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Adsorption of short-chained organic acids on stannic oxide
Colloids and Surfaces, 34 (1988/89) 255-269 Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands
255
A d s o r p t i o n of S h o r t - C h a i n e d Organic Acids on S t a n n i c Oxide J.B. FARROW and L.J. WARREN
CSIRO Division of Mineral Products, Curtin University Site, Perth, Western Australia (Australia) (Received 10 February 1988; accepted 2 August 1988)
ABSTRACT The uptake of styryl phosphonic acid (SPA) by stannic oxide was studied using three independent measures of the adsorption process: (i) adsorption isotherms, (ii) zeta potentials of stannic oxide particles before and after adsorption and (iii) the release of hydroxyl ions from the stannic oxide during SPA uptake. The uptake of structurally similar styryl sulfonic acid, benzene arsonic acid and benzene sulfonic acid was also measured. The results are consistent with exchange chemisorption between phosphonate anions and surface hydroxyl groups being the main adsorption mechanism for SPA between pH 3 and 5.
INTRODUCTION
The adsorption of short-chained organic acids, such as styryl phosphonic acid (SPA), CsHsCHCHPO (OH)2, make the surface of cassiterite, SnO2, hydrophobic and hence recoverable by the froth flotation process [ 1 ]. Whilst a number of phosphonic acids could be used, SPA has proved to be the most economic in practice [2 ]. However, the mechanism of action of SPA is not well understood. Wottgen [3] studied the uptake of n-heptyl phosphonic acid (HPA) by a natural cassiterite sample (93% SnO2, 0.62 m 2 g - l ) and concluded that the attachment of phosphonic acids to cassiterite was due to "exchange of OH ions on the cassiterite surface for phosphonate ions as a consequence of acid-base reaction, the strength of which is governed by the pH value. Physical coadsorption of phosphonic acid molecules and phosphonate ions must also be considered". The adsorbed species was said to be similar to Sn (IV) phosphonate, (RPO3) 2Sn, where R = CH3 (CH2) s. However, evidence for exchange chemisorption was indirect and hinged on the observation that heptyl phosphonate anions adsorbed on a negatively charged cassiterite surface. Under these conditions, electrostatic adsorption was said not to be possible and Wottgen invoked ion exchange with surface OH as a reasonable alternative. This was 0166-6622/89/$03.50
© 1989 Elsevier Science Publishers B.V.
256 supported by the IR spectra, although these results are not conclusive because of (i) the possibility of reaction occurring during preparation of KBr discs, and (ii) the poor quality of the spectra. Kuys and Roberts [4] obtained F T I R spectra of SPA and the Sn(II) and Sn (IV) phosphonate salts in KBr discs and compared these spectra with those of SPA and sodium styryl phosphonate in solution and SPA adsorbed on a thin film of Sn02 deposited on a germanium ATR crystal. The fine structure of the phosphonate peak of the adsorbate was interpreted as evidence for a bidentate complex, viz. R
I
O
iS
P
/\ \/ /\
O
- -
Sn
O
- -
Kuys and Roberts concluded that the formation of the bidentate complex and the slow rate of adsorption supported the ion exchange mechanism proposed by Farrow et al. [5 ]. In this paper the preliminary work of Farrow et al. [5] and Houchin [6] has been repeated. The interpretation and comparison of the results from these two earlier studies was difficult because three different stannic oxide samples had been used in the course of the investigations. We have used a BDH synthetic Sn02 as substrate in three independent measures of the adsorption process: adsorption isotherms, zeta potentials and the stoichiometry of the adsorption reaction. We performed our measurements at pH 3 and at pH 5 (approximately one pH unit either side of the isoelectric point) where the stannic oxide surface has a nett positive and negative charge, respectively. Of particular interest was the relative importance of chemical and physical forces in the uptake of SPA by Sn02. Therefore, organic acids structurally similar to SPA but with potentially different chemical affinities for SnO2 were also studied. These acids were benzene arsonic acid CsHsAsO3H2 (BAA), styryl sulfonic acid, C6HsCHCHSO3H (SSA) and benzene sulfonic acid, C6HsSO3H (BSA). EXPERIMENTAL Materials
A sample of BDH GPR stannic oxide (99.5%, As < 0.002%, Pb < 0.05% ) was used and had a B E T surface area (N2) of 6.5 m 2 g-1 and an isoelectric point (i.e.p.) at pH 4.1.
257
Benzene arsonic acid and the sodium salt of styryl sulfonic acid were obtained from Tokyo Kasei Kogyo; the sodium salt of benzene sulfonic acid was a BDH reagent. These reagents were used without further purification. Styryl phosphonic acid was obtained from Hoechst, and was purified by several recrystallizations from an acetone/diethyl ether mixture. The melting point of the purified material was 152-153 °C in good agreement with that reported by Schmutzler [7 ]. All other reagents were of Analytical Grade. Doubly distilled water was used. Styryl phosphonic acid is dibasic. Analysis of SPA titration curves by Houchin and Warren [8] indicated that pK1 ~ 2 and pK2 = 7.0, in agreement with Chavane [9] who gave pKI=2.00 and pK2=7.10. Thus the dominant species in the range o f p H studied in the present work (pH 3 to 5) is the monovalent anion. Houchin and Warren's [8 ] measurements also showed that BAA is dibasic with pK1 = 3.4 and pK2 = 8.4. These values are similar to those given for p-tolyl arsonic acid by Kortiim et al. [10], viz. pK1=3.77 and pK2=8.73. The acidity constants for BAA imply that the neutral acid species would be dominant at pH 3.0 (monoanion:neutral acid= 1:2.5) whilst the monoanion would be the most abundant species at pH 5.0 ( monoanion: neutral acid = 80:1 ). By contrast sulfonic acids are strong monoprotic acids (pKa < 1) so that the main species in SSA and BSA solutions between pH 3 and 5 would be the monoanion.
Organic acid analysis Organic acid concentrations before and after contact with stannic oxide suspensions were obtained from measurements of the optical density of the supernatant at the appropriate UV wavelength of the reagent's benzenoid chromophore, following the method of Houchin and Warren [8]. The wavelengths of maximum absorption ()[max) and the corresponding extinction coefficients (emax) for the four organic acids used are given in Table 1. As expected, SPA, SSA and BSA obeyed the Beer-Lambert Law and their extinction coefficients were independent of the pH in the pH range 3-5. HowTABLE 1 E x t i n c t i o n coefficients of organic acids
Organic
)lmax (nm)
6max
acid SPA SSA BAA BSA
256 256 262 262
20,600 15,450 ~ 650 380
258 ever, for BAA, e at 262 nm increased with decreasing pH, suggesting that at this wavelength the neutral acid form had a higher extinction coefficient than the monoanion. Optical densities were measured to three decimal places using a Unicam SP6-550 visible-UV spectrophotometer. Hence SPA and SSA concentrations could be measured down to about 5.10 -7 mol dm -3, BAA concentrations to 10 -5 mol dm -~ and BSA concentrations to about 2-10 -5 mol dm -~.
Measurement of adsorption isotherm The procedure used to determine the uptake of each reagent at approximately constant ionic strength is described below. Pyrex stoppered conical flasks (250 cm 3) were thoroughly cleaned, rinsed with doubly distilled water and dried. To each flask was added a known small volume of organic acid solution, 6 cm a of 10 -2 mol dm -3 KC1 solution (to maintain constant ionic strength) and sufficient doubly distilled water to give 60 cm 3 of solution. The pH of these three components was previously adjusted with either HC1 or KOH to the desired value (pH 3.0 or 5.0). A 10 ml aliquot was withdrawn from each flask and analysed to determine the initial organic acid concentration. Approximately 0.4 g of SnO2 (weighed accurately) was added to all flasks, except one which acted as a blank solution, prepared and treated in the same way as the others apart from the addition of SnO2. The SnO2-organic acid suspensions were allowed to equilibrate for 100 hours at ambient temperature (20 _+2 ° C ) during which time they were mechanically shaken in an equivalent fashion. After this time the suspensions were filtered through a membrane filter (Millipore, 0.22/~m ) to obtain the supernatant free of particulate SnO2. Although the adsorption process very slightly altered the supernatant pH, no readjustments were made.
Measurement of zeta potentials Electrokinetic measurements were made using a Rank Brothers Mark II Microelectrophoresis apparatus in the flat cell configuration. The i.e.p, of the BDH SnO2 powder was determined as follows. Approximately 0.08 g SnO2 powder was added to 500 cm 3 of 10 -2 mol dm -3 KC1 solution and dispersed with ultrasound. The suspension was equilibrated for 20 minutes, after which the pH was measured (5.6). An aliquot was removed and placed in the microelectrophoresis cell and thermostatted at 25.0 ° C. A total of 10 readings were taken, the mobilities determined, then averaged, from which the zeta potential (~) was calculated using the Smoluchowski equation. The suspension pH was then adjusted using 0.1 mol dm -3 HC1, and the zeta potential measured after 30 minutes at lower pH values. The procedure was repeated using 10- 3 and 10- 4 mol d m - 3 KC1 as background electrolyte.
259 The zeta potential of the BDH Sn02 sample after organic acid adsorption at pH 3.0 and 5.0 was determined as follows. Approximately 0.03 g of Sn02 was added to 200 cm 3 of 2.10-3 mol dm-3 KC1 and dispersed with ultrasound. The pH was adjusted to either pH 3.0 or 5.0 using HC1. To aliquots of this suspension was added an equal volume of solution at the same pH containing SPA, BAA, BSA or SSA at one of the following concentrations: 2.10-a, 2.10 -4 or 2-10 -5 mol dm -3. Each suspension was then equilibrated at ambient temperature for 100 hours, after which the zeta potential was measured as before.
Measurement of hydroxyl ion release 2.00 g of Sn02 was added to 100 cm 3 of 2-10 -3 tool dm -3 KC1 and dispersed with ultrasound. The pH was adjusted to 5.00 using HC1. The suspension was equilibrated for 36 hours at ambient temperature during which time the pH was periodically checked and found to be unchanged (pH 5.00 + 0.01 ). The pH of 100 cm 3 of a 2-10 -3 mol dm -8 SPA solution was adjusted to 5.00 with KOH. The solution was then equilibrated for 36 hours at ambient temperature during which time the pH remained at 5.00 _+0.01. These two samples were then combined and stirred for 120 hours at ambient temperature. The pH of the mixture was periodically readjusted to pH 5.00 by the measured addition of 0.010 mol dm -3 HC1 from a Radiometer automatic titration assembly. This volume was used to calculate the number of moles of acid required to maintain the pH at 5.00, which equates to the number of moles of O H - released during the adsorption of SPA onto SnO2. It was not possible to quantify the hydroxyl ion release at pH 3.0 for either SPA or BAA since the adsorption of these acids onto SnO2 at this pH caused only minute pH changes. The pH measurements were made using a Titron Instruments combined glass/reference electrode which was frequently standardized at pH 4.01 and 7.00 using commercial buffer solutions. RESULTS
Rate and amount of uptake ("adsorption") Neither of the sulfonic acids, SSA and BSA, showed any measureable uptake by stannic oxide under the range of conditions used. For SSA, the lowest measureable uptake would have been ~ 0.01/~mol m-2; for BSA ~ 0.4 #mol m -2. Both the phosphonic and arsonic acids were taken up by stannic oxide, to a degree which depended on their solution concentrations, time of contact and pH. Two tests were carried out over 120 hours to determine the rate of adsorption
260
of SPA onto Sn02, one at pH 3.0 and one at pH 5.0, using an initial SPA concentration of 10 -3 mol dm -3 (Fig. 1 ). At both pH 3.0 and 5.0, the solution concentration of SPA decreased relatively slowly upon contact with SnO2, corresponding to the slow adsorption of SPA onto SnO2. At pH 5.0, the uptake after 2 hours was about 80% of the equilibrium amount and about 90% after 4 hours. Uptake continued however for at least 30-40 hours. At pH 3.0, the uptake after 2 hours was only about 50% of the equilibrium amount; after 4 hours this had increased to about 60%. The adsorption process at pH 3.0 required approximately 50-60 hours to reach completion. Similar tests were performed to determine the rate of adsorption of BAA onto SnO2 at pH 3.0 and pH 5.0. However at pH 5.0 there was evidence that the adsorption process is not straightforward. It was found that the interaction of BAA with Sn02, although giving the expected decrease in the solution concentration of BAA (as judged from the slight decrease in the relative peak heights in the BAA spectrum, see Fig. 2), also resulted in the release into solution of one or more UV active species as evident from the broad adsorption band developing with time below 380 nm (Fig. 2). These UV active species were presumably the products from the decomposition of BAA which was apparently catalysed by SnO2, since equivalent BAA solutions were found to be totally stable at pH 5.0 in the absence of SnO2 as were aqueous suspensions of Sn02. This phenomenon made determination of the true adsorption of BAA onto SnO2 at pH 5.0 uncertain and as such no data for the adsorption of BAA at this pH is reported here. The rate of adsorption of BAA onto SnO2 at pH 3.0 for an initial BAA concentration of 10 -3 mol dm -3 is shown in Fig. 3. The uptake after 2 hours was about 50% of the equilibrium amount; after 4 hours about 70%. Adsorption of BAA at pH 3.0 continued for approximately 30-40 hours before equilibrium was achieved. There was no evidence of the decomposition of BAA at pH 3.0 in the presence of SnO2. The adsorption isotherms of SPA at pH 3.0 and pH 5.0 and BAA at pH 3.0
10.2
9.8
o c
-~ E
)~o-o
o
o-
pH 5.0
9.4
g ,o pH 3.0
8.6 I
/ 40
I Times
I 80
I
I 120
(hrs)
Fig. 1. Rate of adsorption of SPA onto SnO2 at pH 3.0 and pH 5.0.
261 1.0
0.8
0.6 .a
0.4
0.2
0
I
i
I
i
i
I
240
1.0
280 Wavelength
320
360
400
(nm)
0.8
0.6
o
<
0.4
o2
0
I
240
I
280 Wavelength
320 (nm)
360
400
Fig. 2. UV spectra (230-300 nm) of BAA solutions after contact with Sn02 at pH 5.0. (a) 0 hours, (b) 2 hours, (c) 24 hours.
were determined after 100 hours contact with SnO2 and are given in Figs 4 and 5, respectively. The shapes of the SPA adsorption isotherms are similar at pH 3.0 and pH 5.0. As the pH increased from 3 to 5, the amount of SPA adsorbed at a constant residual solution concentration decreased. The amount of SPA adsorbed decreased less than the corresponding increase in the hydroxyl ion concentration, e.g. at 10 -3 mol dm -3 the amount adsorbed decreased ~ 7-9 times (2.45 to 0.33/~mol m -2) for a 100-fold increase in [ O H - ] (pH 3 to 5); with similar results at 5-10 -4 and 10 -4 mol dm -3 where the amount adsorbed also de-
262
to -6 E
o 5
8
'
~ 7 c_
~
'
4'0
'
8
'o
'
'
120
Time (hrs)
Fig. 3. Rate of adsorption of BAA onto Sn02 at pH 3.0.
/
4 E "5 E
3
10 H 30
E
6
o
c
8
E
4
pH 5.0
E
E 0
I -6
-5
-4
-3
I -2
log (Residual solution concentration (tool din-3))
-5
I -4
L -3
-2
log (Residual solution concentration (mol dm-3))
Fig. 4. (left) Adsorption isotherms for SPA after 100 hours contact with SnO2 at pH 3.0 and 5.0. Points shown by (m) are from data given in Fig. 1. Fig. 5. (right) Adsorption isotherm for BAA after 100 hours contact with Sn02 at pH 3.0. Point shown by (m) is from the data given in Fig. 3.
creased ~ 7 - 9 times (1.80 to 0.20 Hmol m -2 and 1.06 to 0.13 Hmol m -2) for the same increase in pH. The BAA adsorption isotherm at pH 3.0 was similar in shape to that of SPA at the same pH. However BAA adsorbed more strongly, the amount of uptake being approximately twice that of SPA for the same residual solution concentration and pH. For example at 10 -2 and 10 -4 mol dm -3, the uptake of BAA was 1.9 and 2.5 times that of SPA (4.70 to 2.45 and 2.70 to 1.06 Hmol m -2 respectively) at pH 3.0.
Effect of adsorption on zeta potentials The i.e.p, of the B D H SnO2 was determined to lie at pH 4.1 from the intersection of the ~ versus pH curves at 10 -2, 10 -3 and 10 -4 mol dm -3 KC1 where ~ = 0. The B D H Sn02 value compares well with the value o f p H 4.5 determined previously by Houchin and Warren [11 ] for SnO2.
263
.
2O
.
.
.
lO o s%
¢1 -t0
~u
-20
-so
--@ .... -~^=o=---~o---o~-;~ I
I
I
0
10-5
10-4
Organic acid concentration
50
I 10 -3 (mol dm - 3 )
Fig. 6. Zeta potential of S n 0 2 a s a function of organic acid concentration (10-8, 10-4 10-5, 0 tool dm -3) at pH 3.0 [SPA (@), BAA ( T ) , SSA (m), BSA (A), none ( ~ ) ] and at pH 5.0 [SPA ( O ), BAA ( V ), SSA ([]), BSA (/x ), none (~>) ]. 0.4
0.3 E
0.1 0 0
0
0
120
SPA adsorption time (hrs)
Fig. 7. Rate of release of hydroxyl ions during the adsorption of SPA (10-3 mol din-3) onto Sn02 at pH 5.0.
Figure 6 shows the effect of adding SPA, BAA, BSA and SSA to suspensions of B D H SnO2. Zeta potentials were measured after 100 hours contact between the organic acid (at 10 -3, 10 -4 and 10 -5 mol dm -3) and SnO2, in 10 -3 mol dm -3 KC1. At pH 5.0, the values in the absence and in the presence of organic acid concentrations up to 10 -3 mol dm -3 were not significantly different and were probably within experimental error ( + 3 mV). This conclusion is supported by the fact that there was no trend towards higher or lower zeta potentials as the acid concentration was increased. At pH 3.0, treatment of B D H SnO2 with SPA or BAA made the particles more negative, SPA changing the zeta potential from + 16 to - 1 0 mV and BAA changing the value from + 11 to +2.5 mV over the range 10 -5 to 10 -3 mol dm -3, compared to 21 mV in the absence of the organic acids. BSA and SSA had no effect (Fig. 6. )
Release of hydroxyl ions As SPA was adsorbed by SnO2, hydroxyl ions were released, as shown in Fig. 7 for the adsorption of 10 -3 mol dm -3 SPA onto SnO2 at pH 5.0. The rate of
264 release of hydroxyls slowed as the adsorption proceeded and was virtually zero after approximately 72 hours. The total amount of O H - released during this adsorption experiment was 0.36/lmol m -2 which can be compared to the SPA adsorption density at pH 5.0 (see Fig. 4) for a 10 -3 mol dm -~ SPA solution of 0.33/~mol dm -3. DISCUSSION Adsorption versus reaction at the surface
The observed decrease in the solution concentration of SPA after it is mixed with stannic oxide may result from one or more of the following processes: (i) adsorption of phosphonate anions onto the SnO2 surface without lattice rearrangement, (ii) reaction with "bulk SnO2" near the surface, probably after lattice rearrangement, to form a separate phase of tin (IV) phosphonate or a mixed tin (IV) oxide-phosphonate, (iii) a precipitation reaction between phosphonate anions and cationic species in solution, e.g., dissolved Sn. The last process, (iii), is improbable because of the very insoluble nature of stannic oxide at normal pH values [12]. However, process (ii), reaction with bulk SnO2, is a possibility since the uptake of SPA was slow requiring at least 30 hours to reach completion, whilst the accompanying release of hydroxyl ions took ~ 70 hours to reach completion. Simple adsorption processes (process (i)) are not normally so slow unless the adsorption involves a chemical interaction with a significant activation energy. The shape of the apparent absorption isotherm can be useful in distinguishing adsorption, process (i), from reaction, process (ii). Simple adsorption is often indicated by the isotherm starting off as Henry's Law and becoming less sensitive to concentration, often reaching a definite plateau as found by Edwards and Ewers [13] for uptake of sodium hexadecyl sulfate (SHDS) by stannic oxide at pH 2.6. Reaction with the bulk is often indicated by the isotherm increasing steadily with increasing concentration, well past the point of equivalent monolayer coverage, as found by Zimmels et al. [ 14 ] and Kitchener [15] for oleate on calcite or Wottgen [3] for HPA on cassiterite at pH 2. Our isotherms for SPA and BAA match neither of the above extreme cases. Total uptake of SPA was always less than that of a close-packed monolayer (assuming a molecule of SPA occupies an area of ~ 0.25 nm 2 on the SnO2 surface, then close-packed monolayer coverage would correspond to an adsorption density of approximately 6.5/~mol m -2). Another argument against reaction as an uptake process for SPA and BAA is that there are many examples of adsorption processes which produce isotherms of unusual shape a n d / o r without a saturation plateau, e.g. sodium dodecyl sulfonate on alumina [16], cobalt cations on silica [17], fluoride on alumina [18,19]. In some of these cases the adsorption mechanism changed as
265 the solution concentration increased, or as the pH changed, leading in turn to a change in isotherm shape. In summary, whilst the data on isotherm shape and magnitude do not allow us to decide unequivocally between adsorption and reaction with the bulk, we believe the evidence favours adsorption as the main uptake process. This is strongly supported by the IR spectra of Kuys and Roberts [4]. Reaction, if it occurs, is probably more likely at low pH and with BAA than with SPA, given the known reactivity of phenylarsonic acid with quadrivalent metal ions in acidic media [20 ].
Mechanism of adsorption of SPA At pH 5.0, the main SPA species in solution is the monoanion, RPO (OH) O -, and the surface of SnO2 is negatively charged ( - 2 8 mV, see Fig. 6), yet significant adsorption occurred against the electrostatic repulsion (Fig. 4). The mechanism therefore must involve a chemical interaction, with or without hydrophobic association. The chain length of SPA is only 5.5 equivalent CH2 groups which is shorter than the critical value of 6 equivalent CH2 groups required before hydrophobic association substantially enhances adsorption [21]. Thus, the adsorption mechanism is primarily a chemical interaction between SPA and the SnO2 surface. Further, since the structurally similar sulfonates SSA and BSA were not adsorbed at any pH value, the chemical interaction is specific for phosphonate and arsonate head groups. It was expected that at low pH, organic acid anions would be adsorbed into the electrical double layer around the positively charged SnO2 particles. If such electrostatic adsorption occurred at pH 3.0, it was not detected and must therefore be less than 0.01 Hmol m-2. Thus, electrostatic adsorption cannot account for the amount of SPA adsorbed which was of the order of i Hmol m -2. The key observation in explaining the nature of the adsorption process was the release of hydroxyl ions which accompanied the uptake of SPA, but quantification of the exchange reaction was difficult because the total amount of OH - released was small. At pH 5.0, for each RPO ( OH ) O - adsorbed, one OH (within experimental error) was released, suggesting that the adsorption process can be described by an equation of the type: -SnOH(s~ + RPO (OH) O~-~)~--SnO (OH) OPR(s) + OH ~-d~)
(1)
Equation (1) refers to an equilibrium between surface species (s) and exchanging anions in the double layer (dl) adjacent to the surface. Our adsorption isotherms for SPA at pH 3 and 5 (see Fig. 4) show that equilibrium for this ion exchange reaction lies well towards the side of the reactants, i.e. it is very difficult to exchange phosphonate ions for surface hydroxyl groups. At pH 5.0, zeta potential measurements of SnO2 before and after treatment with SPA were, within experimental error, the same irrespective of the con-
266 centration of SPA (Fig. 6). This is consistent with Eqn (1). At pH 3.0, the zeta potential of the Sn02 particles became increasingly negative as more SPA was adsorbed. However, the overall change was small (see Fig. 6), and this in a region where the zeta potential is very sensitive to small changes in surface charge density [11 ]. Neutralization of the positive surface charge, which arises from a small number of - S n O H + groups (the dominant surface species is always -SnOH; see Houchin and Warren [11 ] ) is consistent with reactions such as:
- S n O H + + RPO (OH) O- ~ - S n O (OH) OPR + H20
(2 )
- S n O H + + RPO (OH) 2 ~ - S n O (OH) OPR + H3 0 +
(3)
If we assume that Eqn (1) alone describes the adsorption of phosphonate and has a pH-independent equilibrium constant, then SPA adsorption should fall in proportion to the increase in [ O H - ] , as Edwards and Ewers [13] observed for SHDS on SnO2. We did not observe this behaviour for SPA or BAA nor did Wottgen [3 ] for HPA. Our results show that adsorption of phosphonic and arsonic acids by stannic oxide is not as sensitive to pH as expected from simple ion exchange, but the reasons for this are not clear. With diprotic acids such as SPA and BAA it is possible that the initial onepoint attachment to the SnO2 surface is followed by ring closure to form the surface analogue of a bidentate chelate, as shown below: 0
O-
HO
\1/
+
/X R
Sn
.o/ OH
I\
0 ~,,
\\/°\1/ P
Sn
/\
+.
OH-
+
H20
.o/I x
R
OH
p R
/°\ I/
/\/I
Sn
0
\
Similar structures have been proposed for the adsorption of phosphate onto goethite [22,23]. Our evidence does not allow us to distinguish between the bidentate bridging complex and the bidentate chelation of SPA to a single surface Sn atom, although recent F T I R spectra favour the latter [4 ]. Ring closure should be favoured by the fact that the second proton of SPA would become more acidic after SPA was adsorbed. Ring closure might still be a slow process and contribute to the overall slow rate of uptake of SPA.
Mechanism of adsorption of BAA The shape of the adsorption isotherm for BAA on stannic oxide at pH 3.0 is similar to that for SPA, although the magnitude of adsorption is about two
267 times greater. The differences in adsorption density may be related to the greater abundance of the neutral arsonic acid species, RAsO (OH)2, and/or a difference in the chemical reactivity of arsonic and phosphonic acids towards -SnOH. Presumably the major adsorption reaction for the arsonic acid at pH • 3.0 is Eqn (5) rather than ion exchange following Eqn (1): -SnOH + RAsO (OH) 2 ~--SnO (OH) OAsR + H20
(5)
Equation (5) is an acid-base reaction in which the stannic oxide surface acts as the base. The apparent catalytic effect of stannic oxide on the decomposition of BAA at pH 5.0 (Fig. 2 ) but not at pH 3.0 requires further investigation before the mechanism of adsorption of BAA onto SnO2 can be established more precisely. GENERALCOMMENTS The uptake of organic acid anions onto a negatively charged SnO2 surface, the dependence of uptake on the nature of the head group (sulfonates not adsorbed) and the release of about one O H - for each SPA anion adsorbed are all consistent with ion exchange chemisorption, Eqn ( 1 ), even though Eqn (1) is not obeyed quantitatively. A number of other systems show exchange chemisorption. Phosphate anions chemisorb on iron oxide, by exchange with surface hydroxyl groups, to produce a bidentate bridging complex like that proposed for the phosphonate in Eqn (4) [22]. Sulfate anions and alkyl sulfates chemisorb on iron oxides by ion exchange with surface OH [24,25] showing that exchange chemisorption occurs even when an organic tail is attached to the head group. The long-chained SHDS adsorbed by ion exchange with surface OH on stannic oxide [13], although hydrophobic association between the organic tails presumably enhanced the adsorption. Marinakis and Kelsall [26] concluded that dodecyl sulphate and decyl phosphonate ions adsorb via anion exchange with tungstate anions from scheelite (CAW04). These examples show that exchange chemisorption of anionic surfactants on oxides and salt-type minerals may be of more general importance than previously believed. The question remains as to why the sulphonates, SSA and BSA, do not chemisorb onto SnO2 whereas phosphonates, arsonates and sulphates do. Two factors may be involved: (i) sulphonates are weaker ligands for metal ions in aqueous solutions than phosphonates, arsonates and sulphates [27], and (ii) ring closure or bidentate chelation to a single surface Sn atom is not possible with monoprotic sulphonic acids.
268 CONCLUSIONS (1) T h e main process by which suspensions of stannic oxide abstracted styryl phosphonic acid from aqueous solutions at p H 5 was t h a t of exchange chemisorption between p h o s p h o n a t e m o n o a n i o n s and hydroxyl groups a t t a c h e d to tin atoms in the oxide surface. (2) For each p h o s p h o n a t e anion adsorbed at p H 5, approximately one hydroxyl ion was released. However, not all results were consistent with a simple exchange process, e.g. changing the p H did not cause a proport i onat e change in the a m o u n t of SPA adsorbed. (3) T h e a m o u n t of SPA adsorption was relatively low { ~ 20% monolayer) reflecting the preference of surface tin atoms for hydroxyl over phosphonate. Also, adsorption was slow, requiring at least 20 hours. (4) At p H 3.0, the uptake of benzene arsonic acid by stannic oxide was approximately twice t h a t of SPA but the m e c ha ni sm of adsorption was not established. (5) T h e short-chained sulfonic acids SSA and BSA were not adsorbed to a significant e x t e n t ( < 0.01/~mol m -2) consistent with their lower complexing ability with metal ions compared with phosphonates, arsonates and sulfates. (6) Exchange chemisorption of anionic surfactants with oxides and salttype minerals may be of more general importance t h a n previously believed.
REFERENCES 1 G.D. Senior and G.W. Poling, The chemistry of cassiterite flotation, in P. Somasundaran (Ed.), Advancesin Mineral Processing,A Half Century of Progress in Applicationand Practice, SME-AIME, Littleton, CO, 1986, pp. 229-254. 2 E. Wottgen, Freiberg. Forschungsh. A, 621 (1980) 3. 3 E. Wottgen, Trans. Inst. Min. Metall. Sect. C, 75 (1969) 91. 4 K.J. Kuys and N.K. Roberts, Colloids Surfaces, 24 (1987) 1. 5 J.B. Farrow, M.R. Houchin and L.J. Warren, Mechanism of adsorption of organic acids on stannic oxide. Abstract of paper deliveredat the 5th Australian Conferenceon Colloidsand Surfaces, Brisbane, Feb. 1986. 6 M.R. Houchin, Surface Chemical Studies Relevant to Cassiterite Flotation Using Modem Analytical Techniques, in P.K. Jena and K.S. Narasimhan (Eds), Proceedings 2nd International Symposium on Beneficiation and Agglomeration,Bhubaneswar, 17-19 Dec. 1986 (INSDOC, New Delhi, 1986), paper ILl, pp. 65-71. 7 R. Schmutzler, Inorg. Chem., 3 (1964) 410. 8 M.R.Houchin and L.J. Warren, Tin flotation: the surface chemistry of Australian cassiterite ore minerals (Final Report on AMIRA Project 79/Pl14). Investigation Report (Restricted) No. IR 1532R, CSIRO Division of Mineral Chem., Port Melbourne, Australia, 1984, 169 pp. 9 V. Chavane, Ann. Chim. (Paris), 12 (1949) 383. 10 G. Korttim, W. Vogeland W. Andrusso, Dissociation Constants of Organic Acids in Aqueous Solution, Butterworths, London, 1961. 11 M.R.Houchin and L.J. Warren, J. ColloidInterface Sci., 100 (1984) 278. 12 N.N. Greenwoodand A. Eamshaw, Chemistry of the Elements, Pergamon, Oxford, 1984, p. 447.
269 13 G.R. Edwards and W.E. Ewers, Austral. J. Sci. Res. Ser. A, 4 (1951) 627. 14 Y. Zimmels, I.J. Lin and J.P. Friend, Colloid Polym. Sci., 253 (1975) 404. 15 J.A. Kitchener, The froth flotation process: past, present, and future - - in brief, in K.J. Ives (Ed.), The Scientific Basis of Flotation, Martinus Nijhoff, The Hague, 1984, pp. 3-51. 16 P. Somasundaran and D.W. Fuerstenau, J. Phys. Chem., 70 (1966) 90. 17 R.O. James and T.W. Healy, J. Colloid Interface Sci., 40 (1972) 42. 18 L.J. Warren and J.A. Kitchener, Trans. Inst. Min. Metall. Sect. C, 81 (1972) 137. 19 F.J. Hingston, A review of anion adsorption, in M.A. Anderson and A.J. Rubin (Eds.), Adsorption of Inorganics at Solid-Liquid Interfaces, Ann Arbor Science, Michigan, 1981, pp. 51-90. 20 I.M. Kolthoff, E.B. Sandell and E.J. Meehan, Quantitative Chemical Analysis, 4th edn, Macmillan, London, 1969, p. 269. 21 M.J. Rosen, Surfactants and Interfacial Phenomena, Wiley-Interscience, New York, 1978, p. 26. 22 R.J. Atkinson, R.L. Parfitt and R.St.C. Smart, J. C~em. Soc. Faraday Trans. 1, 70 (1974) 1472. 23 P.W. Schindler, Surface complexes at oxide-water interfaces, in M.A. Anderson and A.J. Rubin (Eds), Adsorption of Inorganics at Solid-Liquid Interfaces, Ann Arbor Science, Michigan, 1981, pp. 1-50. 24 H.L. Shergold and O. Mellgren, Trans. Inst. Min. Metall. Sect. C, 78 (1969) 121. 25 L. Sigg and W. Stumm, Colloids Surfaces, 2 (1981) 101. 26 K.I. Marinakis and G.H. Kelsall, Colloids Surfaces, 25 (1987) 369. 27 J. MacB. Harrowfield, Dept. Phys. Inorg. Chem., Univ. Western Australia, personal communication, 1986.
255
A d s o r p t i o n of S h o r t - C h a i n e d Organic Acids on S t a n n i c Oxide J.B. FARROW and L.J. WARREN
CSIRO Division of Mineral Products, Curtin University Site, Perth, Western Australia (Australia) (Received 10 February 1988; accepted 2 August 1988)
ABSTRACT The uptake of styryl phosphonic acid (SPA) by stannic oxide was studied using three independent measures of the adsorption process: (i) adsorption isotherms, (ii) zeta potentials of stannic oxide particles before and after adsorption and (iii) the release of hydroxyl ions from the stannic oxide during SPA uptake. The uptake of structurally similar styryl sulfonic acid, benzene arsonic acid and benzene sulfonic acid was also measured. The results are consistent with exchange chemisorption between phosphonate anions and surface hydroxyl groups being the main adsorption mechanism for SPA between pH 3 and 5.
INTRODUCTION
The adsorption of short-chained organic acids, such as styryl phosphonic acid (SPA), CsHsCHCHPO (OH)2, make the surface of cassiterite, SnO2, hydrophobic and hence recoverable by the froth flotation process [ 1 ]. Whilst a number of phosphonic acids could be used, SPA has proved to be the most economic in practice [2 ]. However, the mechanism of action of SPA is not well understood. Wottgen [3] studied the uptake of n-heptyl phosphonic acid (HPA) by a natural cassiterite sample (93% SnO2, 0.62 m 2 g - l ) and concluded that the attachment of phosphonic acids to cassiterite was due to "exchange of OH ions on the cassiterite surface for phosphonate ions as a consequence of acid-base reaction, the strength of which is governed by the pH value. Physical coadsorption of phosphonic acid molecules and phosphonate ions must also be considered". The adsorbed species was said to be similar to Sn (IV) phosphonate, (RPO3) 2Sn, where R = CH3 (CH2) s. However, evidence for exchange chemisorption was indirect and hinged on the observation that heptyl phosphonate anions adsorbed on a negatively charged cassiterite surface. Under these conditions, electrostatic adsorption was said not to be possible and Wottgen invoked ion exchange with surface OH as a reasonable alternative. This was 0166-6622/89/$03.50
© 1989 Elsevier Science Publishers B.V.
256 supported by the IR spectra, although these results are not conclusive because of (i) the possibility of reaction occurring during preparation of KBr discs, and (ii) the poor quality of the spectra. Kuys and Roberts [4] obtained F T I R spectra of SPA and the Sn(II) and Sn (IV) phosphonate salts in KBr discs and compared these spectra with those of SPA and sodium styryl phosphonate in solution and SPA adsorbed on a thin film of Sn02 deposited on a germanium ATR crystal. The fine structure of the phosphonate peak of the adsorbate was interpreted as evidence for a bidentate complex, viz. R
I
O
iS
P
/\ \/ /\
O
- -
Sn
O
- -
Kuys and Roberts concluded that the formation of the bidentate complex and the slow rate of adsorption supported the ion exchange mechanism proposed by Farrow et al. [5 ]. In this paper the preliminary work of Farrow et al. [5] and Houchin [6] has been repeated. The interpretation and comparison of the results from these two earlier studies was difficult because three different stannic oxide samples had been used in the course of the investigations. We have used a BDH synthetic Sn02 as substrate in three independent measures of the adsorption process: adsorption isotherms, zeta potentials and the stoichiometry of the adsorption reaction. We performed our measurements at pH 3 and at pH 5 (approximately one pH unit either side of the isoelectric point) where the stannic oxide surface has a nett positive and negative charge, respectively. Of particular interest was the relative importance of chemical and physical forces in the uptake of SPA by Sn02. Therefore, organic acids structurally similar to SPA but with potentially different chemical affinities for SnO2 were also studied. These acids were benzene arsonic acid CsHsAsO3H2 (BAA), styryl sulfonic acid, C6HsCHCHSO3H (SSA) and benzene sulfonic acid, C6HsSO3H (BSA). EXPERIMENTAL Materials
A sample of BDH GPR stannic oxide (99.5%, As < 0.002%, Pb < 0.05% ) was used and had a B E T surface area (N2) of 6.5 m 2 g-1 and an isoelectric point (i.e.p.) at pH 4.1.
257
Benzene arsonic acid and the sodium salt of styryl sulfonic acid were obtained from Tokyo Kasei Kogyo; the sodium salt of benzene sulfonic acid was a BDH reagent. These reagents were used without further purification. Styryl phosphonic acid was obtained from Hoechst, and was purified by several recrystallizations from an acetone/diethyl ether mixture. The melting point of the purified material was 152-153 °C in good agreement with that reported by Schmutzler [7 ]. All other reagents were of Analytical Grade. Doubly distilled water was used. Styryl phosphonic acid is dibasic. Analysis of SPA titration curves by Houchin and Warren [8] indicated that pK1 ~ 2 and pK2 = 7.0, in agreement with Chavane [9] who gave pKI=2.00 and pK2=7.10. Thus the dominant species in the range o f p H studied in the present work (pH 3 to 5) is the monovalent anion. Houchin and Warren's [8 ] measurements also showed that BAA is dibasic with pK1 = 3.4 and pK2 = 8.4. These values are similar to those given for p-tolyl arsonic acid by Kortiim et al. [10], viz. pK1=3.77 and pK2=8.73. The acidity constants for BAA imply that the neutral acid species would be dominant at pH 3.0 (monoanion:neutral acid= 1:2.5) whilst the monoanion would be the most abundant species at pH 5.0 ( monoanion: neutral acid = 80:1 ). By contrast sulfonic acids are strong monoprotic acids (pKa < 1) so that the main species in SSA and BSA solutions between pH 3 and 5 would be the monoanion.
Organic acid analysis Organic acid concentrations before and after contact with stannic oxide suspensions were obtained from measurements of the optical density of the supernatant at the appropriate UV wavelength of the reagent's benzenoid chromophore, following the method of Houchin and Warren [8]. The wavelengths of maximum absorption ()[max) and the corresponding extinction coefficients (emax) for the four organic acids used are given in Table 1. As expected, SPA, SSA and BSA obeyed the Beer-Lambert Law and their extinction coefficients were independent of the pH in the pH range 3-5. HowTABLE 1 E x t i n c t i o n coefficients of organic acids
Organic
)lmax (nm)
6max
acid SPA SSA BAA BSA
256 256 262 262
20,600 15,450 ~ 650 380
258 ever, for BAA, e at 262 nm increased with decreasing pH, suggesting that at this wavelength the neutral acid form had a higher extinction coefficient than the monoanion. Optical densities were measured to three decimal places using a Unicam SP6-550 visible-UV spectrophotometer. Hence SPA and SSA concentrations could be measured down to about 5.10 -7 mol dm -3, BAA concentrations to 10 -5 mol dm -~ and BSA concentrations to about 2-10 -5 mol dm -~.
Measurement of adsorption isotherm The procedure used to determine the uptake of each reagent at approximately constant ionic strength is described below. Pyrex stoppered conical flasks (250 cm 3) were thoroughly cleaned, rinsed with doubly distilled water and dried. To each flask was added a known small volume of organic acid solution, 6 cm a of 10 -2 mol dm -3 KC1 solution (to maintain constant ionic strength) and sufficient doubly distilled water to give 60 cm 3 of solution. The pH of these three components was previously adjusted with either HC1 or KOH to the desired value (pH 3.0 or 5.0). A 10 ml aliquot was withdrawn from each flask and analysed to determine the initial organic acid concentration. Approximately 0.4 g of SnO2 (weighed accurately) was added to all flasks, except one which acted as a blank solution, prepared and treated in the same way as the others apart from the addition of SnO2. The SnO2-organic acid suspensions were allowed to equilibrate for 100 hours at ambient temperature (20 _+2 ° C ) during which time they were mechanically shaken in an equivalent fashion. After this time the suspensions were filtered through a membrane filter (Millipore, 0.22/~m ) to obtain the supernatant free of particulate SnO2. Although the adsorption process very slightly altered the supernatant pH, no readjustments were made.
Measurement of zeta potentials Electrokinetic measurements were made using a Rank Brothers Mark II Microelectrophoresis apparatus in the flat cell configuration. The i.e.p, of the BDH SnO2 powder was determined as follows. Approximately 0.08 g SnO2 powder was added to 500 cm 3 of 10 -2 mol dm -3 KC1 solution and dispersed with ultrasound. The suspension was equilibrated for 20 minutes, after which the pH was measured (5.6). An aliquot was removed and placed in the microelectrophoresis cell and thermostatted at 25.0 ° C. A total of 10 readings were taken, the mobilities determined, then averaged, from which the zeta potential (~) was calculated using the Smoluchowski equation. The suspension pH was then adjusted using 0.1 mol dm -3 HC1, and the zeta potential measured after 30 minutes at lower pH values. The procedure was repeated using 10- 3 and 10- 4 mol d m - 3 KC1 as background electrolyte.
259 The zeta potential of the BDH Sn02 sample after organic acid adsorption at pH 3.0 and 5.0 was determined as follows. Approximately 0.03 g of Sn02 was added to 200 cm 3 of 2.10-3 mol dm-3 KC1 and dispersed with ultrasound. The pH was adjusted to either pH 3.0 or 5.0 using HC1. To aliquots of this suspension was added an equal volume of solution at the same pH containing SPA, BAA, BSA or SSA at one of the following concentrations: 2.10-a, 2.10 -4 or 2-10 -5 mol dm -3. Each suspension was then equilibrated at ambient temperature for 100 hours, after which the zeta potential was measured as before.
Measurement of hydroxyl ion release 2.00 g of Sn02 was added to 100 cm 3 of 2-10 -3 tool dm -3 KC1 and dispersed with ultrasound. The pH was adjusted to 5.00 using HC1. The suspension was equilibrated for 36 hours at ambient temperature during which time the pH was periodically checked and found to be unchanged (pH 5.00 + 0.01 ). The pH of 100 cm 3 of a 2-10 -3 mol dm -8 SPA solution was adjusted to 5.00 with KOH. The solution was then equilibrated for 36 hours at ambient temperature during which time the pH remained at 5.00 _+0.01. These two samples were then combined and stirred for 120 hours at ambient temperature. The pH of the mixture was periodically readjusted to pH 5.00 by the measured addition of 0.010 mol dm -3 HC1 from a Radiometer automatic titration assembly. This volume was used to calculate the number of moles of acid required to maintain the pH at 5.00, which equates to the number of moles of O H - released during the adsorption of SPA onto SnO2. It was not possible to quantify the hydroxyl ion release at pH 3.0 for either SPA or BAA since the adsorption of these acids onto SnO2 at this pH caused only minute pH changes. The pH measurements were made using a Titron Instruments combined glass/reference electrode which was frequently standardized at pH 4.01 and 7.00 using commercial buffer solutions. RESULTS
Rate and amount of uptake ("adsorption") Neither of the sulfonic acids, SSA and BSA, showed any measureable uptake by stannic oxide under the range of conditions used. For SSA, the lowest measureable uptake would have been ~ 0.01/~mol m-2; for BSA ~ 0.4 #mol m -2. Both the phosphonic and arsonic acids were taken up by stannic oxide, to a degree which depended on their solution concentrations, time of contact and pH. Two tests were carried out over 120 hours to determine the rate of adsorption
260
of SPA onto Sn02, one at pH 3.0 and one at pH 5.0, using an initial SPA concentration of 10 -3 mol dm -3 (Fig. 1 ). At both pH 3.0 and 5.0, the solution concentration of SPA decreased relatively slowly upon contact with SnO2, corresponding to the slow adsorption of SPA onto SnO2. At pH 5.0, the uptake after 2 hours was about 80% of the equilibrium amount and about 90% after 4 hours. Uptake continued however for at least 30-40 hours. At pH 3.0, the uptake after 2 hours was only about 50% of the equilibrium amount; after 4 hours this had increased to about 60%. The adsorption process at pH 3.0 required approximately 50-60 hours to reach completion. Similar tests were performed to determine the rate of adsorption of BAA onto SnO2 at pH 3.0 and pH 5.0. However at pH 5.0 there was evidence that the adsorption process is not straightforward. It was found that the interaction of BAA with Sn02, although giving the expected decrease in the solution concentration of BAA (as judged from the slight decrease in the relative peak heights in the BAA spectrum, see Fig. 2), also resulted in the release into solution of one or more UV active species as evident from the broad adsorption band developing with time below 380 nm (Fig. 2). These UV active species were presumably the products from the decomposition of BAA which was apparently catalysed by SnO2, since equivalent BAA solutions were found to be totally stable at pH 5.0 in the absence of SnO2 as were aqueous suspensions of Sn02. This phenomenon made determination of the true adsorption of BAA onto SnO2 at pH 5.0 uncertain and as such no data for the adsorption of BAA at this pH is reported here. The rate of adsorption of BAA onto SnO2 at pH 3.0 for an initial BAA concentration of 10 -3 mol dm -3 is shown in Fig. 3. The uptake after 2 hours was about 50% of the equilibrium amount; after 4 hours about 70%. Adsorption of BAA at pH 3.0 continued for approximately 30-40 hours before equilibrium was achieved. There was no evidence of the decomposition of BAA at pH 3.0 in the presence of SnO2. The adsorption isotherms of SPA at pH 3.0 and pH 5.0 and BAA at pH 3.0
10.2
9.8
o c
-~ E
)~o-o
o
o-
pH 5.0
9.4
g ,o pH 3.0
8.6 I
/ 40
I Times
I 80
I
I 120
(hrs)
Fig. 1. Rate of adsorption of SPA onto SnO2 at pH 3.0 and pH 5.0.
261 1.0
0.8
0.6 .a
0.4
0.2
0
I
i
I
i
i
I
240
1.0
280 Wavelength
320
360
400
(nm)
0.8
0.6
o
<
0.4
o2
0
I
240
I
280 Wavelength
320 (nm)
360
400
Fig. 2. UV spectra (230-300 nm) of BAA solutions after contact with Sn02 at pH 5.0. (a) 0 hours, (b) 2 hours, (c) 24 hours.
were determined after 100 hours contact with SnO2 and are given in Figs 4 and 5, respectively. The shapes of the SPA adsorption isotherms are similar at pH 3.0 and pH 5.0. As the pH increased from 3 to 5, the amount of SPA adsorbed at a constant residual solution concentration decreased. The amount of SPA adsorbed decreased less than the corresponding increase in the hydroxyl ion concentration, e.g. at 10 -3 mol dm -3 the amount adsorbed decreased ~ 7-9 times (2.45 to 0.33/~mol m -2) for a 100-fold increase in [ O H - ] (pH 3 to 5); with similar results at 5-10 -4 and 10 -4 mol dm -3 where the amount adsorbed also de-
262
to -6 E
o 5
8
'
~ 7 c_
~
'
4'0
'
8
'o
'
'
120
Time (hrs)
Fig. 3. Rate of adsorption of BAA onto Sn02 at pH 3.0.
/
4 E "5 E
3
10 H 30
E
6
o
c
8
E
4
pH 5.0
E
E 0
I -6
-5
-4
-3
I -2
log (Residual solution concentration (tool din-3))
-5
I -4
L -3
-2
log (Residual solution concentration (mol dm-3))
Fig. 4. (left) Adsorption isotherms for SPA after 100 hours contact with SnO2 at pH 3.0 and 5.0. Points shown by (m) are from data given in Fig. 1. Fig. 5. (right) Adsorption isotherm for BAA after 100 hours contact with Sn02 at pH 3.0. Point shown by (m) is from the data given in Fig. 3.
creased ~ 7 - 9 times (1.80 to 0.20 Hmol m -2 and 1.06 to 0.13 Hmol m -2) for the same increase in pH. The BAA adsorption isotherm at pH 3.0 was similar in shape to that of SPA at the same pH. However BAA adsorbed more strongly, the amount of uptake being approximately twice that of SPA for the same residual solution concentration and pH. For example at 10 -2 and 10 -4 mol dm -3, the uptake of BAA was 1.9 and 2.5 times that of SPA (4.70 to 2.45 and 2.70 to 1.06 Hmol m -2 respectively) at pH 3.0.
Effect of adsorption on zeta potentials The i.e.p, of the B D H SnO2 was determined to lie at pH 4.1 from the intersection of the ~ versus pH curves at 10 -2, 10 -3 and 10 -4 mol dm -3 KC1 where ~ = 0. The B D H Sn02 value compares well with the value o f p H 4.5 determined previously by Houchin and Warren [11 ] for SnO2.
263
.
2O
.
.
.
lO o s%
¢1 -t0
~u
-20
-so
--@ .... -~^=o=---~o---o~-;~ I
I
I
0
10-5
10-4
Organic acid concentration
50
I 10 -3 (mol dm - 3 )
Fig. 6. Zeta potential of S n 0 2 a s a function of organic acid concentration (10-8, 10-4 10-5, 0 tool dm -3) at pH 3.0 [SPA (@), BAA ( T ) , SSA (m), BSA (A), none ( ~ ) ] and at pH 5.0 [SPA ( O ), BAA ( V ), SSA ([]), BSA (/x ), none (~>) ]. 0.4
0.3 E
0.1 0 0
0
0
120
SPA adsorption time (hrs)
Fig. 7. Rate of release of hydroxyl ions during the adsorption of SPA (10-3 mol din-3) onto Sn02 at pH 5.0.
Figure 6 shows the effect of adding SPA, BAA, BSA and SSA to suspensions of B D H SnO2. Zeta potentials were measured after 100 hours contact between the organic acid (at 10 -3, 10 -4 and 10 -5 mol dm -3) and SnO2, in 10 -3 mol dm -3 KC1. At pH 5.0, the values in the absence and in the presence of organic acid concentrations up to 10 -3 mol dm -3 were not significantly different and were probably within experimental error ( + 3 mV). This conclusion is supported by the fact that there was no trend towards higher or lower zeta potentials as the acid concentration was increased. At pH 3.0, treatment of B D H SnO2 with SPA or BAA made the particles more negative, SPA changing the zeta potential from + 16 to - 1 0 mV and BAA changing the value from + 11 to +2.5 mV over the range 10 -5 to 10 -3 mol dm -3, compared to 21 mV in the absence of the organic acids. BSA and SSA had no effect (Fig. 6. )
Release of hydroxyl ions As SPA was adsorbed by SnO2, hydroxyl ions were released, as shown in Fig. 7 for the adsorption of 10 -3 mol dm -3 SPA onto SnO2 at pH 5.0. The rate of
264 release of hydroxyls slowed as the adsorption proceeded and was virtually zero after approximately 72 hours. The total amount of O H - released during this adsorption experiment was 0.36/lmol m -2 which can be compared to the SPA adsorption density at pH 5.0 (see Fig. 4) for a 10 -3 mol dm -~ SPA solution of 0.33/~mol dm -3. DISCUSSION Adsorption versus reaction at the surface
The observed decrease in the solution concentration of SPA after it is mixed with stannic oxide may result from one or more of the following processes: (i) adsorption of phosphonate anions onto the SnO2 surface without lattice rearrangement, (ii) reaction with "bulk SnO2" near the surface, probably after lattice rearrangement, to form a separate phase of tin (IV) phosphonate or a mixed tin (IV) oxide-phosphonate, (iii) a precipitation reaction between phosphonate anions and cationic species in solution, e.g., dissolved Sn. The last process, (iii), is improbable because of the very insoluble nature of stannic oxide at normal pH values [12]. However, process (ii), reaction with bulk SnO2, is a possibility since the uptake of SPA was slow requiring at least 30 hours to reach completion, whilst the accompanying release of hydroxyl ions took ~ 70 hours to reach completion. Simple adsorption processes (process (i)) are not normally so slow unless the adsorption involves a chemical interaction with a significant activation energy. The shape of the apparent absorption isotherm can be useful in distinguishing adsorption, process (i), from reaction, process (ii). Simple adsorption is often indicated by the isotherm starting off as Henry's Law and becoming less sensitive to concentration, often reaching a definite plateau as found by Edwards and Ewers [13] for uptake of sodium hexadecyl sulfate (SHDS) by stannic oxide at pH 2.6. Reaction with the bulk is often indicated by the isotherm increasing steadily with increasing concentration, well past the point of equivalent monolayer coverage, as found by Zimmels et al. [ 14 ] and Kitchener [15] for oleate on calcite or Wottgen [3] for HPA on cassiterite at pH 2. Our isotherms for SPA and BAA match neither of the above extreme cases. Total uptake of SPA was always less than that of a close-packed monolayer (assuming a molecule of SPA occupies an area of ~ 0.25 nm 2 on the SnO2 surface, then close-packed monolayer coverage would correspond to an adsorption density of approximately 6.5/~mol m -2). Another argument against reaction as an uptake process for SPA and BAA is that there are many examples of adsorption processes which produce isotherms of unusual shape a n d / o r without a saturation plateau, e.g. sodium dodecyl sulfonate on alumina [16], cobalt cations on silica [17], fluoride on alumina [18,19]. In some of these cases the adsorption mechanism changed as
265 the solution concentration increased, or as the pH changed, leading in turn to a change in isotherm shape. In summary, whilst the data on isotherm shape and magnitude do not allow us to decide unequivocally between adsorption and reaction with the bulk, we believe the evidence favours adsorption as the main uptake process. This is strongly supported by the IR spectra of Kuys and Roberts [4]. Reaction, if it occurs, is probably more likely at low pH and with BAA than with SPA, given the known reactivity of phenylarsonic acid with quadrivalent metal ions in acidic media [20 ].
Mechanism of adsorption of SPA At pH 5.0, the main SPA species in solution is the monoanion, RPO (OH) O -, and the surface of SnO2 is negatively charged ( - 2 8 mV, see Fig. 6), yet significant adsorption occurred against the electrostatic repulsion (Fig. 4). The mechanism therefore must involve a chemical interaction, with or without hydrophobic association. The chain length of SPA is only 5.5 equivalent CH2 groups which is shorter than the critical value of 6 equivalent CH2 groups required before hydrophobic association substantially enhances adsorption [21]. Thus, the adsorption mechanism is primarily a chemical interaction between SPA and the SnO2 surface. Further, since the structurally similar sulfonates SSA and BSA were not adsorbed at any pH value, the chemical interaction is specific for phosphonate and arsonate head groups. It was expected that at low pH, organic acid anions would be adsorbed into the electrical double layer around the positively charged SnO2 particles. If such electrostatic adsorption occurred at pH 3.0, it was not detected and must therefore be less than 0.01 Hmol m-2. Thus, electrostatic adsorption cannot account for the amount of SPA adsorbed which was of the order of i Hmol m -2. The key observation in explaining the nature of the adsorption process was the release of hydroxyl ions which accompanied the uptake of SPA, but quantification of the exchange reaction was difficult because the total amount of OH - released was small. At pH 5.0, for each RPO ( OH ) O - adsorbed, one OH (within experimental error) was released, suggesting that the adsorption process can be described by an equation of the type: -SnOH(s~ + RPO (OH) O~-~)~--SnO (OH) OPR(s) + OH ~-d~)
(1)
Equation (1) refers to an equilibrium between surface species (s) and exchanging anions in the double layer (dl) adjacent to the surface. Our adsorption isotherms for SPA at pH 3 and 5 (see Fig. 4) show that equilibrium for this ion exchange reaction lies well towards the side of the reactants, i.e. it is very difficult to exchange phosphonate ions for surface hydroxyl groups. At pH 5.0, zeta potential measurements of SnO2 before and after treatment with SPA were, within experimental error, the same irrespective of the con-
266 centration of SPA (Fig. 6). This is consistent with Eqn (1). At pH 3.0, the zeta potential of the Sn02 particles became increasingly negative as more SPA was adsorbed. However, the overall change was small (see Fig. 6), and this in a region where the zeta potential is very sensitive to small changes in surface charge density [11 ]. Neutralization of the positive surface charge, which arises from a small number of - S n O H + groups (the dominant surface species is always -SnOH; see Houchin and Warren [11 ] ) is consistent with reactions such as:
- S n O H + + RPO (OH) O- ~ - S n O (OH) OPR + H20
(2 )
- S n O H + + RPO (OH) 2 ~ - S n O (OH) OPR + H3 0 +
(3)
If we assume that Eqn (1) alone describes the adsorption of phosphonate and has a pH-independent equilibrium constant, then SPA adsorption should fall in proportion to the increase in [ O H - ] , as Edwards and Ewers [13] observed for SHDS on SnO2. We did not observe this behaviour for SPA or BAA nor did Wottgen [3 ] for HPA. Our results show that adsorption of phosphonic and arsonic acids by stannic oxide is not as sensitive to pH as expected from simple ion exchange, but the reasons for this are not clear. With diprotic acids such as SPA and BAA it is possible that the initial onepoint attachment to the SnO2 surface is followed by ring closure to form the surface analogue of a bidentate chelate, as shown below: 0
O-
HO
\1/
+
/X R
Sn
.o/ OH
I\
0 ~,,
\\/°\1/ P
Sn
/\
+.
OH-
+
H20
.o/I x
R
OH
p R
/°\ I/
/\/I
Sn
0
\
Similar structures have been proposed for the adsorption of phosphate onto goethite [22,23]. Our evidence does not allow us to distinguish between the bidentate bridging complex and the bidentate chelation of SPA to a single surface Sn atom, although recent F T I R spectra favour the latter [4 ]. Ring closure should be favoured by the fact that the second proton of SPA would become more acidic after SPA was adsorbed. Ring closure might still be a slow process and contribute to the overall slow rate of uptake of SPA.
Mechanism of adsorption of BAA The shape of the adsorption isotherm for BAA on stannic oxide at pH 3.0 is similar to that for SPA, although the magnitude of adsorption is about two
267 times greater. The differences in adsorption density may be related to the greater abundance of the neutral arsonic acid species, RAsO (OH)2, and/or a difference in the chemical reactivity of arsonic and phosphonic acids towards -SnOH. Presumably the major adsorption reaction for the arsonic acid at pH • 3.0 is Eqn (5) rather than ion exchange following Eqn (1): -SnOH + RAsO (OH) 2 ~--SnO (OH) OAsR + H20
(5)
Equation (5) is an acid-base reaction in which the stannic oxide surface acts as the base. The apparent catalytic effect of stannic oxide on the decomposition of BAA at pH 5.0 (Fig. 2 ) but not at pH 3.0 requires further investigation before the mechanism of adsorption of BAA onto SnO2 can be established more precisely. GENERALCOMMENTS The uptake of organic acid anions onto a negatively charged SnO2 surface, the dependence of uptake on the nature of the head group (sulfonates not adsorbed) and the release of about one O H - for each SPA anion adsorbed are all consistent with ion exchange chemisorption, Eqn ( 1 ), even though Eqn (1) is not obeyed quantitatively. A number of other systems show exchange chemisorption. Phosphate anions chemisorb on iron oxide, by exchange with surface hydroxyl groups, to produce a bidentate bridging complex like that proposed for the phosphonate in Eqn (4) [22]. Sulfate anions and alkyl sulfates chemisorb on iron oxides by ion exchange with surface OH [24,25] showing that exchange chemisorption occurs even when an organic tail is attached to the head group. The long-chained SHDS adsorbed by ion exchange with surface OH on stannic oxide [13], although hydrophobic association between the organic tails presumably enhanced the adsorption. Marinakis and Kelsall [26] concluded that dodecyl sulphate and decyl phosphonate ions adsorb via anion exchange with tungstate anions from scheelite (CAW04). These examples show that exchange chemisorption of anionic surfactants on oxides and salt-type minerals may be of more general importance than previously believed. The question remains as to why the sulphonates, SSA and BSA, do not chemisorb onto SnO2 whereas phosphonates, arsonates and sulphates do. Two factors may be involved: (i) sulphonates are weaker ligands for metal ions in aqueous solutions than phosphonates, arsonates and sulphates [27], and (ii) ring closure or bidentate chelation to a single surface Sn atom is not possible with monoprotic sulphonic acids.
268 CONCLUSIONS (1) T h e main process by which suspensions of stannic oxide abstracted styryl phosphonic acid from aqueous solutions at p H 5 was t h a t of exchange chemisorption between p h o s p h o n a t e m o n o a n i o n s and hydroxyl groups a t t a c h e d to tin atoms in the oxide surface. (2) For each p h o s p h o n a t e anion adsorbed at p H 5, approximately one hydroxyl ion was released. However, not all results were consistent with a simple exchange process, e.g. changing the p H did not cause a proport i onat e change in the a m o u n t of SPA adsorbed. (3) T h e a m o u n t of SPA adsorption was relatively low { ~ 20% monolayer) reflecting the preference of surface tin atoms for hydroxyl over phosphonate. Also, adsorption was slow, requiring at least 20 hours. (4) At p H 3.0, the uptake of benzene arsonic acid by stannic oxide was approximately twice t h a t of SPA but the m e c ha ni sm of adsorption was not established. (5) T h e short-chained sulfonic acids SSA and BSA were not adsorbed to a significant e x t e n t ( < 0.01/~mol m -2) consistent with their lower complexing ability with metal ions compared with phosphonates, arsonates and sulfates. (6) Exchange chemisorption of anionic surfactants with oxides and salttype minerals may be of more general importance t h a n previously believed.
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