1196
Ind. Eng. Chem. Res. 1995,34, 1196-1202
The Use of Hydrotalcite as an Anion Absorbent Linda M. Parker,* Neil B. Milestone, and Roger H. Newman New Zealand Institute for Industrial Research and Development, P.O. Box 31-310, Lower Hutt, New Zealand The high affinity of Mg/Al hydrotalcite for carbonate anions means that it cannot be reversibly exchanged and so prevents its use as a n anion-exchange material. However, carbonate can be removed by thermal decomposition, evolving C02. The resultant material sorbs anions when placed in solution and returns to the hydrotalcite structure. This work showed that hydrotalcite can be used as a n anion absorbent by firing, anion sorption, exchanging in carbonate, and repeating the cycle. Ion exchange, XRD, TGA/MS, and 27AlNMR results showed that after the first firing and rehydration cycle, the hydrotalcite structure becomes disordered, with its sorption capacity reduced by -50%. Further firinghorption cycles cause only a slow degradation in structure.
Introduction Hydrotalcite (Frondel, 1941)is a mineral that displays anion-exchange properties, in contrast to the more common cation-exchange properties of clays. It is composed of octahedral Mg(OH12 layers where a maximum of about one in three trivalent aluminum sites are substituted by divalent magnesium (Brown and Gastuche, 1967). This substitution results in positively charged layers separated by charge-balancing anions and water. Although hydrotalcite shows anion-exchange properties, it cannot readily be used for this purpose. The carbonate is preferentially sorbed and prevents further significant anion exchange (Sat0 et al., 1986). Only hydrotalcite which has been carefully prepared in carbonate-free solutions shows an anionexchange capacity of -3 mequivlg (Miyata, 19831, similar to that obtained for organic ion-exchangeresins. Hydrotalcite shows the greatest affinity for anions with the highest charge density, with multicharged anions sorbed more strongly than monovalent anions (Dutta and Puri, 1989; Miyata, 1983). This is because the ionexchange reaction occurs with bare, unhydrated anions. Organic anion-exchange resins induce an ion-exchange reaction with the hydrated anion and so have a reversed order of affinities. Although the carbonate anion is almost impossible to remove by ion exchange, it can be removed by heating the hydrotalcite above 400 "C, decomposing C032- to C02 and 02-,leaving 02-anions between the layers (Miyata, 1980;Reichle et al., 1986; Sat0 et al., 1986 and Miyata, 1983). The shape of the crystals remains unchanged although numerous fine pores are observed on the surfaces as a result of the degassing (Reichle et al., 1986). However, as shown by XRD, the crystalline hydrotalcite structure has collapsed to give an amorphous magnesium oxide with aluminum ions dispersed as a solid solution. Some of the aluminum has changed from octahedral to tetrahedral coordination, along with some migration of aluminum to the surface of the crystals (Reichle et al., 1986). Some short-range order, not detected by X-rays must still exist in this solid solution since addition of water restores the original hydrotalcite structure, allowing sorption of anions from solution. When hydrotalcite is heated above 900 "C, exothermic peaks are observed by differential thermal analysis and some MgA1204 appears in the XRD pattern. These changes are irreversible as addition of water does not re-form hydrotalcite. The ability of hydrotalcite heated between 400 and 900 "C to re-adsorb water and anions as it returns to 0888-5885/95/2634-1196$09.00/0
its original structure means that it can be used as an anion sorbent. This property has been investigated by Sat0 et al. (1986) who showed that when heat treated hydrotalcite was added in excess to an anion solution, the HP042-, Si032-, sod2-and Cr042- anions were almost completely removed. Selectivities were higher for divalent anions than monovalent anions, as is observed for the parent hydrotalcite. In this work we have investigated the possibility of exploiting the anion sorption capacity of fired Mg/Al hydrotalcite by reusing it in a cycle involving heat treatment, anion sorption, and carbonate exchange to remove the sorbed anions, and then repeating the cycle. XRD, TGA/MS, and 27AlNMR techniques were applied to study the processes occurring.
Experimental Section Materials Studied. A sample of commercial, synthetic hydrotalcite (DHT4) was obtained from the Kyowa Chemical Industry Co. Ltd., Japan. Its molecular formula was given as Mgd2(0H)16*C03*4H20. Chemical analysis by atomic absorption spectroscopy after alkali fusion gave 21.2 wt % Mg and 10.2 wt % Al. The mole ratio Mg/Al was 2.4, slightly lower than the ratio of 3 give in the formula above. From these results the molecular formula is Mg5.65A12,35(OH)16'1.2co3. 6.5H20. Synthesis of Hydrotalcite Samples without Carbonate. To exclude carbonate ions, all solutions were prepared with CO2-free distilled water, and the syntheses were prepared in a nitrogen atmosphere in a glovebag. Aluminum salt (0.01 mol) and magnesium salt (0.02 mol) were dissolved in 15 mL of water in a stainless steel bomb. Alkali solution (25 mL of 25% NaOH) was added slowly with stirring. The bomb was sealed, rotated, and heated at 80 "C for 2 days. For prep A, aluminum sulfate and magnesium sulfate salts were used. For prep B, the chloride salts were used. Both syntheses produced crystalline hydrotalcite, as shown by sharp peaks in their XRD pattern. Determination of Anion-Exchange Properties. A hydrotalcite sample (-0.5 g) was accurately weighed and added to 25 mL of a multianion solution containing 2.3 mmoYL NazB40~10Hz0,2.2mmoYL NaF, 1.3 m m o K NaBr, 1.7 mmol/L m03,and 1.82 mmoVL KzHPO4 with a pH of 9.16. The samples were shaken in the solution for 1 h and then left for 18 h. The resultant solution was then filtered. The concentrations of F-, C1-, Br-, NOS-, HP042-, and s04'- anions remaining in solution 0 1995 American Chemical Society
Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 1197 Table 1. Results for Sorption in Multianion Solutiona ~
sample DHT4 as received DHT4 C1- exchanged DHT4 sod2-exchanged DHT4 1st cycle (FWb DHT4 2nd cycle (FSFSb DHT4 4th cycle prep A as synthesized prep A 1st cycle prep A 2nd cycle prep B as synthesized
F-
Br-
.3 1.8 .4 10.0 5.6 8.8 5.4 8.2 7.7 10.8
-.2
~~
(AmmoVg sorbantlx 100 B(OH)4NO3-
.2 .2 5.3 2.1 1.3 1.5 0.9 1.2 1.2
-1.4 -5.7
-.7
-1.1
-.7 5.1
.8
42.5 24.9 12.0 6.1 6.5 7.5 25.3
-.4
0 5.6 7.1* -1.1 0
HP04’2 1.5 2.0 8.3* 8.7* 8.4 8.3 4.2 4.7 10.2*
Zmequivlg (excluding C032- and OH-) 0.02 0.09 0.06 0.85 0.5 0.39 0.39 0.38 0.26 0.56
a An asterisk notes that the concentration of the anion was reduced to below the level of detection, therefore maximum sorption capacity was not reached. F = fired, S = anion sorption, C = Carbonate exchanged.
were determined by the Dionex method. Borate was analyzed by atomic adsorption spectroscopy. Carbonate and hydroxide anions could not be analyzed by AA or Dionex and were not determined in this survey.
Anion Sorption Capacities for Single Anions. A sample of freshly fired (400 “Cfor ’2 h) DHT4 (1.00 g) was added to 100 mL of decarbonated anion solution. The solution concentrations used were 1.07 mmol/L for nitrate; 49.9 mmol/L for borate; 44.4 mmol/L for phosphate; 51 mmol/L for fluoride, and 50 mmol/L for both sulfate and chloride. The solutions were continuously bubbled with nitrogen to provide mixing. Solution samples (2 mL) were withdrawn, filtered, and then analyzed using Dionex anion-exchange chromatography for fluoride, bromide, and nitrate, and phosphate and by AA spectroscopy for borate. XRD. The X-ray powder diffraction patterns were obtained using a Philips PW 1700 APD system with a diffracted beam monochromator and Co Ka radiation.
Thermogravimetric Analysis/Mass Spectrometry (TGNMS). Thermogravimetric analysis was carried out with about 10 mg of sample heated in an argon gas flow (15 mumin). Effluent gas was sampled via a capillary inlet into a mass spectrometer and continuously analyzed. The system is described in more detail in earlier reports (Parker and Patterson, 1983). Nuclear Magnetic Resonance (NMR). Between 0.1 and 0.4 g of powder was packed in a 7 mm diameter cylindrical silicon nitride rotor with Vespel end caps. Rotors were spun a t 7 kHz in a magic angle spinning (MAS) probe manufactured by Doty Scientific Inc. for aluminum-27 NMR at 52.1 MHz in a Varian XL-200 spectrometer. Each 0.1 ,us radio frequency pulse was followed by 20 ms of data acquisition time and a delay of 0.5 s. Transients from 4096 pulses were averaged. The chemical-shift scale was referenced relative to the signal for aluminum-27 in aqueous aluminum nitrate.
Results and Discussion Anion Exchange. The anion-exchange properties of the hydrotalcite-like materials were initially studied by observing the amounts of different anions sorbed from a multianion solution containing added fluoride, bromide, borate, nitrate, and phosphate. An indication of the relative affinities and exchange capacities for each anion, in competition with the other anions, was obtained. Carbonate ion was not deliberately added, but carbon dioxide was readily sorbed from the atmosphere into the alkaline solution as carbonate ions. Hydroxide anions were also present. The results are presented in Table 1. In some cases, the concentration of an anion was reduced to below the
level of detection. These are shown by an asterisk in the table and indicate that the maximum sorption capacity was not reached for that anion. A negative number shows an increase in concentration of the anion in the anion solution. For hydrotalcite DHT4 in its “as received” carbonate form, only a small amount of anion exchange occurred for the anions studied. Only 0.02 mequiv/g of anions was exchanged compared to -3 mequivlg shown possible by Miyata (1983) and a theoretical maximum of 3.6 mequiv/g if all the carbonate in the molecular formula was exchanged. Pretreatment in 1 M Na2S04 or 1 M NaCl provided only a slight increase in ion-exchange capacity. For some materials, such as zeolites, increasing the ion-exchange temperature increases the degree of ion exchange (Ward, 1984). However, while severe conditions such as boiling in 1 M Na2S04 (pH 3.2) dissolved some of the hydrotalcite, there was no significant decrease in the amount of C02 desorbed as observed by TGA/MS, showing that the hydrotalcite had remained in its carbonate form. Therefore the removal of carbonate by simple ion exchange is not practical. Careful synthesis of Mg/M hydrotalcite under “carbonate-free” conditions provided a material with a greater ion-exchange capacity. For hydrotalcite synthesized with sulfate anions (prep A), 0.38 mequiv/g anions (excluding carbonate) was exchanged. For hydrotalcite synthesized with chloride anions (prep B), 0.56 mequiv/g anions (excluding carbonate) were exchanged. TGA/MS results (discussed in more detail later) showed that C 0 2 was evolved when these exchanged hydrotalcites were heated. This meant that carbonate anions had also been absorbed from the multianion solution, reducing the anion-exchange capacity measured excluding carbonate. Therefore, even when the hydrotalcite is synthesized in a carbonate-free environment, anion exchange must be carried out in carbonate-free solutions to obtain maximum anion sorption capacity. Any carbonate ion present in the exchange solutions is sorbed preferentially and irreversibly, preventing further exchange. An alternative method used by Sat0 et al. (1986) is to remove the carbonate ion by heat treatment. The fired hydrotalcite can be used to sorb anions, returning to its original hydrotalcite structure. In this work, we have investigated carrying out further exchange in a carbonate solution to replace sorbed anions with carbonate so that the hydrotalcite can be refired and reused as a sorbent. The changes in sorption capacities for DHT4 and prep A after repeated cycles are shown in Table 1. For DHT4, the phosphate ion was completely removed for
1198 Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 I
2,
Table 2. TGA/MS Data for the M g M Hydrotalcite (DHT4) after Different Treatments" mlz17 mlz44 o/" weight loss integralb integralb 30-220 220-700 50-700 250-700 sample treatment "C "C "C "C 1120 100 as received 13 27 boiled in 0.5 M NaZS04 14 27 1400 97 1075 45 FS' 21 22 FSCb,' 23 22 18 FSCFSb>' 23 20 1140 FSCFSCbac 16 25 845 45 F 380 "C 24 h 5 4 191 0 F 2 days later 6 10 449 40 F 43 days later 11 20 1240 92 "The weight losses are given as a percentage of the initial sample weight. The integrals of the water (m/z 17) and COz (mlz 44) desorption peaks have been scaled to constant sample weight and mass spectrometer sensitivity. Arbitrary units. F = fired, S = anion sorption, C = carbonate exchanged.
"-
2
0
4
6
8 10 12 14 16 18 20 22 24 Time (h)
Figure 1. Amount of anion sorbed by freshly fired hydrotalcite versus time. The experiments for each anion were carried out separately in decarbonated solutions.
the first cycles. A similar amount of fluoride ion was removed each cycle, but nitrate ion was removed on the first cycle only. Bromide and borate ions were removed in decreasing amounts. For prep A, the total amount of anions sorbed was the same for both the as synthesized and fired (first cycle) samples. However the amount of phosphate sorbed was reduced and fluoride and nitrate increased. After the second cycle no nitrate was sorbed (a small amount desorbed), with sorption of the other anions remaining almost constant. Further work is required t o understand these changes in affinities with cycle. These results mean that, for the conditions and anions studied here, the use of recycled hydrotalcite is limited to sorption of phosphate, fluoride, and perhaps borate. Anion Sorption Capacities for "Single"Anions. Variations in the anion concentration with time were measured for freshly fired DHT4 in carbonate-free solutions containing a single added anion. The results (Figure 1) clearly show the differences in sorption capacities expressed as milliequivalents of anions sorbed per gram of original hydrotalcite. Initial rapid sorption was observed for all anions (except chloride), with phosphate and borate showing an oscillation. This was followed by a slower increase to the final value observed. After 24 h, the total amounts sorbed were in the order SO -:
>
F-
>
HP0;-
>
C1- > B(OH),- > NO,-
This order is similar to the order of ion exchange equilibrium constants given by Miyata (1983) as
OH-
>
C0,2-
F>
>
SO-;
C1- > Br- > NO,- =- Ifor monovalent anions for divalent anions
The more strongly bound anions are those with the smaller ionic radii which result in a decreased spacing between the crystal hydroxide layers.
The maximum theoretical anion capacity of 3.6 mequiv/g was not achieved. This is because hydroxide anions would have also been present and competing with sorption of the anion studied. Sat0 et ai. (1986) observed the pH of solutions increased to 9.5 soon after addition of fired hydrotalcite. Powder X-ray Diffraction (XRD). XRD of the "as received" hydrotalcite DHT4 (Figure 2A) shows very sharp peaks, indicating a very crystalline, ordered material. After firing and rehydration in an anion solution, the hydrotalcite structure returns but the peaks are broadened (Figure 2B), showing some disorder remains in the crystal structure. After cycling four times (Figure 2C), the peaks continue to broaden and reduce in intensity, showing a further loss of crystallinity. The fired material slowly sorbs carbon dioxide and water from the atmosphere, reverting to hydrotalcite. A sample exposed on the laboratory bench exhibits broad peaks similar to brucite (Mg(0H)Z)several hours after firing (Figure 3A). One day after firing (Figure 3B), the XRD remains unchanged except for the appearance of an extra peak at 2.35 A. After 7 days of ambient exposure (Figure 3C), peaks associated with hydrotalcite were observed, although a large proportion of the fired material remained. The peak at 2.35 A was replaced by a hydrotalcite peak at 2.29 After 43 days (Figure 3D), peaks from hydrotalcite dominate with a smaller amount of the fired material still present.
A.
Thermogravimetric Analysis/Mass Spectrometry (TGA/MS). TGA/MS of the as received hydrotalcite DHT4 is shown in Figure 4. The weight losses and integrals of the C02 and water desorption peaks are given in Table 2. The first weight loss from 50 to 230 "C is due to physisorbed water. The XRD results of Brindley and Kikkawa (1979) showed that only a small change in the interlayer spacing from 7.8 to 7.9 A occurs. From 260 to 370 "C, further interlayer water is lost with an accompanying decrease in the layer spacing to 6.4 A. From 375 to 650 "C, both water and COz are lost from between the layers, with the layer spacing approaching that of brucite (Mg(0H)z). When a sample was boiled for 10 min in -0.5 M NaaS04, no significant changes were observed in the weight losses or shapes of the water and carbon dioxide desorption peaks. Therefore, no significant sulfate exchange had occurred for the fully carbonated material. All the carbonate is lost by firing in air at 380 "C for 24 h, as shown by TGA/MS (Figure 5A). The small
Ind. Eng. Chem. Res., Vol. 34,No. 4, 1995 1199 x 103 1.50
1.35
A
1.20 1.05 0.90 0.75
0.60 0.45
g3 -
0’30 0.15
0.0
0.3 40 5
10.0
1
20.0
30.0
40.0
.50.0
60.0
70.0
80.0
B
0.15-
0.45’
C
0.30’ 0.15-
0.0
10.0
20.0
30.0
40.0
50.0
60 0
70 0
80 0
wFigure 2. Powder X-ray difiaction patterns of hydrotalcite showing changes resulting from firing and rehydration. (A) The “as received” hydrotalcite. (B)The hydrotalcite after firing at 400 “C and rehydration. (C) The hydrotalcite after four firinghehydration cycles.
weight losses from 12 to 300 “C and 450 to 700 “C are due to water desorption from MgO. One day aRer firing (Figure 5B), a small COS desorption peak with a maximum a t -300 “C showed that some COShad been sorbed from the atmosphere. This desorption temperature is lower than that observed for hydrotalcite and may be due to decomposition of an intermediate material formed enroute to hydrotalcite. This intermediate material is also appears to be associated with an additional XRD peak (Figure 3B). After 43 days (Figure 5 0 , the COS desorption peak has shifted to the same temperature range as that observed for the “as received” hydrotalcite, although a different peak shape is observed. The temperature of the peak maximum is at the same temperature as that observed by us for the decomposition of magnesium carbonate (MgCO3-Mg(OH)z*5Hz0)under the same conditions. After the initial anion sorption in the first cycle, the full capacity of the hydrotalcite is never restored. This is also shown in both the percentage weight losses and in the size of the evolved COS integrals given in Table 2. After the first firing, the high-temperature weight
loss decreases from 27 to about 22%. The evolved COS decreased from 100 to 45 (arbitrary units). TGA/MS of the samples after sorption in the multianion solution show evolution of some COS,demonstrating that carbonate had been sorbed along with the other anions. For the FS sample, TGA/MS (Table 2) showed the amount of COS desorbed was almost half that desorbed from the original sample, implying that almost half the anion sorption capacity had been consumed by carbonate. A comparison of the amount of carbon dioxide desorbed from the sample after exchange in anion solution (FSCFS, Figure 6) and after exchange in carbonate solution (FSCFSC, Figure 7) showed that 40% of the anion sorption capacity had been consumed by carbonate sorbed from the anion solution. A comparison of the water desorption integrals shows that less low-temperature physisorbed water is desorbed for the fully carbonated sample than the sample containing a mixture of anions (compare Figure 7 with Figure 6). This may be due to an optimal packing arrangement of the carbonate anions between the layers resulting in less additional water present than with
1200 Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995
1
0.4 50
A
0.30 0.20-
0.10-
0 0
io0
200
300
500
4 0 0
600
700
eo0
0 401
{
0.30 0.20
i
0 401
0.30 0.20.
0.10-
J
.
,
,
.
.
.
,
,
,
,
,
0.0
10.0
20.0
30.0
40.0
50.0
60.0
70.0
eo o
o o
10.0
20.0
30.0
40.0
50.0
60.0
70.0
80.0
oegFigure 3. Powder X-ray diffraction patterns showing the effect of ambient exposure to hydrotalcite fired a t 400 "C for 2 h. (A) 8 h after firing, (B)24 h after firing, (C) 7 days after firing, and (D)43 days later.
other anions and would help account for the high affinity for carbonate anions compared to any other anion. TGA/MS of the carbonate exchanged sample d e r two firing cycles (Figure 7) showed a more complex CO2 desorption over a wider temperature range than the original hydrotalcite (Figure 4). This implies a variety of sites for the carbonate anions. The weight loss data also shows a reduced anion sorption capacity. NMR Results. The 27Al solid state NMR spectrum of hydrotalcite DHT4 (Figure 8) shows a peak with a single quadrupolar coupling constant arising from wellordered, octahedral aluminum. A similar line shape can be simulated with parameters for a single site. The simulation was based on a trial-and-error selection of a nuclear quadrupolar coupling constant x = 1.4 MHz, asymmetry parameter 11 = 0.1 and true chemical shift of 6 = 10 ppm. The broken line in Figure 8 was then generated by convoluting the simulated curve with a Gaussian broadening functions of width W = 0.14 kHz. Mg and Al atoms in a mole ratio of 3:l can be arranged
in a regular pattern in the octahedral layer so that the
AI atoms are always separated by Mg atoms (Brindley and Kikkawa, 1979). Elemental analysis indicated that the mole ratio for DHT4 fell below the nominal value of 3:l (see the Experimental Section). The discrepancy between simulated and experimental curves in Figure 8 might have resulted from additional signals from Al atoms that could not be accommodated in a regular 3:l structure. A signal from tetrahedral Al was observed after the sample was fired (Figure 9A). The fraction of Al involved is difficult to estimate because of overlap between bands, but 18%of the signal area was found on the low-field side of a vertical line drawn arbitrarily at 6 = 25 ppm. The dominant signal assigned to octahedral Al shows at tail spreading as far as 6 = -100 ppm. Such low-frequency (high-field) tails have been observed in aluminum-27 MAS NMR spectra of clay minerals (Woessner, 1989) and are attributed to distributions of values of x associated with disordered structures. Meinhold et al. (1992)have described a procedure
Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 1201
'"1
A
6-
. ,.-. -$I
Temperature
( O C
. ..-.-.-
~
100
200
300 400 500 Temperature (OC 1
600
7t)O
1bO
2do
3d0 400 500 Temperature PC)
6dO
?A0
)
Figure 4. TGAMS results for the as received hydrotalcite. The weight loss (-) and ion signals from desorbed water (- - -1 and carbon dioxide (- -) are plotted versus sample temperature.
for computer simulation of such tails. They introduced a parameter defined by
e
6 = ea,Q/h where a, is the standard deviation of random perturbations superimposed on electric field gradients. The ratio $11 is then a convenient measure of the effects of disorder on the NMR spectrum. The broken line in figure 9A was simulated with 61%= 2.0, W = 0.70 kHz, 6 = 13 ppm, and x and r] as for Figure 8. The best fit ratio of $ 1 was ~ so large that variation of x and r] had little effect on line shapes. Heating the sample seems to have caused movement of Al atoms. Some have acquired tetrahedral coordination, while others have formed a disordered arrangement within the lattice. Spectra of fired hydrotalcite recorded on a 500 MHz NMR machine with faster spinning speeds and better resolution show no evidence of other peaks between those observed for octahedral and tetrahedral aluminum (McKenzie, 1992). Reformed hydrotalcite showed no tetrahedral Al (Figure 9B). Chemical analysis showed that the MgIAl mole ratio remained constant over four firingfexchange cycles, implying that the tetrahedral Al returned to octahedral coordination in the hydroxide lattice. The signal assigned to octahedral Al retained a tail, but not as pronounced as that observed for freshly-fired material. The broken line in Figure 9b was simulated for $11 = 1.5, 6 = 13 ppm, and W = 0.53 kHz. The observation of a large ratio 611 suggests that the Al and Mg atoms remain highly disordered within the hydrotalcite layers rather than returning to the regular pattern associated with the original material. Aluminum-27 MAS NMR spectroscopy has been used in an earlier study of Zn-Al hydrotalcite-like material (Thevenot et al., 1989). The published spectra show asymmetric line shapes which we attribute to Zn-Al disorder analogous to the Mg-A1 disorder discussed above.
Conclusions Hydrotalcite has a high selectivity for carbonate anions, making it ineffective as an anion-exchange material unless further treatment is made. Carbonate anions are not even exchanged in boiling 0.5 M NaZS04
L
10
-
-9 F
v
-
e .C
Pa %
9-
ul
6-
-5L
0
I
100
I
200
I
300
I
400
Temperature
1
500
I
600
1
700
('C 1
Figure 5. (A) TGAMS results for hydrotalcite DHT4 immediately after firing at 380 "C for 24 h. (B)TGAMS results for hydrotalcite one day after firing at 380 "C for 24 h. (C) TGAMS results for hydrotalcite 43 days after firing at 380 "C for 24 h. The weight loss (-) and ion signals from desorbed water (- - -) and carbon dioxide (- * -1 are plotted versus sample temperature.
solution a t pH 3.5. However, heating to 300 "C causes decarboxylation as the carbonate anion decomposes, resulting in an amorphous material that will sorb anions and return to its original hydrotalcite structure. The sorbed anions can be exchanged with carbonate which can then be removed by firing, regenerating the
1202 Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995
e
n
j
J._____ T ___.. ~
100
L A ; , L7 -100
0 6 1PPm
100
-100
0 6 / ppm
Figure 9. Z 7 A l MAS NMR spectra of (a) fired hydrotalcite and (b) hydrotalcite after firing and rehydration. A signal assigned to tetrahedral sites is labeled “T”. Computer-simulated line shapes are shown as broken lines. (“C)
Temperature
Figure 6. TGA/MS results for hydrotalcite after two firing cycles and sorption in a multianion solution (FSCFS). The weight loss (-1 and ion signals from desorbed water (- - -1 and carbon dioxide (- -1 are plotted versus sample temperature.
-
Literature Cited
(‘C)
Temperature
Figure 7. TGA/MS results of the carbonate exchanged hydrotalcite after two firing cycles (FSCFSC). The weight loss (-) and ion signals from desorbed water (- - -) and carbon dioxide (- * -) are plotted versus sample temperature.
n --.
20
10
0
The XRD, TGA/MS, and NMR results show that although the hydrotalcite structure returns after firing and rehydration, the hydrotalcite is irreversibly altered. This could account for the decrease in sorption capacity. NMR observations show the disorder is due to the rearrangement of Mg and Al atoms within the octahedral layers. This is confirmed by XRD where an amorphous material results, showing that the longrange order of the crystal structure has been broken. Further firinghorption cycles cause only a slow degradation in structure. The fired material slowly sorbs carbon dioxide and water from the atmosphere over a period of days, reverting to hydrotalcite and losing its anion sorption capacity. Hydrotalcite can be used as an anion sorbant if it is fired to remove the carbonate anions where its higher affinity for multiply charged anions may be useful in some applications. A possible application may be for preconcentration of anions before analysis in a manner similar to that described by Andersson and Ingri (1991).
-10
6 1 PPm Figure 8. 27AlMAS NMR spectrum of hydrotalcite (solid line) and a computer-simulated line shape (broken line).
anion sorbent material. This sorption, carbonate exchange, and firing cycle can be repeated. After the first cycle the sorption capacity is reduced by -50%, but then shows only a slight decline with further cycles.
Andersson, P.; Ingri, J. Water. Res. 1991,25,617. Brindley, G. W.; Kikkawa, S. Am. Mineral. 1979,64,836. Brown, G.; Gastuche, M. C. Clay Miner. 1967,7 , 193. Dutta, P. K.; Puri, M. J . Phys. Chem. 1989,93, 367. Frondel, C. Am. Mineral. 1941,26,295. McKenzie, K. J. D., New Zealand Institute for Industrial Research and Development, Personal communication, 1992. Meinhold, R. H.; Slade, R. C. T.; Newman, R. H. Appl. Magn. Reson. 1992,in press. Miyata, S. Clays Clay Miner. 1980,28, 50. Miyata, S.Clays Clay Miner. 1983,31, 305. Parker, L.M.; Patterson, J. E. Chemistry Division Report No. CD 2330, 1983, DSIR, Private Bag, Petone, New Zealand. Reichle, W. T.; Kang, S. Y.; Everhardt, D. S. J. Catal. 1986,101, 352. Sato, T.; Wakabayashi, T.; Simada, M. Ind. Eng. Chem. Prod. Res. Dev. 1986,25,89. Serna, C. J.;Rendon, J. L.; Iglesias, J. E. Clays Clay Miner. 1982, 30, 180. Thevenot, F.; Szymanski, R.; Chaumette, P. Clays Clay Miner. 1989,37, 369. Ward, J. W. Appl. Ind. Catal. 1984,3 , 287. Woessner, D. E. Am. Mineral. 1989,74, 203. Received for review December 5, 1994 Accepted December 13, 1994 @
IE940718L
* Abstract published in Advance ACS Abstracts, February 15, 1995.
