.356 0.233 0.110 0.012 0.006 -0.003 -0.006 -0.010
a Initial volume, 25.00 ml; molarity of base, 0.1676 M; grams of exchanger, 0.3034; grams FeOOH, 0.2712. bShown as the lower curve in Fig. 2. c Vb' Volume of base, ml; Vs, volume of salt solution, ml; AX, mEq base taken up/g FeOOH;X; Chloride capacity mEq/g FeOOH. Journal of Colloid and Interface Science, Vol. 94, N o . 1, J u l y 1983
64
PATERSON AND RAHMAN TABLE II Titration of a Suspension of/3FeOOH(100) in Water with Sodium Hydroxide, with the Addition of 1.004 mmole of Sodium Chloride~'b Vbc
V~"
pH
pA
0.00 0.00 0.20 0.40 0.60 0.80 1.00 1.20 1.41 1.60 1.80 2.00 2.20
0.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00
3.337 4.118 5.146 5.978 6.575 7.272 8.057 8.924 10.641 11.069 11.320 11.469 11.577
6.674 5.606 6.625 7.447 8.035 8.723 9.499 10.358 12.071 12.500 12.754 12.905 13.016
AX c
0.000 -0.035 0.081 0.204 0.327 0.451 0.574 0.697 0.774 ~-0.801 j 0.809 Av. 0.809 /0.823 ~ 0.838
Xc
0.809 0.844 0.728 0.605 0.482 0.358 0.235 0.112 0.035 0.008 0.000 -0.014 -0.029
a Initial volume, 25.00 ml; molarity of base, 0.1676 M; molarity of salt, 1.004 M; grams of exchanger, 0.3035; grams FeOOH, 0.2713. b Shown as the upper curve in Fig. 2. c Vb, Volume of base, ml; Vs, volume of salt solution, ml; AX, mEq base taken up/g FeOOH; X; chloride capacity, mEq/g FeOOH.
from the apparent imbalance of acidity, viewed from the solution phase (Eq. [7]). ~ x = (B - A + [H]- V -
[OH]- V
- [H0].
Vo)/G.
l o g ~ = -0.5(1/I/(1 + ~/-/)- 0.3•). [7]
In Eq. [7] AX is expressed in mmole g-i of F e O O H in the air dried sample and [Ho] is the initial hydrogen ion concentration of the suspension in water. For these calculations hydrogen ion concentrations are required. They were obtained by an iterative calculation from pH. In the first approximation, activity coefficients were set to unity to estimate the hydrogen concentration and so to make a first approximation to the ionic strength. Using the Davies (14) expression for activity coefficients (Eq. [8]), an improved estimate of hydrogen and hydroxyl ion concentrations were made. These, in turn, modified the ionic strength somewhat. This iterative cycle was repeated until the concentrations thus calculated were constant to 0.1% or better. The calculation was obviously sensitive to the particular ionic strength expression used. The Davies equaJournal o f Colloid a n d Interface Science.
Vol.94, No. 1, July 1983
tion was chosen after tests on two systems where X is known. [81
Both tests involved solution titrations. The first was the application of this m e t h o d to a titration o f hydrochloric acid with sodium hydroxide. In this case AX should be zero at all pHs, since both reagents are fully dissociated and consequently there are no hidden "sources" or "sinks" for either acid or base. The results showed that AX was indeed zero at all p H values of the titration, which ranged from pH 1.7 to 11.3. In all cases AX calculated was 0 _ 0.005 mmole. In the second test acetic acid was titrated. In that case AX was positive and rose linearly to a constant, m a x i m u m value, AXmax, which corresponded to the calculated a m o u n t of undissociated acid in the original solution (within 0.5%). The m e t h o d is particularly suited to ~FeOOH titrations where the titration endpoints are not obvious (see Fig. 2) and may be further masked in titrations where salt has been added.
65
ION EXCHANGE OF OXIDE-HYDROXIDES
FIG. 1. Electron micrograph of flFeOOH crystals prepared by hydrolysis of ferric chloride at 100°C, by the method of Matijevic and Scheiner (13).
Journal of Colloid and Interface Science.
Vol.94, No. 1, July 1983
66
PATERSON AND R A H M A N RESULTS A N D DISCUSSION
The electron micrograph (Fig. 1) shows the characteristic shape of the 13FeOOH crystals, obtained by the hydrolysis of ferric chloride at 100°C by the method of Matijevic and Scheiner (13). Those obtained by hydrolysis at 90°C were identical to those shown here. Electron diffraction on single crystals gave only #FeOOH spacings (9). The crystals as shown, were very regular in size and shape. Their dimensions were 2000/~ (a +_ 400 A) in the direction of the c-axis, and 400 A (a + 70 h) in width. A number of titrations were made on #FeOOH with alkali and typical pH titration curves are shown in Fig. 2. In that figure there appear two distinct titration plots, which merge in alkaline solution above pH 10. The lower curve was obtained when #FeOOH was dispersed in water and titrated with sodium hydroxide, with no added sodium chloride present. The pH curve rises steeply from pH 3 to 9 and gives finally an indistinct end point, unsuitable for quantitative work. Titration data are given in Table I. If the standard sample of 0.303 g of air dried/3FeOOH was again dispersed in 25 ml of water and 12 ./, ?
lo
[3FeOOH
8 •
4
/! / , ' ~" / "# / "
pH TITRATION WITH
ADDEDSALT, NaCl
j'+/' /
~,,
2 0
i
I
1
2
mt OF BASE FIG. 2. pH titration curves for flFeOOH(100). Samples of the air dried material (0.3 g) were titrated with 0.1676 M NaOH. Initial point for all titrations pH 3.3. Lower curve with no added salt as Table I. Upper curve, after addition of I mmole NaCI (Table II). Dotted curve, after addition of 2 mmole NaC1. Arrows, transitions from lower to upper curves by addition of 1 mmole of NaC1 during titration. Journal of Colloid and InterfaceScience, Vol. 94, No. 1, July 1983
now 1 mmole of sodium chloride added (1 ml × 1 Msalt), the pH of the suspension rose by almost one pH unit (Table II and Fig. 2). Thereafter, on addition of base, the pH followed the upper curve to the end point. The transition from lower to upper curves could be made at any point by addition of salt. Three such titrations were made and the transitional jumps recorded as vertical arrows on Fig. 2. An addition of 2 mmole of sodium chloride gave only a marginal additional increase in pH (Fig. 2). Dispersions of #FeOOH in water usually gave a pH of approximately 3 but were variable from batch to batch, dependent upon the degree of washing during preparation. Solubility tests showed however that in the titration range of pH, greater than 3, soluble iron was not detectable and was therefore less than 10-6 M. In alkaline solution, at pH greater than 10, salt addition caused no change in the pH (other than the very small effect due to ionic strength effects). Salt addition in acidic solutions, however, caused proton uptake by t3FeOOH. This is a very sensitive test of uptake, since a large pH change, say between 5 and 6, as observed in Fig. 2, would amount to a transfer of only some 3 × 10-4 mmole of acid from the solution into the solid. A more exact account of the acid/base balance between solution and #FeOOH is obtained by calculation of AX, obtained from Eq. [7] and defined in the titration theory as the release of protons (mmole/g FeOOH), during the titration, taken relative to the initial capacity of the #FeOOH sample dispersed in water. AX values for the six titrations, shown in Fig. 2, were calculated. Data for the two titrations defining the entire upper and the lower curves are given in Tables I and II. The proton release (AX) on addition of base, although variable with salt addition, reaches a constant, maximum value for all samples. This AXmax expressed in mmole/g FeOOH, represents the capacity of the samples of #FeOOH(100) in equilibrium with water at pH 3. The true capacity of X, at
ION E X C H A N G E O F O X I D E - H Y D R O X I D E S
each titration point is A X m a x - - AX (Tables I and II). In the titration of t3FeOOH, the release of protons was shown to be accompanied by an equal amount of chloride, using chloride-reversible ion selective electrodes. This was confirmed by direct analysis of total chloride released in alkali and showed that under all conditions proton and chloride are released stoichiometrically as hydrochloric acid. Isotopic exchange of 36C1- under conditions of constant pH and salt concentration gave identical capacities. This also showed that chloride was freely exchangeable by an anion exchange process. Plots of capacity, X, against pH for the six titrations of Fig. 2 are shown in Fig. 3. Once more all the points lie on two curves and transitions between them made an addition of 1 mmole of salt, indicated by arrows. It is obvious that, at a given pH, the anion capacity is much larger if sodium chloride is added. The arrows, indicating transitions, also have slightly positive slopes because the process of salt addition raises pH and removes protons and chloride from solution to raise the capacity of the exchanger. Addition of twice as much salt has little further effect and is scarcely detectable. The capacity of/3FeOOH for acid (its anX -080 >-
0060
CA o.~o tLl O
CE 0 -d
I o
0.20
ion capacity also) is therefore clearly not a single-valued function of pH as implied in the literature on hydrous oxides (1, 2). It is obviously a function of both pH and salt concentration. In this series of papers on the ion exchange properties of microcrystals of hydrous oxides, similar salt/pH dependency has been observed in acidic (and alkaline) solutions with both aFeOOH (12) and monoclinic zirconia (15). Such effects have also been reported for other amphoteric colloidal oxides and often used to estimate the pH of the zero point of charge (z.p.c.). The literature does not however provide a simple explanation of the effects. In this paper a pore model mechanism is proposed. The model has thermodynamic consequences which are supported by the experimental data.
The Pore Model for Anionic Capacity of l3FeOOH The explanation for the major influence of chloride on capacity lies in the basic mechanism for exchange. This involves a thermodynamic coupling between the "protonation" reaction, by which positive charges are created in the matrix of the crystal and uptake of chloride (or other anion) to compensate these charges. The term protonation is not precisely defined, since the actual interaction between proton and ~FeOOH is not known. It is probably a reaction between a bound hydroxyl group, yielding a coordinated water (Eq. [9]). For thermodynamics it is only necessary to know that such an interaction exists. FeOOH + H + ~- FeO(OH2) +
o
I
I 5
I
• "~.~..,.~..2.~.. I I I 7
9
67
[9]
e." I
I 11
pH FIG. 3. Chloride capacity, X ( m m o l e / g F e O O H ) against p H using the titration data of Fig. 2. O, Initial point for all titrations, which is also the arbitrary comm o n reference for all 2xX calculations. Arrows represent additions of 1 m m o l e of salt. T h e u p p e r m o s t curve obtained after 2 m m o l e addition.
Bars over species in Eq. [9] indicate that the reactions are assumed to occur within the pores of 13FeOOH in this model. The pores may be the small hollandite pores (9) or the macropores of Gallagher (16). Equation [9] is a base protonation reaction, at least formally, for which an acid dissociation constant,/Ca, may be defined (Eq. [10]). Journal of Colloid and Interface Science, Vol. 94, No. I, July 1983
68
PATERSON AND RAHMAN
g a = (aFeoO H • dH+)/dFeOOHg.
[
10]
The protons of Eq. [9] are pore protons (H ÷) which are in a Donnan equilibrium with those of the external solution (H +) (Eq. [11 ]). H+ ~ H+
[11]
Because this is an interfacial equilibrium involving charge separation, and driven by the affinity of the protonation, the activities of the proton in each phase are not equal, but depend upon the Donnan potential (ff - ~)
(Eq. [1 la]) aH+/dH+ = exp(F(~/- ~ ) / R T ) .
[ 11 al
If only the equilibria of Eqs. [9] and [11] applied, then the protonation would not proceed because uncompensated transfer of positive charge from solution to pore would create an enormous positive Donnan potential (~ ~k), so that in effect aH would become zero (Eq. [1 la]). There would therefore be no protonation of the pore sites, regardless of the favorability of Ka (Eq. [10]) to protonation. The process can occur only if chloride (or other counterion) may also enter the pores. This allows effective electroneutrality to be achieved in the pores, while large numbers of protons (now accompanied by chloride) are transferred. The equilibrium between chloride in solution (CF) and in the -
Cl- ~ Cl-
[121
pore (CI-) (Eq. [I 2]) may be combined with the proton equilibrium (Eq. [11 ]) to give the well-known Donnan equilibrium relationship for electrolyte uptake of acid (Eq. [13]). aH+'acx- = dn+'dcl •
[13]
activities of the pore species, FeOOH~, CI-, and FeOOH, will be determined solely by the activity of the acid in solution, aH'acl. It is assumed that there is no salt uptake in the pore. Under such conditions, a particular value of the activity of acid, a H ' a o or equivalently in logarithmic form, pA = pH + pCl, however constituted by combinations of pH and chloride activities, would define only one pore condition and hence only one capacity. The experimental capacities equal the concentration of pore species FeOOH~ and chloride and would therefore be a single-valued function of all" a a or more conveniently, of pA. This conclusion is independent of the detailed thermodynamic model which might be used, since it is independent of the standard states which might be used or the method by which activity coefficients might be evaluated. It is noteworthy that protonation of surface sites would not require such coupling to the anion and there protonation should be solely a function of the pH of the external solution. Plots of capacity, X, against pA for the titration data of Fig. 2 and Tables I and II, are given in Fig. 4. The results show conclusively that chloride capacity for BFeOOH is a single-valued function of the activity of the total
X
>_I'--
Ka = -
-
dwooH " (aN+" acv) a-veOOH~"dcl-
[14]
Since, by definition, Ka is a constant, the Journal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
o.6o
£3
0
.N
'~
"3,
0.40
(For simplicity the same standard states are assumed for pore and solution.) For Eqs. [10] and [13]
0.80
-lL)
0.20
E E
"k ":k
-,<. o I
I
7
I
I
4
g
i
I~
I
I
13
15
pA FIG. 4. Chloride capacity, X (mmole/g FeOOH), against pA (-log aH• acl), using the data of Fig. 2. Data presented refer to/3FeOOH(100) (13) samples prepared at 90°C/3FeOOH(90) by the same method and samples prepared by hydrolysis of ferric chloride at 25°C (17) gave identical results.
ION EXCHANGE OF OXIDE-HYDROXIDES acid, pA, as p r e d i c t e d b y t h e p o r e m o d e l above. C a p a c i t y a g a i n s t p A c u r v e s were det e r m i n e d u s i n g f l F e O O H (90) a n d w i t h /3FeOOH prepared by the method of Fryer et al. (17) f r o m 0.01 M ferric c h l o r i d e a l l o w e d to h y d r o l y z e at r o o m t e m p e r a t u r e . I n o n e case s a m p l e s h a d b e e n left for 7 years a n d in a n o t h e r t h e h y d r o l y s i s p e r i o d was a few m o n t h s . I n all these cases t h e c a p a c i t i e s for c h l o r i d e ( e x p r e s s e d as m m o l e / g F e O O H ) were i d e n t i c a l f u n c t i o n s o f p A a n d superp o s a b l e o n Fig. 4. T h e r e p r o d u c i b i l i t y o f these s a m p l e p r e p a r a t i o n s is t h e r e f o r e excellent a n d m a n y b a t c h e s w i t h i d e n t i c a l p r o p erties h a v e b e e n m a d e in the c o u r s e o f this work. S p e c u l a t i o n t h a t t h e l o c a t i o n o f exc h a n g e sites w i t h i n t h e h o l l a n d i t e p o r e s is n o w s u p p o r t e d b y e v i d e n c e o f a n i o n exclusion o f p e r c h l o r a t e , discussed in p a r t II (7). REFERENCES 1. Kraus, K. A., Phillips, H. O., Carlson, T. A., and Johnson, J. S., Proc. Second Int. Conf. "Peaceful Uses of Atomic Energy," United Nations, Geneva, 28, 3 (1958).
69
2. Amphlett, C. B., "Inorganic Ion Exchangers." Elsevier, Amsterdam, 1964. 3. Clearfield, A., Nancollas, G. H., and Blessing, R. H., in "Ion Exchange and Solvent Extraction" (J. A. Marinsky and Y. Marcus, Eds.), Vol. 5, Ch. 1. Dekker, New York, 1973. 4. Vesely, V., and Pekarek, V., Talanta 19, 1219 (1972). 5. Nancollas, G. H., and Paterson, R., J. Inorg. Nucl. Chem. 29, 565 (1967). 6. Clearfield, A., and Stynes, J. A., J. Inorg. Nucl. Chem. 26, 117 (1964). 7. Paterson, R., and Rahman, H., Part II, J. Colloid Interface Sci., submitted for pubfication. 8. B/3hm, J., Z. Anorg. Allgem-Chem. 149, 203 (1925). 9. Mackay, A. L., Mineral Magn. 32, 545 (1960). 10. Mackay, A. L., Mineral Magn. 33, 270 (1962). 11. Mackay, A. L., J. Phys. Soc. Japan 17, 317 (1962). 12. Paterson, R., and Rahman, H., Part II1, J. Colloid Interface Sci., in press. 13. Matijevic, E., and Seheiner, P., J. Colloid Interface. Sci. 63, 509 (1978). 14. Davies, E. W., "Ion Association." Butterworths, London, 1962. 15. Paterson, R., and Rahman, H., Part IV, manuscript in preparation. 16. Gallagher, K. J., Nature (London) 226, 1225 (1970). 17. Fryer, J. R., Gildawie, A. M., and Paterson, R., Nature (London) 226, 149 (1974).
Journalof Colloidand InterfaceScience, Vol.94, No. 1, July 1983
Insoluble hydrous oxides of tri- and tetravalent metals have been of interest as ion exchangers since the mid 1950s (1), for possible use as ion exchangers in radiochemical separations of fission products. Typical members of the class were the hydrous oxides of Mn n, Fe m, A1m, Zr Iv, Sn Iv, Ti w, Th w, and Siw. These hydrous oxides have no well-defined stiochiometric formulas, since they contain not only oxide bridges, but also hydroxyl groups, bound water, and free water. the relative proportions of each are variable, depending on many factors, which include aging and history of drying or heat treatment, when the elements of water are lost irreversibly. Amphlett (2) has reviewed the literature on their exchange properties up to 1964 and, since that time, other reviews have appeared (3, 4). Most hydrous oxides have pseudo-amphoteric properties, with pH-variable capacity, anion exchange in acid and cation exchange in alkali. Experimental curves relating capacity to pH have been widely reported, but show wide variations, even between samples prepared batch-wise in the same laboratory. Possibly for these reasons, the mech-
anism of ion exchange has not been fully investigated. Selectivities for simple univalent anions are unremarkable (1, 2, 5). In contrast, polyvalent anions are strongly, often irreversibly bound. Phosphate uptake by zirconia is a good example. Indeed, this affinity has been exploited and has led to a parallel and successful development of acid-salt cation exchangers (3, 6), of which zirconium phosphate is the most celebrated. These disadvantages have limited the development and use of hydrous oxides for theoretical, commercial, and analytical purposes. In the present series of papers the thermodynamic and mechanistic aspects of exchange have been investigated using crystalline analogs of the original hydrous oxides. These are the chemically well-defined oxidehydroxides, which, in contrast to the gels, may be prepared reproducibly. Being crystalline, they might be expected to exhibit well-defined selectivity patterns, unattainable in gels, and even ion-sieve properties (7), all related to the inherent regularity of exchange sites in a crystalline matrix. The slower ki60
0021-9797/83 $3.00 Copyright © 1983 by Academic Press, Inc. All rights of reproduction in any form reserved.
Journal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
I O N E X C H A N G E OF O X I D E - H Y D R O X I D E S
netics of exchange, which might also be expected, due to diffusion through regular crystalline pores, would, in part, be off-set by their microcrystalline forms. Single crystals of 2000 A are typical of those prepared for this series. This first paper is concerned with the ion exchange properties of /3Iron(III)oxide-hydroxide (/3FeOOH). It was first prepared by Brhm (8), but not studied in detail until 1960, when a series of papers by Mackay (911) described the preparation, crystal structure, and morphology. It occurs in nature as the mineral akagrnite and can be prepared by various synthetic methods. The most important of these involve hydrolysis of dilute aqueous ferric chloride; chloride, it appears, is essential for the formation of the beta form of the oxide-hydroxide. When prepared by such methods, the/3FeOOH is found to consist of cigarshaped crystals between 1000 and 5000 /k long and 400-600 /~ wide, with square or circular cross-section. (For any particular preparative method, the size distribution is very narrow: almost uniform crystals are formed.) X-Ray powder diffraction studies (9) show the crystals to have the hollandite (BaMn8016) structure, with a tetragonal unit cell (a = 10.48 A and c = 3.02 A) and space group I4/m. The small crystal pore (5 × 5/~2) which, in hollandite, contains barium ions, in flFeOOH contains chloride and water in variable, nonstoichiometric amounts (9). Single crystal electron diffraction patterns showed the long axis of the cigar-like crystal to be the crystallographic c-axis, along which lie the hollandite pores. Chloride is present in large amounts in freshly prepared flFeOOH. It may be washed out by water (as HC1) until there remains some 2% of chloride (9), which, if removed by prolonged treatment with base, causes structural rearrangement, giving the s-form of FeOOH (geothite). The ion exchange properties of c~FeOOH is the subject of a separate paper (12). In this work /3FeOOH is shown to be a variable capacity anion exchanger with a
61
maximum useable capacity for chloride of approximately 1 mmole g - 1 (or some 3 mmole cm -3 of air-dried crystals. Unlike the hydrous (Fe m) gel or crystalline a F e 2 0 3 o r aFeOOH (12), which all have both cationic and anionic exchange properties, it was found that/3FeOOH is purely an anion exchanger with no cation exchange branch, even in alkali. Using a new titration procedure, it was found that the capacity for chloride (counteflon) was not only pH-dependent, but was also equally dependent on the concentration of chloride (or more precisely, its activity) in the external solution phase. This is a major and significant effect, not attributable to ionic strength effects upon activity coefficients. It is shown that this fundamental observation is a consequence of the mechanism of exchange, which involves coupled participation of both proton and chloride. EXPERIMENTAL
Preparation of/3FeOOH. The method of preparation of t3FeOOH microcrystals was based on the method of Matijevic and Scheiner (13). Three litres of a stock solution, 0.09 M in ferric chloride and 0.01 M in hydrochloric acid, was prepared. This solution was passed through a 0.45-~m Millipore filter to remove any colloidal contaminants which might be present before hydrolysis was begun. (The filter was prewashed to remove possible surfactants introduced as wetting agents in their manufacture.) Hydrolyses were performed at 90 and 100°C and maintained to +0.01°C. The solutions were continuously stirred using a Teflon-coated stirring bar. After 24 hr, they were cooled to room temperature and centrifuged. The precipitated microcrystals were redispersed in water, slightly acidified with hydrochloric acid, to prevent possible precipitation of any partly hydrolyzed iron, and centrifuged once more. This procedure was repeated several times to wash the crystals thoroughly. The final dispersions were Journal of Colloid and Interface Science. Vol. 94, No. I, July 1983
62
PATERSON
AND
made using distilled water. The washings showed no detectable, soluble iron species on addition of base or using orthophenanthrolein indicator. The final washings remained acidic due to the slow release of hydrochloric acid absorbed during the hydrolysis by the mechanism which is the subject of this paper. The crystalline precipates, thus obtained, were air dried overnight at 42°C and stored in sealed glass bottled until required. These preparations were examined by electron microscopy for size distribution and morphology and for single crystal electron diffraction. [In all cases these agreed exactly with those expected for/3FeOOH (9).] Samples of the air-dried /3FeOOH were dispersed in water and in solutions of hydrochloric acid of increasing strength to determine the practical range of usable acidities before dissolution was detectable. pH titration. Samples of between 0.3 and 0.5 g were dispersed in 20-30 ml of water. (Suspensions were usually left overnight, with gentle stirring, before use.) Burettes with very fine capillary tips were employed in the titration procedures and these tips were immersed in the suspension during the titration and measurement. Control experiments showed that diffusion oftitrants into the suspension was not detectable when the capillaries were carefully prepared. Full precautions were taken to prevent carbon dioxide uptake by the sodium hydroxide which was dispensed from an automatic burette protected by "Sofnolite" guard tube and the suspension itself was flushed with presaturated nitrogen. The titration vessel was thermostarted to 25°C and the glass electrode system calibrated using a series of five buffers in the range 2-12. pH measurements were made using an Orion Ion Analyzer, type 801, and were reproducible to 0.002 pH units. Chloride release. When chloride release was implied from the interpretation of pH titration curves, it was verified by direct chloride analysis, using a Philips chloride selective electrode. These were used only for confirmation because the accuracy and reproJournal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
RAHMAN
ducibility of these measurements were inferior to the glass electrode system. All titrations were begun with dispersions of/~FeOOH in water. A number of batches were prepared in the course of these studies. They differed slightly from one another in terms of water and chloride content. For comparison therefore, all batches were ignited to constant weight at 1000°C, forming Fe203. The capacities were then expressed per gram of formula weight of FeOOH in the air-dried sample. Isotopic equilibrations. Samples of air-dried /3FeOOH were equilibrated with dilute solutions of sodium chloride (~0.08 M), labeled with chloride-36. Samples sizes and solution volumes were adjusted, such that the amount of chloride in solution was less than or equal to that expected in the exchanger. After overnight equilibration, the suspension was filtered and aqueous samples counted, using standard B-scintillation methods. TITRATION
THEORY
As prepared, /3FeOOH contained leachable hydrochloric acid. When a sample of the crystals was equilibrated with a volume of water, this acid was released in part, reducing the solution pH to approximately 3. In titration this equilibrium was successively displaced by addition of sodium hydroxide, hydrochloric acid, or sodium chloride. Preliminary tests showed that these equilibria were obtained in a few minutes after addition of reagent and so were suitable for a titration technique. Consider initially only the addition of base, In this process acid in solution is neutralized, causing release of more acid from the exchanger, by an amount AX (mmole g-~) from its initial, but unknown capacity of Xo (mmole g-l). Addition of acid or salt was found to increase acid uptake, in which case, &X would be negative. AX is therefore the change in capacity of the sample relative to its initial equilibrium value when dispersed in water. The absolute value of the capacity, X, corresponding to the amount of
ION EXCHANGE OF OXIDE-HYDROXIDES chloride/proton uptake per gram o f F e O O H content at any point in the titration is therefore u n k n o w n unless X0 m a y be found (Eq. Ill). z x x = ( x 0 - Jr).
[1]
F o r / J F e O O H the value of A X increased to a constant value in alkaline solutions above p H 10 (Tables I and II). This was designated AXmax and indicated that all the acid originally present in the exchanger in equilibrium with the water had been released and neutralised. Consequently from Eq. [1], X0 = AXn,ax.
63
was negligible D o n n a n uptake of salt at the ionic strengths used in this work (<0.1). In addition, it was also assumed that solution volumes were additive so that the total volume o f solution, V, is given by Eq. [2]. V = Vo+ Vb+ Va+ Vs.
[2]
Since sodium ion is assumed excluded from the exchanger, its concentration is determined solely by the a m o u n t of base and salt added (Eq. [3]). (Molar concentrations are represented in square brackets and for clarity, charges are omitted.) [Na] = (S + B)/V.
Calculation o f Capacity Consider an initial equilibration o f a weight of air-dried material, corresponding to G g of FeOOH, with a volume, V0 ml of water at 25°C, giving an acidic solution. Subsequent additions of B, A, or S m m o l e of sodium hydroxide, hydrochloric acid, or sodium chloride were each made using V ml of the respective reagent. Subscripts b, a, and s, used below, distinguish base, acid, and salt, respectively. In the calculation of X, several assumptions were made. It was assumed that there
[3]
From the electroneutrality of the solution, [Na] + [H] = [CI] + [OH].
[4]
In consequence, chloride concentration and ionic strength,/, are defined by Eqs. [5] and [6]. [CI] = (S + B + [H]. V - [ O H ] . V ) / V
[51
and
I = (S + B ) / V + [H].
[6]
AX which is the acid release, if positive, or uptake, if negative, is therefore obtained
TABLE I Titration of a Suspension of flFeOOH(100) in Water with Sodium Hydroxide and without Added Salff'b Vb ~
V~ ~
0.00 0.20 0.40 0.60 0.80 1.00 1.20 1.40 1.60 1.80 2.00 2.20
0 0 0 0 0 0 0 0 0 0 0 0
pH
3.332 3.985 4.885 5.707 6.550 7.420 8.445 10.394 11.120 11.370 11.530 11.644
pA
AX ~
6.664 0.000 6.845 0.090 7.485 0.205 8,141 0.327 8.867 0.451 9.647 0.574 10.599 0.697 12.500 0.795 13.228 ~ 0.801 13.480 Av. 0.807 j 0.810 13.643 | 0.813 13.760 L 0.817 /
/
Xc
0.807 0.717 0.602 0.480 0.356 0.233 0.110 0.012 0.006 -0.003 -0.006 -0.010
a Initial volume, 25.00 ml; molarity of base, 0.1676 M; grams of exchanger, 0.3034; grams FeOOH, 0.2712. bShown as the lower curve in Fig. 2. c Vb' Volume of base, ml; Vs, volume of salt solution, ml; AX, mEq base taken up/g FeOOH;X; Chloride capacity mEq/g FeOOH. Journal of Colloid and Interface Science, Vol. 94, N o . 1, J u l y 1983
64
PATERSON AND RAHMAN TABLE II Titration of a Suspension of/3FeOOH(100) in Water with Sodium Hydroxide, with the Addition of 1.004 mmole of Sodium Chloride~'b Vbc
V~"
pH
pA
0.00 0.00 0.20 0.40 0.60 0.80 1.00 1.20 1.41 1.60 1.80 2.00 2.20
0.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00
3.337 4.118 5.146 5.978 6.575 7.272 8.057 8.924 10.641 11.069 11.320 11.469 11.577
6.674 5.606 6.625 7.447 8.035 8.723 9.499 10.358 12.071 12.500 12.754 12.905 13.016
AX c
0.000 -0.035 0.081 0.204 0.327 0.451 0.574 0.697 0.774 ~-0.801 j 0.809 Av. 0.809 /0.823 ~ 0.838
Xc
0.809 0.844 0.728 0.605 0.482 0.358 0.235 0.112 0.035 0.008 0.000 -0.014 -0.029
a Initial volume, 25.00 ml; molarity of base, 0.1676 M; molarity of salt, 1.004 M; grams of exchanger, 0.3035; grams FeOOH, 0.2713. b Shown as the upper curve in Fig. 2. c Vb, Volume of base, ml; Vs, volume of salt solution, ml; AX, mEq base taken up/g FeOOH; X; chloride capacity, mEq/g FeOOH.
from the apparent imbalance of acidity, viewed from the solution phase (Eq. [7]). ~ x = (B - A + [H]- V -
[OH]- V
- [H0].
Vo)/G.
l o g ~ = -0.5(1/I/(1 + ~/-/)- 0.3•). [7]
In Eq. [7] AX is expressed in mmole g-i of F e O O H in the air dried sample and [Ho] is the initial hydrogen ion concentration of the suspension in water. For these calculations hydrogen ion concentrations are required. They were obtained by an iterative calculation from pH. In the first approximation, activity coefficients were set to unity to estimate the hydrogen concentration and so to make a first approximation to the ionic strength. Using the Davies (14) expression for activity coefficients (Eq. [8]), an improved estimate of hydrogen and hydroxyl ion concentrations were made. These, in turn, modified the ionic strength somewhat. This iterative cycle was repeated until the concentrations thus calculated were constant to 0.1% or better. The calculation was obviously sensitive to the particular ionic strength expression used. The Davies equaJournal o f Colloid a n d Interface Science.
Vol.94, No. 1, July 1983
tion was chosen after tests on two systems where X is known. [81
Both tests involved solution titrations. The first was the application of this m e t h o d to a titration o f hydrochloric acid with sodium hydroxide. In this case AX should be zero at all pHs, since both reagents are fully dissociated and consequently there are no hidden "sources" or "sinks" for either acid or base. The results showed that AX was indeed zero at all p H values of the titration, which ranged from pH 1.7 to 11.3. In all cases AX calculated was 0 _ 0.005 mmole. In the second test acetic acid was titrated. In that case AX was positive and rose linearly to a constant, m a x i m u m value, AXmax, which corresponded to the calculated a m o u n t of undissociated acid in the original solution (within 0.5%). The m e t h o d is particularly suited to ~FeOOH titrations where the titration endpoints are not obvious (see Fig. 2) and may be further masked in titrations where salt has been added.
65
ION EXCHANGE OF OXIDE-HYDROXIDES
FIG. 1. Electron micrograph of flFeOOH crystals prepared by hydrolysis of ferric chloride at 100°C, by the method of Matijevic and Scheiner (13).
Journal of Colloid and Interface Science.
Vol.94, No. 1, July 1983
66
PATERSON AND R A H M A N RESULTS A N D DISCUSSION
The electron micrograph (Fig. 1) shows the characteristic shape of the 13FeOOH crystals, obtained by the hydrolysis of ferric chloride at 100°C by the method of Matijevic and Scheiner (13). Those obtained by hydrolysis at 90°C were identical to those shown here. Electron diffraction on single crystals gave only #FeOOH spacings (9). The crystals as shown, were very regular in size and shape. Their dimensions were 2000/~ (a +_ 400 A) in the direction of the c-axis, and 400 A (a + 70 h) in width. A number of titrations were made on #FeOOH with alkali and typical pH titration curves are shown in Fig. 2. In that figure there appear two distinct titration plots, which merge in alkaline solution above pH 10. The lower curve was obtained when #FeOOH was dispersed in water and titrated with sodium hydroxide, with no added sodium chloride present. The pH curve rises steeply from pH 3 to 9 and gives finally an indistinct end point, unsuitable for quantitative work. Titration data are given in Table I. If the standard sample of 0.303 g of air dried/3FeOOH was again dispersed in 25 ml of water and 12 ./, ?
lo
[3FeOOH
8 •
4
/! / , ' ~" / "# / "
pH TITRATION WITH
ADDEDSALT, NaCl
j'+/' /
~,,
2 0
i
I
1
2
mt OF BASE FIG. 2. pH titration curves for flFeOOH(100). Samples of the air dried material (0.3 g) were titrated with 0.1676 M NaOH. Initial point for all titrations pH 3.3. Lower curve with no added salt as Table I. Upper curve, after addition of I mmole NaCI (Table II). Dotted curve, after addition of 2 mmole NaC1. Arrows, transitions from lower to upper curves by addition of 1 mmole of NaC1 during titration. Journal of Colloid and InterfaceScience, Vol. 94, No. 1, July 1983
now 1 mmole of sodium chloride added (1 ml × 1 Msalt), the pH of the suspension rose by almost one pH unit (Table II and Fig. 2). Thereafter, on addition of base, the pH followed the upper curve to the end point. The transition from lower to upper curves could be made at any point by addition of salt. Three such titrations were made and the transitional jumps recorded as vertical arrows on Fig. 2. An addition of 2 mmole of sodium chloride gave only a marginal additional increase in pH (Fig. 2). Dispersions of #FeOOH in water usually gave a pH of approximately 3 but were variable from batch to batch, dependent upon the degree of washing during preparation. Solubility tests showed however that in the titration range of pH, greater than 3, soluble iron was not detectable and was therefore less than 10-6 M. In alkaline solution, at pH greater than 10, salt addition caused no change in the pH (other than the very small effect due to ionic strength effects). Salt addition in acidic solutions, however, caused proton uptake by t3FeOOH. This is a very sensitive test of uptake, since a large pH change, say between 5 and 6, as observed in Fig. 2, would amount to a transfer of only some 3 × 10-4 mmole of acid from the solution into the solid. A more exact account of the acid/base balance between solution and #FeOOH is obtained by calculation of AX, obtained from Eq. [7] and defined in the titration theory as the release of protons (mmole/g FeOOH), during the titration, taken relative to the initial capacity of the #FeOOH sample dispersed in water. AX values for the six titrations, shown in Fig. 2, were calculated. Data for the two titrations defining the entire upper and the lower curves are given in Tables I and II. The proton release (AX) on addition of base, although variable with salt addition, reaches a constant, maximum value for all samples. This AXmax expressed in mmole/g FeOOH, represents the capacity of the samples of #FeOOH(100) in equilibrium with water at pH 3. The true capacity of X, at
ION E X C H A N G E O F O X I D E - H Y D R O X I D E S
each titration point is A X m a x - - AX (Tables I and II). In the titration of t3FeOOH, the release of protons was shown to be accompanied by an equal amount of chloride, using chloride-reversible ion selective electrodes. This was confirmed by direct analysis of total chloride released in alkali and showed that under all conditions proton and chloride are released stoichiometrically as hydrochloric acid. Isotopic exchange of 36C1- under conditions of constant pH and salt concentration gave identical capacities. This also showed that chloride was freely exchangeable by an anion exchange process. Plots of capacity, X, against pH for the six titrations of Fig. 2 are shown in Fig. 3. Once more all the points lie on two curves and transitions between them made an addition of 1 mmole of salt, indicated by arrows. It is obvious that, at a given pH, the anion capacity is much larger if sodium chloride is added. The arrows, indicating transitions, also have slightly positive slopes because the process of salt addition raises pH and removes protons and chloride from solution to raise the capacity of the exchanger. Addition of twice as much salt has little further effect and is scarcely detectable. The capacity of/3FeOOH for acid (its anX -080 >-
0060
CA o.~o tLl O
CE 0 -d
I o
0.20
ion capacity also) is therefore clearly not a single-valued function of pH as implied in the literature on hydrous oxides (1, 2). It is obviously a function of both pH and salt concentration. In this series of papers on the ion exchange properties of microcrystals of hydrous oxides, similar salt/pH dependency has been observed in acidic (and alkaline) solutions with both aFeOOH (12) and monoclinic zirconia (15). Such effects have also been reported for other amphoteric colloidal oxides and often used to estimate the pH of the zero point of charge (z.p.c.). The literature does not however provide a simple explanation of the effects. In this paper a pore model mechanism is proposed. The model has thermodynamic consequences which are supported by the experimental data.
The Pore Model for Anionic Capacity of l3FeOOH The explanation for the major influence of chloride on capacity lies in the basic mechanism for exchange. This involves a thermodynamic coupling between the "protonation" reaction, by which positive charges are created in the matrix of the crystal and uptake of chloride (or other anion) to compensate these charges. The term protonation is not precisely defined, since the actual interaction between proton and ~FeOOH is not known. It is probably a reaction between a bound hydroxyl group, yielding a coordinated water (Eq. [9]). For thermodynamics it is only necessary to know that such an interaction exists. FeOOH + H + ~- FeO(OH2) +
o
I
I 5
I
• "~.~..,.~..2.~.. I I I 7
9
67
[9]
e." I
I 11
pH FIG. 3. Chloride capacity, X ( m m o l e / g F e O O H ) against p H using the titration data of Fig. 2. O, Initial point for all titrations, which is also the arbitrary comm o n reference for all 2xX calculations. Arrows represent additions of 1 m m o l e of salt. T h e u p p e r m o s t curve obtained after 2 m m o l e addition.
Bars over species in Eq. [9] indicate that the reactions are assumed to occur within the pores of 13FeOOH in this model. The pores may be the small hollandite pores (9) or the macropores of Gallagher (16). Equation [9] is a base protonation reaction, at least formally, for which an acid dissociation constant,/Ca, may be defined (Eq. [10]). Journal of Colloid and Interface Science, Vol. 94, No. I, July 1983
68
PATERSON AND RAHMAN
g a = (aFeoO H • dH+)/dFeOOHg.
[
10]
The protons of Eq. [9] are pore protons (H ÷) which are in a Donnan equilibrium with those of the external solution (H +) (Eq. [11 ]). H+ ~ H+
[11]
Because this is an interfacial equilibrium involving charge separation, and driven by the affinity of the protonation, the activities of the proton in each phase are not equal, but depend upon the Donnan potential (ff - ~)
(Eq. [1 la]) aH+/dH+ = exp(F(~/- ~ ) / R T ) .
[ 11 al
If only the equilibria of Eqs. [9] and [11] applied, then the protonation would not proceed because uncompensated transfer of positive charge from solution to pore would create an enormous positive Donnan potential (~ ~k), so that in effect aH would become zero (Eq. [1 la]). There would therefore be no protonation of the pore sites, regardless of the favorability of Ka (Eq. [10]) to protonation. The process can occur only if chloride (or other counterion) may also enter the pores. This allows effective electroneutrality to be achieved in the pores, while large numbers of protons (now accompanied by chloride) are transferred. The equilibrium between chloride in solution (CF) and in the -
Cl- ~ Cl-
[121
pore (CI-) (Eq. [I 2]) may be combined with the proton equilibrium (Eq. [11 ]) to give the well-known Donnan equilibrium relationship for electrolyte uptake of acid (Eq. [13]). aH+'acx- = dn+'dcl •
[13]
activities of the pore species, FeOOH~, CI-, and FeOOH, will be determined solely by the activity of the acid in solution, aH'acl. It is assumed that there is no salt uptake in the pore. Under such conditions, a particular value of the activity of acid, a H ' a o or equivalently in logarithmic form, pA = pH + pCl, however constituted by combinations of pH and chloride activities, would define only one pore condition and hence only one capacity. The experimental capacities equal the concentration of pore species FeOOH~ and chloride and would therefore be a single-valued function of all" a a or more conveniently, of pA. This conclusion is independent of the detailed thermodynamic model which might be used, since it is independent of the standard states which might be used or the method by which activity coefficients might be evaluated. It is noteworthy that protonation of surface sites would not require such coupling to the anion and there protonation should be solely a function of the pH of the external solution. Plots of capacity, X, against pA for the titration data of Fig. 2 and Tables I and II, are given in Fig. 4. The results show conclusively that chloride capacity for BFeOOH is a single-valued function of the activity of the total
X
>_I'--
Ka = -
-
dwooH " (aN+" acv) a-veOOH~"dcl-
[14]
Since, by definition, Ka is a constant, the Journal of Colloid and Interface Science, Vol. 94, No. 1, July 1983
o.6o
£3
0
.N
'~
"3,
0.40
(For simplicity the same standard states are assumed for pore and solution.) For Eqs. [10] and [13]
0.80
-lL)
0.20
E E
"k ":k
-,<. o I
I
7
I
I
4
g
i
I~
I
I
13
15
pA FIG. 4. Chloride capacity, X (mmole/g FeOOH), against pA (-log aH• acl), using the data of Fig. 2. Data presented refer to/3FeOOH(100) (13) samples prepared at 90°C/3FeOOH(90) by the same method and samples prepared by hydrolysis of ferric chloride at 25°C (17) gave identical results.
ION EXCHANGE OF OXIDE-HYDROXIDES acid, pA, as p r e d i c t e d b y t h e p o r e m o d e l above. C a p a c i t y a g a i n s t p A c u r v e s were det e r m i n e d u s i n g f l F e O O H (90) a n d w i t h /3FeOOH prepared by the method of Fryer et al. (17) f r o m 0.01 M ferric c h l o r i d e a l l o w e d to h y d r o l y z e at r o o m t e m p e r a t u r e . I n o n e case s a m p l e s h a d b e e n left for 7 years a n d in a n o t h e r t h e h y d r o l y s i s p e r i o d was a few m o n t h s . I n all these cases t h e c a p a c i t i e s for c h l o r i d e ( e x p r e s s e d as m m o l e / g F e O O H ) were i d e n t i c a l f u n c t i o n s o f p A a n d superp o s a b l e o n Fig. 4. T h e r e p r o d u c i b i l i t y o f these s a m p l e p r e p a r a t i o n s is t h e r e f o r e excellent a n d m a n y b a t c h e s w i t h i d e n t i c a l p r o p erties h a v e b e e n m a d e in the c o u r s e o f this work. S p e c u l a t i o n t h a t t h e l o c a t i o n o f exc h a n g e sites w i t h i n t h e h o l l a n d i t e p o r e s is n o w s u p p o r t e d b y e v i d e n c e o f a n i o n exclusion o f p e r c h l o r a t e , discussed in p a r t II (7). REFERENCES 1. Kraus, K. A., Phillips, H. O., Carlson, T. A., and Johnson, J. S., Proc. Second Int. Conf. "Peaceful Uses of Atomic Energy," United Nations, Geneva, 28, 3 (1958).
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2. Amphlett, C. B., "Inorganic Ion Exchangers." Elsevier, Amsterdam, 1964. 3. Clearfield, A., Nancollas, G. H., and Blessing, R. H., in "Ion Exchange and Solvent Extraction" (J. A. Marinsky and Y. Marcus, Eds.), Vol. 5, Ch. 1. Dekker, New York, 1973. 4. Vesely, V., and Pekarek, V., Talanta 19, 1219 (1972). 5. Nancollas, G. H., and Paterson, R., J. Inorg. Nucl. Chem. 29, 565 (1967). 6. Clearfield, A., and Stynes, J. A., J. Inorg. Nucl. Chem. 26, 117 (1964). 7. Paterson, R., and Rahman, H., Part II, J. Colloid Interface Sci., submitted for pubfication. 8. B/3hm, J., Z. Anorg. Allgem-Chem. 149, 203 (1925). 9. Mackay, A. L., Mineral Magn. 32, 545 (1960). 10. Mackay, A. L., Mineral Magn. 33, 270 (1962). 11. Mackay, A. L., J. Phys. Soc. Japan 17, 317 (1962). 12. Paterson, R., and Rahman, H., Part II1, J. Colloid Interface Sci., in press. 13. Matijevic, E., and Seheiner, P., J. Colloid Interface. Sci. 63, 509 (1978). 14. Davies, E. W., "Ion Association." Butterworths, London, 1962. 15. Paterson, R., and Rahman, H., Part IV, manuscript in preparation. 16. Gallagher, K. J., Nature (London) 226, 1225 (1970). 17. Fryer, J. R., Gildawie, A. M., and Paterson, R., Nature (London) 226, 149 (1974).
Journalof Colloidand InterfaceScience, Vol.94, No. 1, July 1983