A critical assessment of the methods for intercalating

A critical assessment of the methods for intercalating
A critical assessment of the methods for intercalating
anionic surfactants in layered double hydroxides
Lumbidzani Moyo1, Nontete Nhlapo1 and Walter W. Focke2
(1) Department of Chemistry, Institute of Applied Materials, University of Pretoria,
Lynnwood Road, Pretoria, 0002, South Africa
(2) Department of Chemical Engineering, University of Pretoria, Lynnwood Road,
Pretoria, 0002, South Africa
Walter W. Focke
Email: [email protected]
Abstract
Anionic surfactant intercalated layered double hydroxides (LDH) of high purity are easily
prepared via direct coprecipitation and also by the ion exchange method provided that the
precursor contains a monovalent anion, e.g., LDH–Cl or LDH–NO3. However, LDH–
CO3 is an attractive starting material as it is commercially available in bulk form owing
to large-scale applications as a PVC stabilizer and acid scavenger in polyolefins. Thus,
intercalation of dodecyl sulfate and dodecylbenzenesulfonate into a commercial (LDH)
with approximate composition [Mg0.654Al0.346(OH)2](CO3)0.173 · 0.5H2O] was explored.
Direct ion exchange is difficult as the carbonate is held tenaciously. In the regeneration
method it is removed by thermal treatment and the surfactant form obtained by reaction
with the layered double hydroxide that forms in aqueous medium. Unfortunately the
resulting products are impure, poorly crystallized and only partial intercalation is
achieved. Better results were obtained using water-soluble organic acids, e.g., acetic,
butyric, or hexanoic acid, to aid decarbonation of LDH–CO3. Intercalation proceeded at
ambient temperatures with the precursor powder suspended in an aqueous dispersion of
the anionic surfactant. The carboxylic acids are believed to assist intercalation by
facilitating the elimination of the carbonate ions present in the anionic clay galleries.
Introduction
Layered double hydroxides
The Mg–Al layered double hydroxides (LDH) are anionic clays with the general formula
where A is a charge balancing anion and x is the fractional aluminum substitution in the
layers. It usually varies from 0.20 to 0.36 [1–6]. LDH are synthetic analogues of the
natural mineral hydrotalcite [Mg6Al2(OH)16CO3 · 4H2O]. These materials feature a
brucite-like [Mg(OH)2] stacked sheet structure in which the cations are octahedrally
coordinated by six oxygen atoms as hydroxides. A net positive charge of the sheets arises
from the partial replacement of Mg2+ with Al3+ ions. The interlayer contains water and
charge balancing anions, e.g., carbonate ions in LDH–CO3 [7]. The three dimensional
structure of the clay is maintained by a combination of electrostatic forces and hydrogen
bonding interactions between the layer and interlayer anions or molecules [6].
Hydrophobic interactions also play a role when the inserted anions contain long aliphatic
chains [8–10]. The basal spacing of the (003) planes in brucite is 0.47 nm [11]. In LDH–
CO3 structures it increases slightly with a decrease in the fractional aluminum
substitution. Bellotto et al. [12] give values of 0.76 and 0.79 nm for x = 0.36 and x = 0.17,
respectively.
Surfactants and surfactant aggregates
Surface active agents (surfactants) are molecules with an amphiphilic nature [13]. Their
chemical structure contains two parts with very different polarities: a polar hydrophilic
head group (e.g., a sulfonate ion) and a non-polar hydrophobic tail. The latter is often
represented by a linear alkane chain segment. The hydrophilic head group shows strong
affinity for water but is not very compatible with non-polar solvents. The hydrophobic
segment is soluble in non-polar solvents but has a low affinity for water and other polar
solvents. Surfactants usually have a relatively low solubility in water and show a distinct
preference to adsorb on available surfaces and interfaces. This leads to reduced surface
and interfacial tensions, respectively.
Surfactant molecules self-organize into micellar aggregates above the critical micelle
concentration or CMC. Just above the CMC the micelles usually have a spherical shape
with surfactant molecules arranged such that the hydrophilic part is on the outside and the
hydrophobic part in the center. As the concentration is increased, the micelles coalesce to
first form elongated ‘worm-like’ tubes and later convert into lamellar sheets of organized
molecules [14].
Surfactant adsorption on clay surfaces
Clay particles present two different surfaces for interaction with surfactants. It is usual to
use the expression “adsorption” when the interaction is limited to aggregation of the
surfactant on the outside surfaces of the clay particles. The phrase “intercalation” refers
to situation where the surfactant molecules additionally aggregate inside the galleries,
i.e., between pairs of adjacent clay sheets. Crepaldi et al. [15] gives a short overview of
surfactant intercalation into LDH.
The adsorption of surfactants is a type of aggregate formation on the mineral surface.
Harwell et al. [16] refer to these as admicelles to emphasize the micelle-like aspects of
their structure and behavior. Bitting and Harwell [17] found that the degree of adsorption
of dodecyl sulfate salts on oxide minerals is a function of pH, counterion type, and
counterion concentration. Ionic surfactant aggregate formation is favored at higher
counterion concentrations. There is also a tendency for monovalent counterions to adsorb
between the surfactant aggregate and the mineral surface. The extent of adsorption
process depends on the pH as well as the nature of the counterion type as it is determined
by a combination of steric and surface complexation effects. The planar geometry of
admicelles present is expected to provide for more favorable steric interactions compared
to spherical micelles in the bulk solution.
Clay intercalation
Self-assembly is the process whereby small pre-existing subunits spontaneously organize
themselves into an ordered state or structural arrangement. The formation of lamellar
micelles by surfactants molecules in solution is a typical example [18]. Interactions that
promote self-assembly include electrostatic attractions, hydrogen bonding, and
hydrophobic interactions among others [19]. Intercalation is defined as a reversible
insertion of mobile guest species into a crystalline host lattice during which the structural
integrity of the latter is formally conserved [20].
Adsorption and intercalation of surfactants in LDH
Pavan et al. [21] concluded that adsorption on LDH mirrors the surfactant adsorption
behavior on mineral oxides with respect to the effects of pH, counter ion type, and ionic
strength [17]. Dèkány and Haraszti [22], and Pavan et al. [21, 23, 24] found that anionic
surfactants such as sodium dodecyl sulfate (SDS) and sodium dodecylbenzenesulfonate
(SDBS) adsorb on the LDH–CO3 crystal surfaces rendering them hydrophobic. Pavan
et al. [21] explicitly state that these surfactants do not intercalate when the LDH contains
the difficult to exchange carbonate anion. This was confirmed by Ulibarri et al. [25] with
respect to the interaction between dodecylbenzenesulfonate (SDBS) and Mg3Al–LDH–
CO3 at ambient conditions. It seems that temperature plays an important role too. While
at 25 °C SDS only absorbs on LDH–CO3 [21] some intercalation was observed when the
mixture was heated to 70 °C [26]. However, SDBS still did not intercalate at the latter
temperature [26], but partial replacement of carbonate did occur under reflux conditions
and very long reaction times [10]. Apparently the high affinity of the layered double
hydroxide for the carbonate ion prevents the latter’s displacement by the sulfonate ions
even at high contact temperatures. Xu and Braterman [10] argue that the replacement of
carbonate with RSO3 − is kinetically rather than thermodynamically controlled as the
carbonate is very tightly bonded [27, 28], and its removal requires the scaling of a high
activation energy barrier.
Direct intercalation of LDH clays involves ion exchange. It is regarded as a form of
spontaneous self assembly of the guest molecules or ions between the brucite-like layers
of the crystal lattice [29]. Layered host lattices can adapt to the geometry of the inserted
guest species by adjustment of the interlayer separation. Intercalation of organic
compounds creates diverse types of supramolecular structures in the clay interlayer [30].
Linear molecules with appropriate functional groups self-assemble into monolayers or
bilayers between the sheets of LDH [20]. The intercalation of surfactants such as SDS
from aqueous solution can be viewed as a change in the nature of the micelle structure
from spherical to lamellar [15]. The internal hydoxy sheet surfaces are similar to the
surface planes presented to the outside. Thus, it is expected that intercalation behavior
will mimic that of the adsorption process. For example, it is found that monovalent
anions and water co-intercalate [8, 9, 31]. However, there are key differences. The
intercalating guest molecule is now affected by, and must interact with, two parallel clay
surfaces in its vicinity. This constraint enforces a greater degree of order within the
galleries than is required at the free outside clay surfaces. Displacing multivalent resident
ions may also be more difficult. They may neutralize charges on opposite sheets thereby
making parting, to allow larger guests to move in, more difficult. In addition, carbonate
ions also interact strongly with both surfaces via hydrogen bonding forces [32, 33].
Guest surfactant molecules usually assemble in either a monolayer or a bilayer format.
The actual monolayer arrangement of surfactant chains in intercalated LDH corresponds
to two interdigitated anti-parallel half-monolayers [10, 34]. Kopka et al. [3] derived
equations for estimating the basal spacing of mono- and bilayer intercalated clays based
on the following assumptions: (i) Alkyl chain substituents assume an extended chain
conformation; (ii) the methylene bond length equals to 0.127 nm, and (iii) the slant angle
is independent of the chain length.
(1)
(2)
Here, d L is the basal spacing, n represents the number of carbon atoms in the aliphatic
chain; α is the chain tilt angle to LDH layer plane; d o measures the vertical dimension of
the head group taking into account its relative intercalated orientation; d 1 and d 2 are the
distances between the center plane of the brucite-like sheets and the terminal (ionized)
head group and tail ends respectively; d 3 is the distance between the two facing terminal
methyl groups in the bilayer structure. Note that both d 1 and d 2 can be affected by the
presence adsorbed water, solvents molecules or other ions in the galleries.
The extended length of SDBS is about 2.26 nm [35]. Meyn et al. [36] suggest that
intercalated dodecylbenzenesulfonate ions orient their extended alkyl chains into a
perpendicular position with the benzene ring tilted toward the layer. This proposed
arrangement was supported by You et al. [37]. Instead, Xu and Braterman [10] and Zhao
and Nagy [38] contend that (i) the benzene rings are oriented perpendicularly to allow for
three point attachment of the sulfonate group to the hydroxide layer, and (ii) that the alkyl
chains are tilted at ca. 56° with respect to the layer planes in order to facilitate their close
packing. They argue that such anti-parallel arrangement also reduces the electrostatic
repulsion between the anion head groups and effectively maintains the hydrophobic
interactions between the hydrocarbon chains. These two proposals can be tested by
plotting the basal spacing against the chain carbon number (Fig. 1). According to Eqs. 1
and 2, the tilt angle can be calculated from knowledge of the type of intercalation and the
slope of the d L versus carbon number plot. Least square data fitting of the data shown in
Fig. 1 yielded α values of 60.0°, 68.8°, and 61.8° for alkyl sulfonates in Mg2Al–LDH
[39]; alkyl sulfates in Zn2Al–LDH [3], and alkylbenzene sulfonates in Mg2Al–LDH [36],
respectively. These values, while somewhat higher than the expected angle of 56°, do
provide support for the Xu and Brateman’s [10] proposal.
Fig. 1 Effect of alkyl chain length on the basal spacing: (▲) Alkyl sulfonates in Mg2Al–
LDH [39]; (○) alkyl sulfates in Zn2Al–LDH [3]; (◊) alkylbenzene sulfonates in Mg2Al–
LDH [36]; LDH–”carboxylates” prepared in the presence of SDS (●) or SDBS(Δ) (this
work and Nhlapo et al. [61])
The LDH–DS and LDH–DBS data collated in Table 1 show a considerable spread in the
reported basal spacing values. Clearly the observed d-spacing of LDH–CO3 and other
intercalated derivatives depends on a number of factors. Much of the variation can be
attributed to differences in the degree of hydration, i.e., the presence or absence of
interlayer water [3, 36, 40, 41]. It also explains the variations induced by the drying
procedure including the effect of temperature drying temperature: Meyn et al. [36]
indicate that vacuum drying at 65 °C reduces the basal spacing of about 0.3 nm owing to
removal of the adsorbed water from the interlayer. Interestingly, the aliphatic chain tilt
angle is also affected by the presence of the interlayer water [41]. Zhao and Nagy [38]
state that the intercalation pH influences the d-spacing of intercalates prepared by the
coprecipitation method. For LDH–DS in particular it is claimed that the interlayer
spacing is affected by the method of synthesis [41], the LDH Mg:Al ratio [38], and the
intercalation pH [31, 38]. Clearfield et al. [31] found that, as the exchange pH was
increased, so did the basal spacing although, surprisingly, the amount of SDS in the
interlayer decreased. This communication will attempt to provide a rationalization for this
unexpected and unexplained pH dependence.
Table 1 Effect of the intercalation method and Mg:Al ratio on the d-spacing of LDH–DS
and LDH–DBS
Intercalation method d-Spacinga (nm)
Reference
LDH Mg:Al ratio
LDH–DS LDH–DBS
Reconstitution
[54]
0.250
2.64
[50]
0.250
3.01
[25]
0.330
2.68
2.95
[57]
0.333
2.62
2.95
[56]
0.340
2.6
[21]
0.171
4.03
[10]
0.205
3.66
0.250
2.63
Coprecipitation
0.250
[42]
2.68
[70]
0.254
2.54
[38]
0.290
2.63
[92]
0.301
2.60
[38]
2.92
[38]
Ion exchange
0.171
0.175
2.74
0.204
2.74
0.205
2.78
0.250
2.27
0.250
0.254
[38]
2.66
[37]
2.95
[36]
2.43
0.256
[50]
[38]
2.66
[50]
0.301
2.58
0.323
2.09
Failedb
[26]
0.333
2.42
2.96
[49]
0.333
2.6
0.333
[38]
[93]
2.87
[50]
0.333
3.05
[10]
Where applicable as prepared before drying
b
Intercalation failed
a
LDH surfactants intercalation methods
Miyata and Kumura [7] first intercalated LDH with different organic anions. Newman
and Jones [41] and Crepaldi et al. [15] give overviews of surfactant intercalation in LDH.
Intercalated LDHs can be prepared by direct synthesis methods, e.g., hydrothermal
crystallization of gels formed by the coprecipitation of the M2+ and M3+ hydroxides in the
presence of the required organic anion [10, 15, 31, 38, 42, 43]. Hussein et al. [44] used
microwave heating to accelerate the co-precipitation driven intercalation of SDS into
Zn4Al–LDH.
Indirect methods utilize suitable LDH precursors prepared by direct synthesis. Indirect
intercalation involves the modification and treatment of the host and finally the insertion
of the guest molecules inside the layer. Crepaldi et al. [45] identified three main indirect
techniques: (i) direct anion exchange; (ii) LDH reconstruction from a layered double
oxide form obtained by calcinations of a suitable precursor; and (iii) anion replacement
by elimination of the precursor interlamellar species.
(i) Direct anion exchange: Direct ion exchange was pioneered by Miyata and Kumura
[7]. First consider the exchange of simple anions, e.g., the replacement of nitrate ions in
LDH–NO3 with chloride ions from aqueous solution. The mass action relationship is
written as:
Here {LDH} represents a clay subunit commensurate with a single positive charge or Al
atom. In this communication {LDH} = {Mg2Al(OH)6} is used as archetype for
discussion and illustration purposes. The selectivity coefficient the intercalation of the
chloride anion relative to chloride ion in the ion exchange reaction is defined as follows
[34, 46]:
X Cl and represent the fraction of chloride and nitrate anions present in the clay when
equilibrium has been reached. S Cl and are the equivalent fractions of these ions present in
the aqueous solution. The exchange may also involve heterovalent ions as in Scheme II:
In this case the selectivity coefficient takes the form [46]:
Miyata [46] reports and so that
Thus, direct intercalation of LDH-A is achieved by direct contact with a suitably
concentrated aqueous or non-aqueous solution of the desired anionic surfactant [3, 10, 31,
36–38, 47–49]. Owing to the tenacity by which carbonate is held, it is customary to use
anions such as chloride [37, 50] or nitrate [3, 36, 42, 47]. The intercalation reaction for a
surfactant such as SDS into LDH–Cl as may be expressed as follows [34, 46]:
Kopka et al. [3] describe intercalation by anion exchange of Zn2Cr–LDH–NO3 with alkyl
sulfate ions [C n H2n + 1SO4 − with n = 6, 8, …, 18] as well as dodecyl glycol ether sulfate
ions [C12H25(OCH2CH2) n SO4 − with n = 0, 1, 2, 4]. They claim that the exchange
reaction proceeds to 90–95% of theory. Monolayer intercalation was observed. However,
co-intercalation with alkanols and alkyl amines of similar chain lengths lead to the
formation of bilayer structures. Water and small organic molecules, e.g., diols, NMF,
DMSO, etc., co-intercalated as well. Meyn et al. [36] studied the intercalation of anionic
surfactants by such anion exchange into a wide range of LDH compounds. They observed
that dodecylbenzenesulfonate intercalated as monolayers while secondary alkyl
sulfonates intercalated as bimolecular layers.
You et al. [37] investigated the intercalation of sodium octyl sulfate (SOS), SDS 4octylbenzenesulfonate (SOBS) and SDBS into Mg3Al–LDH–Cl via ion exchange in
aqueous medium. They found that the equilibrium amount of surfactant intercalated
decreased in the order SDS > SOBS > SDBS > SOS. SOS also formed bilayers but the
others exhibited monolayer arrangements. Xu and Braterman [10] prepared Mg2Al–
LDH–DBS. The intercalated dodecylbenzenesulfonate product had a d-spacing of
3.05 nm consistent with anti-parallel monolayer packing of interpenetrating chains.
Crepaldi et al. [45] describe a variation of the ion exchange method based on the
formation and organic phase extraction of a salt between dodecyl sulfate and a cationic
surfactant.
Anbarasan et al. [26] attempted the direct ion exchange reaction using Mg2Al–LDH–CO3
at 70 °C. They found no evidence for the intercalation of SDBS. The XRD spectrum of
the product obtained using SDS features new peaks at lower angles consistent with a
basal spacing of 2.09 nm which Anbarasan et al. [26] interpret as providing evidence for
some monolayer intercalation of SDS.
True ion exchange reactions are topotactic in nature implying that any layer stacking
defects in the precursor will also appear in the pillared LDH product [51]. However, Xu
and Braterman [10] observed changes in crystal habit on intercalating SDS in LDH at
elevated temperatures. This implies that at least some recrystallization must have
accompanied the intercalation process.
(ii) The LDH reconstruction method: Comprises a hydrothermal reconstitution of
calcined LDH carbonates in the presence of the desired anion in carbonate-free water [15,
21, 37, 51–57]. As discussed above, carbonate anions do not readily ion exchange owing
to strong electrostatic and hydrogen-bonding interactions. Calcining a suitable precursor,
e.g., LDH–CO3 at 400–500 °C produces a dehydroxylated and decarbonated layered
double oxide (LDO) form. From this the original clay can be reconstructed, in an
intercalated form, by treatment with an aqueous solution of the required anion. The
mechanism is believed to entail the fast rehydration of the oxide with intercalation of OHanions, followed by a slow anion exchange of the latter with other anions [15]. The
formation of the pure hydroxide (LDH–OH) form requires reconstruction in pure water
and total exclusion of CO2, e.g., a nitrogen atmosphere [51]. Chibwe and Jones [54]
intercalated SDS, p-toluene sulfonate, and other ions using the reconstitution method.
Chibwe and Jones [54], Dimotakis and Pinnavaia [51], and Hansen and Taylor [58] claim
that intercalation of anionic surfactants and other anions into the LDH–OH form is
facilitated by the presence of glycerol. You et al. [37] and Costa et al. [57] point out
shortcomings of the reconstruction method. When applied to organic anions, mixed
phases may be produced [54], and it is difficult to avoid formation of carbonate forms.
Products often show a perforated surface morphology and feature broad XRD peaks
indicative of poorly developed crystallinity. This communication will show that, for the
intercalation of SDS, the reconstruction method holds additional drawbacks.
(iii) Anion exchange by elimination of a precursor interlamellar species: Carlino et al.
[59, 60] and others [26, 61, 62] showed that thermal intercalation takes place when the
LDH–CO3 is brought into direct contact with pure molten organic acids. This method
has, as far as we could ascertain, not yet been used to intercalate anionic surfactants of
the sulfate or sulfonate type.
LDH applications
LDH-type anionic clays, as such or in their calcined form, have existing and many
potential applications [63]. In medicine they are utilized as antacids and antipeptins. In
polymer technology they function as halogen scavengers, flame retardants, and PVC
stabilizers. They are employed as catalysts and catalyst supports and their absorbent and
ion exchange properties are of interest in waste water treatment. Surfactant intercalated
layered double hydroxides are of interest for a variety of reasons [41]. Hydrophobization
of the LDH by ion exchange yields new types of thickening agents. It also facilitates
sorption of nonionic organic compounds [30], e.g., trichloroethylene and
tetrachloroethylene [50]. The distinctive properties of these modified layered double
hydroxides permit a wide range of uses including polymer additives [64], precursors for
catalysts [63, 65], and magnetic materials [64]. Nanocomposites can be prepared by
exfoliation within polymer matrices [57, 65–69]. Their generally non-toxic nature and
membrane-like structure can be harnessed to protect, carry, deliver, and controllably
release active compounds such as pesticides [70], pharmaceuticals and even genes [65,
71]. They can be used as adsorbents to remove contaminants from water [38, 56].
Due to the wide-ranging utility of LDH intercalates, it is of interest to consider
environmentally friendly and energy efficient methods of intercalation that yield products
of an acceptable quality. LDH–CO3 is currently available as a bulk raw material owing to
growing PVC stabilizer applications. Our interest is in upgrading this basic starting
material by suitable intercalation procedures. Previously we reported on the surfactantmediated intercalation of long chain fatty acids into LDH–CO3 [61]. The method is a
refinement of the Carlino et al. [59, 60, 72] melt intercalation method with anion
exchange facilitated by elimination of the carbonate species. This communication reports
on a similar approach based on contacting aqueous suspensions of LDH–CO3 with
combinations of a short chain aliphatic carboxylic acid and an anionic surfactant, e.g.,
SDS or sodium dodecylbenzene sulfate (SDBS). It was found that this simple procedure,
conducted at ambient conditions without exclusion of CO2, provides a facile one-pot
method for intercalating these surfactants in LDH.
Experimental
Materials
LDH–CO3 (Hydrotalcite Grade HT 325) was supplied by Chamotte Holdings. It
contained silica and magnesium carbonate as minor impurities. Distilled water was used
in all experiments. High purity (>98%) SDS and SDBS were purchased from Fluka–
Biochemika. Croda chemicals supplied the octanoic and dodecanoic acids. Myristic acid
was obtained from BDH Chemicals. Butanoic acid (>99%) was supplied by Merck. AR
grade glacial acetic acid and acetone (99.5%) were supplied by Saarchem UnivAR.
Aqueous ammonium hydroxide solution (25%) was procured from Promark chemicals.
Potassium bromide (Uvasol KBr, Merck) was used for preparing samples for FTIR
spectra recording.
Sample preparation
Surfactant intercalation experiments were carried out using variations of the following
representative procedure: 75 g SDS (0.26 mol) and 15 g acetic acid (0.25 mol) were
dispersed in 1.5 L distilled water and the pH adjusted to pH = 10. To this 20 g HT-325
(LDH–CO3 approximating [Mg0.66Al0.34(OH)2](CO3)0.17 · ½H2O (ca. 0.10 mol Al) was
added slowly while stirring. The emulsion–suspension was left to stir overnight. The pH
was again adjusted to pH = 10 each morning by adding dilute ammonia or NaOH solution
if required. It was noted that pH dropped to as low as pH = 7.2 overnight. The mixture
was allowed to react at ambient temperature for a total of 2 days. The product was
recovered by centrifugation, washed four times with distilled water, and once with
acetone. After each washing the solids were separated from the liquid by centrifugation.
The product was allowed to dry at room temperature. This experiment was repeated
leaving out the acetic acid or the SDS. The effect of raising the reaction temperature to
65 °C or 80 °C as well as using reduced amounts of the SDS and/or the acetic acid was
also investigated. Similar experiments were done using sodium dodecyl sulfate (SDBS)
in place of the SDS or magnesium hydroxide in place of LDH–CO3. Fatty acid
intercalation experiments followed procedures similar to the preparation of LDH–laurate:
20 g LDH–CO3 (0.10 mol Al), 40 g SDS (0.26 mol), 76.9 g lauric acid (0.38 mol), and
20 g HT-325 were dispersed in 1 L distilled water at 70 °C and allowed to stir
continuously for 3 days. Lauric acid was divided into three equal portions. One part was
added at the start of the experiment and the two other portions added every subsequent
day. When required, dilute NH4OH was added to the mixture in order to maintain the pH
at pH = 10 ± 0.5.
Additional LDH–DS and LDH–DBS samples were prepared by the regeneration method
described by Costa et al. [57]. In this case LDH–CO3 was first calcined at 450 °C for 3 h
and then stirred in a suspension of the relevant surfactant. These samples were analyzed
by TG and XRF to determine organic and the sodium contents, respectively. Samples
prepared using LDH with x ≈ 0.33 were donated by Dr Costa and analyzed as such.
A mixture of LDH–stearate and magnesium stearate was prepared as follows: LDH was
intercalated with stearate using the standard procedure described above by reacting 20 g
HT 325 with 40 g stearic acid in the presence of 54.4 g SDS at 80 °C. After completion
of the reaction, 56.54 g Mg stearate was added over a period of a further 2 days.
Thereafter the mixture was allowed to stir for another 4 days at 80 °C.
Characterization
The particle size distribution and BET surface area of the precursor LDH were
determined using a Malvern Mastersizer Hydro 2000MY instrument and a Micromeritics
Flowsorb II 2300 instrument, respectively.
Elemental composition was determined by XRF analysis. The intercalated materials were
ashed before analysis in order to reduce their bulk. These samples were ground to
<75 μm in a tungsten carbide mill and roasted at 1,000 °C. Then, 1 g sample was added
to 9 g Li2B4O7 and fused into a glassed bead. Major element analysis was executed on the
fused bead using an ARL9400XP + spectrometer.
Powder samples were viewed on a JEOL 840 SEM scanning electron microscope under
low magnification. They were prepared as follows: A small quantity of the powder
products or the LDH–CO3 precursor was placed onto carbon tape on a metal sample
holder. Excess powder was removed using a single compressed air blast. The samples
were then coated five times with gold under argon gas using the SEM autocoating unit
E5200 (Polaron equipment LTD).
Thermogravimetric analysis was conducted on a Mettler Toledo A851 TGA/SDTA
machine. Powder samples of ca. 10 mg were placed in 70 μL alumina open pans.
Temperature was scanned at 10 °C/min in air range from 25 to 800 °C.
FTIR spectra were recorded on a Bruker Opus Spectrophotometer. Samples were finely
ground and combined with spectroscopic grade KBr in a ratio of 1:50, i.e., approximately
2 mg of sample and 100 mg of KBr. The mixture was pressed into a 13 mm φ pellet. The
reported spectra were obtained over the range 400–4,000 cm−1 and represent the average
of 32 scans at a resolution of 2 cm−1.
XRD
Phase identification was carried out by XRD analysis on a PANalytical X-pert Pro
powder diffractometer. The instrument features variable divergence and receiving slits
and an X’celerator detector using Fe filtered Co K-α radiation (0.17901 nm). The X’Pert
High Score Plus software was used for data manipulation.
Py/GC/MS
Small samples (3–4 mg) were analyzed on an Agilent GC/MS system fitted with a DB17MS intermediate polarity GC column (30 m × 0.25 mm ID), an Agilent MSD 5971
mass spectrometer, and a CDS Instruments Pyroprobe 2000 pyrolyzer. Helium was used
as carrier gas (1 mL/min; split 1:20).
Results
Table 2 lists sample designations and the thermal properties of the major compounds
used or synthesized in this study. Table 3 presents XRF composition data as atom ratios
relative to aluminum. The identification of the compound natures, as implied by the
designations, is justified by the results presented below. Unless otherwise indicated, all
LDH–surfactant sample names refer to products prepared using acetic acid as
intercalation aid and all LDH–carboxylates refer to samples prepared from the
corresponding acid using SDS as mediating surfactant.
Table 2 Sample designations, XRD determined basal spacings, and thermogravimetric
data
TG residual mass
Intercalated compound
Clay Organic
a dL
(%)
Method
(carboxylic acid)
(nm)
(%)
(%)
150 °C 700 °C 800 °C
Inorganic precursors
LDH–CO3 (HT 325)
0.763 98.09
59.15
58.82
100
–
LDH–CO3 (Costa)
0.759 99.03
56.16
55.77
100
–
SDS experiments
LDH + SDS (no acid)
e
0.760 98.41
58.68
–
98.9
1.1
LDH–DS (Costa)
r
2.69
94.09
65.03
64.70
na
na
LDH–DS (this study)
r
2.69
92.92
55.26
55.01
98.7
1.3
LDH–DS (acetic acid)
e
2.60
93.52
41.71
41.37
73.8
26.2
LDH–DS (butyric acid)
e
2.59
92.52
33.33
33.02
59.5
40.5
LDH–octanoate
e
2.72
91.92
23.62
–
42.6
57.4
LDH–laurate
e
3.66
90.67
14.71
–
26.9
73.1
LDH–myristate
e
4.25
93.72
25.93
45.9
54.1
LDH–stearate
e
4.94
92.37
16.09
28.7
71.3
LDH–behenate
e
6.07
95.7
9.00
15.6
84.4
LDH–DBS (Costa)
r
3.07
91.57
41.06
40.85
74.4
25.6
LDH–DBS (this study)
r
3.04
92.41
45.77
45.50
81.5
18.5
LDH–DBS (acetic acid)
e
2.88
93.55
37.99
36.92
65.8
34.2
LDH–DBS (butyric acid)
e
2.84
92.52
33.33
33.02
59.5
40.5
LDH–DBS (hexanoic acid) e
2.84
92.99
38.68
38.24
68.6
31.4
LDH–DBS (lauric acid)
3.64
91.69
18.56
18.30
33.0
67.0
63.1
36.9
15.89
SDBS experiments
e
LDH–(DBS + dodecyl
e
3.15 93.64 36.17 35.84
alcohol)
a
e = acid assisted decarbonation method; r = regeneration method
Table 3 XRF results with composition expressed as atom ratios relative to aluminum
Intercalation Method
Mediated ion exchange Regeneration
Atom
LDH LDH–laurate LDH–DS LDH–DBS LDH–DS LDH–DBS
Mg
1.89 2.31
1.87
2.16
2.27
2.13
S
0.024 0.027
0.356
0.266
0.467
0.127
Na
0.005 0.018
0.007
0.000
0.041
0.039
Si
0.051 0.050
0.041
0.047
0.049
0.062
Ca
0.002 0.021
0.002
0.004
0.018
0.018
Ni
0.003 0.002
0.002
0.004
0.004
0.004
Fe
0.003 0.004
0.003
0.003
0.004
0.003
x
0.346 0.278
0.349
0.317
0.304
0.318
The Malvern particle size analysis of the LDH–CO3 precursor revealed a bimodal particle
size distribution with d(0.1) = 1.0 μm, d(0.5) = 3.5 μm, and d(0.99) = 260 μm. The
measured BET surface areas were 7.9, 21.6, 17.0, 15.7, and 5.3 m2/g for the Mg(OH)2,
LDH–CO3, LDH–DS, LDH–DBS, and LDH–laurate, respectively. Table 3 reports the
chemical composition of the precursor LDH and the intercalated products as determined
by XRF analysis. The results are presented as atom ratios relative to the aluminum
present. The data for the precursor indicate a value for x = 0.346 in the chemical formula
[Mg1–x Al x (OH)2](CO3) x/2 · nH2O. XRF analysis points to values for x of 0.302, 0.349,
and 0.317 for LDH–laurate, LDH–DS, and LDH–DBS, respectively.
Figure 2 shows SEM micrographs of the LDH–CO3 precursor and selected intercalates.
The LDH–CO3 features the typical sand rose structure formed by numerous inter-grown
small crystallites [73]. The surfactant modified powders also show agglomerated platelets
that are similar in size to those seen in the precursor. The LDH–laurate platelets differ in
that they feature significantly larger lateral dimensions. Intercalates obtained by the
regeneration method have an unusual perforated surface morphology.
Fig. 2 SEM micrographs of (a) the LDH–CO3 precursor; the products prepared by acetic
acid aided intercalation (b) LDH–DS, (c) LDH–DBS, (d) LDH–laurate; and products
obtained by the regeneration method: (e) LDH–DS and (f) LDH–DBS
Nhlapo et al. [61] described surfactant-assisted intercalation of fatty acids in the readily
available LDH–carbonate. The carboxylic acids displace the basic carbonate anions and
intercalate in bilayer form. Nhlapo’s [61] method is, in essence, a refinement and
improvement of the Carlino [59, 60] melt intercalation procedure. The intercalation
reactions are conducted at temperatures just above the melting point of the acid
concerned, with the LDH powder suspended in the acid oil-in-water emulsion.
Surfactants, e.g., SDS, facilitate the intercalation process by emulsifying the molten acid
and dispersing the LDH particles. This convenient and environmentally friendly method
for carboxylic acids has several attractive features: Water is used as medium rather than
organic solvents, clay calcinations are not necessary, and there is no need for working
under a CO2-free atmosphere.
It is well-established commercial practice to react magnesium hydroxide with molten
stearic acid to produce magnesium stearate. In addition, LDH–CO3 appears capable of
intercalating variable amounts of stearate. This can exceed the anion exchange capacity
by multiple factors [8, 9, 61]. Even so these products always contain some unreacted
LDH–CO3 as an impurity phase. These observations raise the question as to whether the
Carlino melt method [59, 60] and Nhlapo’s [61] procedure in fact deliver an intercalated
product or just simply mixtures of magnesium and aluminum stearates together with
unreacted LDH–CO3. Figure 3 provides a possible answer to this question. It shows
powder XRD spectra for recrystallized magnesium stearate, LDH–stearate (prepared in
the presence of SDS), and a combination that contains approximately equal amounts of
these two products. In the preparation of the latter mixture, the two compounds were
heated together for an extended period of time in a water suspension containing SDS. The
XRD spectra of the first two samples appear very similar. However, the spectrum for the
mixture shows clear twinning of the basal reflections. This positively demonstrates the
presence of two different phases and supports the assumption that the interaction of
molten stearic acid with LDH yields the intercalated product instead of the metal soaps.
Fig. 3 Powder XRD spectra for recrystallized magnesium stearate, LDH–stearate, and a
mixture of the two compounds. The LDH–stearate contains minor amounts of stearic acid
as an impurity
Figure 1 provides evidence that intercalation favors incorporation of long chain
carboxylic acids above the anionic surfactants SDS and SDBS. The observed basal
spacings (d L) for the products obtained with either surfactant (SDS or SBS) are the same
provided the aliphatic acid chain is sufficiently long. The dependence of d L on the
number of carbons in the carboxylic acid (n) was determined by a least square curve fit to
the SDS data and yielded:
(3)
The magnitude of the observed basal spacing values implies bilayer intercalation. The
slope Δd L/Δn = 0.241 is consistent with a chain tilt angle of ca. 71.6° to the plane of the
clay sheets [61].
Anbarasan et al. [26] previously found that SDBS on its own does not intercalate in LDH
in aqueous suspension even when the reaction temperature is raised to 70 °C. However,
they claim that some intercalation of SDS occurs under similar conditions but the basal
spacing of the product was anomalously low (d L = 2.09 nm cf. 2.67 nm). In the present
study, intercalation of neat SDS and SDBS, as well as acetic acid on its own, was
attempted under ambient conditions and pH = 10. The experimental basal spacing and TG
data obtained for these products are presented in Table 2. They are in substantial
agreement with the values determined for the precursor compounds, and thus indicate
that no discernable intercalation occurred. However, the results were markedly different
when the LDH was suspended in aqueous medium in the presence of mixtures of one of
the surfactants together with a lower acid. Figure 1 shows that the basal spacing of the
products obtained with acetic, butyric, or hexanoic acid, deviate considerably from the
straight line dependence predicted by Eq. 3. For SDBS as surfactant, the corresponding dvalues agree with each other to within experimental error (d L ≈ 2.86 nm). For SDS as
surfactant, d L ≈ 2.58 nm for acetic and butyric acids. The experimental d L values are in
reasonable agreement with basal spacing values reported for LDH–DS and LDH–DBS as
prepared by other methods (See Table 1). This implies that the presence of lower
aliphatic acids facilitates intercalation of SDBS and SDS in LDH–CO3 under mild
conditions (ambient temperature and aqueous medium at pH < 10).
Figure 4 compares the X-ray diffractograms recorded for LDH–DS, LDH–DBS (both
prepared in the presence of acetic acid), and LDH–laurate, with that for LDH–CO3. The
reflections at 0.76 nm (2θ = 13.2o) and 0.38 nm (2θ = 27.2o) are characteristic of LDH–
CO3. They are also present in the LDH–surfactant compounds indicating that they
contain LDH–CO3 as an impurity.
Fig. 4 X-ray diffractograms for LDH–CO3, LDH–DS, LDH–DBS, and LDH–laurate
obtained by the acetic acid mediated carbonate elimination method
Figure 5 shows the thermogravimetric traces for the LDH–CO3 and some intercalates.
The thermal decomposition of LDH–CO3 occurs in three steps corresponding to loss of
adsorbed and interlayer water, dehydroxylation, and a combination dehydroxylation–
decarbonation reaction, respectively [52, 74, 75]. Thermal degradation of the intercalated
layered double hydroxides takes place in several steps too. The first step is attributed to
loss of interlayer water and is assumed complete at a temperature of 150 °C [59, 64, 76,
77]. Mass loss is effectively complete at 700 °C. The final residues may, to a first
approximation, be assumed to have the same composition as the ash of the precursor
provided sodium ions did not co-intercalate. This assumption allows one to make a rough
estimate of the organic content of the initial sample and these are reported in Table 2.
Fig. 5 Thermogravimetric (TG) mass loss curves for LDH–CO3, LDH–DS, LDH–DBS,
and LDH–laurate (prepared by the acetic acid mediated carbonate elimination method)
obtained at a scan rate of 10 °C/min in an air atmosphere
All samples show gradual and progressive mass loss as the temperature is raised above
50 °C. Mass loss rates accelerate above 170, 210, 250, and 270 °C for LDH–laurate,
LDH–DS, LDH–CO3, and LDH–DBS, respectively. Below 550 °C the mass loss of
LDH–DS exceeds that of LDH–DBS. The point where the mass loss rate accelerates
cannot be regarded as a threshold limit for the stability of the LDH. Pyrolysis GC/MS of
LDH–SDBS performed at 200 °C already reveals the liberation of a range of branched
alkyl benzene compounds, i.e., typical SDBS degradation products. Application of this
technique also confirmed that the SDBS did not co-intercalate with stearic acid. No aryl
derivatives were found in the pyrolysis products of LDH–stearate obtained at 245 and
350 °C. Only aliphatic compounds were detected with the two major compounds
identified as similar to C16 and C18 methyl esters.
The degree of aluminum substitution of the clay, characterized by the value that x
assumes in the formula [Mg1–x Al x (OH)2](CO3) x/2 · zH2O, also indicates the anion
exchange capacity of the material. If intercalation had proceeded to completion, the
expected residue levels on a dry, i.e., dehydrated clay basis are 63.4, 29.0, 26.9, and
34.3% for LDH–CO3, LDH–DS, LDH–DBS, and LDH–laurate, respectively. This may
be compared against the experimentally determined values of 60.0, 44.1, 39.1, and 18.9.
The degree of intercalation was estimated using the values for x from the Al/Mg ratios
indicated by XRF analysis (See Table 3) together with the residue levels determined by
TG (See Table 2). The calculated values are 0.37, 0.46, and 2.51 for LDH–DS, LDH–
DBS, and LDH–laurate. Thus, the extent of laurate intercalation exceeded the anionic
exchange capacity of the clay by ca. 2.5 times. This high value is attributed to
concomitant intercalation of non-ionized lauric acid and/or sodium laurate to provide for
tight packing of the alkyl chains inside the clay galleries [8, 9, 61]. The sulfur to
aluminum atom ratio should also provide an indication of the degree of intercalation of
the anionic surfactants. Indeed, the S/Al value for LDH–DS is in good agreement with
the degree of intercalation estimated from TG data. However, the S/Al = 0.266 value
determined via XRF analysis for LDH–DBS is significantly lower than expected. It is
possible that some sulfur may have been lost during the de-bulking heat treatment of the
samples.
Figure 6 compares the FT-IR spectra of the unmodified layered double hydroxide (LDH–
CO3) with the LDH–surfactant intercalates. Costa et al. [57] provide a comprehensive
analysis of the infrared absorption bands relevant to the present compounds. Therefore,
the present discussion is limited to a short overview. The characteristic 446 cm−1 M–O
lattice vibration band is present in all the samples. This is consistent with an intact LDH
sheet structure. A broad band in the region 3,200–3,700 cm−1 is observed in all the
compounds. It is attributed to OH stretching vibrations of the octahedral layer and
intercalated water molecules [32, 78]. The shoulder at 3,063 cm−1 indicates hydrogen
bonding of H2O to CO3 2− ion in the interlayer space [79–81]. As expected, the carbonate
peak located at 1,367 cm−1 is well developed in LDH–CO3. Its presence in the LDH–
laurate indicates the unreacted LDH–CO3 as an impurity. The triplet peaks observed in
the range 2,850–2,965 cm−1 are due to C–H stretching. They confirm the presence of the
alkyl chains of the surfactant anions in the intercalated LDH derivatives [31, 56, 57, 82].
Generally organic sulfate and sulfonate groups exhibit frequencies at 1,200–1,180 cm−1
[15, 31] and 1,420–1,370 cm−1 [26, 56, 70]. The former band is conspicuously absent in
the LDH–laurate spectrum. This indicates that the LDH preferentially intercalated the
laurate and that DS was not co-intercalated. The presence of a small peak at this position
in the LDH–octanoate spectrum indicates that some DS did co-intercalate.
Fig. 6 FT-IR spectra for LDH–CO3, LDH–DS, LDH–DBS, and LDH–laurate prepared
by the acetic acid mediated carbonate elimination method
You et al. [37] used anion exchange to intercalate SDS and SDBS in LDH–Cl. They
obtained AEC levels of ca. 72% using this method even though the solutions were
sparged with nitrogen. Incomplete intercalation of LDH–DS was previously observed by
Zhao and Nagy [38] using ion exchange and coprecipitation and by Costa et al. [57] for
samples prepared using both anionic surfactants according to the regeneration method. In
fact their data (surprisingly) show for their purported LDH–DS, a TG residue value that
is higher than that found for the precursor LDH–CO3. This indicates that the product
obtained could not have been a pure LDH–DS. We therefore repeated these intercalationby-regeneration experiments and reanalyzed samples supplied by Dr Costa. The TG and
XRD results obtained for these samples are presented in Tables 2 and 3 and compared to
the samples prepared using acetic acid as intercalation aid in Fig. 7. The products
obtained using the current elimination methods are characterized by sharper and more
intense XRD peaks than the regeneration-based samples. This suggests improved
ordering and a better developed crystalline structure.
Fig. 7 Comparing the X-ray diffractograms for LDH–DBS and LDH–DS prepared by
reconstitution and the ion exchange by elimination method
Consider the LDH–DS obtained by the regeneration method. The dominant series of
broad peaks centered at 3.82°, 7.81°, and 8.00° (2.64 nm) in Fig. 7 are characteristic for
monolayer intercalated SDS. The 0.74 nm basal spacing indicated by the peak at 14.0° is
attributable to unreacted LDH–CO3. Careful examination of the XRD spectrum for the
regeneration-based LDH–DS reveals two additional series of XRD peaks. The large basal
spacing (3.83 nm) indicated by the low-intensity, but sharp series of peaks at 2.70°,
5.34°, and 8.00° point to bilayer intercalation. This we attribute to co-intercalation of
dodecanol with the dodecyl sulfate. This assumption is able to explain several anomalies
in our experimental data for LDH–DS prepared by the regeneration method as well as
observations by other investigators [26, 31, 37, 38, 57]. Firstly, while our LDH–DS has a
high sulfur content (atom ratio S/Al = 0.467), the TG residue at 800 °C data indicates a
very low organic content (See Table 2). Thus, there is insufficient dodecyl sulfate present
to account for the high concentration sulfur in the sample. Since magnesium sulfate and
aluminum sulfate are soluble, it is likely that the excess sulfur must be accounted for by
the formation of LDH–SO4. Indeed, the peaks at 11.8° and 23.9° (basal spacing 0.87 nm)
are consistent with the presence of this phase as an impurity [52]. Kopka et al. [3] found
that the alkanols co-intercalate with alkyl sulfates into Zn2Cr–LDH forming a bilayer
arrangement. They observed basal spacing of 4.15 nm for the combination of SDS and
dodecanol in this LDH matrix. Thus, the present value for Mg2Al–LDH is at least in the
right ball park.
What could be the origin of the dodecanol and excess of sulfate ions? It is well known
that sodium dodecyl sulfate hydrolyses at low pH [83, 84]. However, Angarska et al. [85]
report that hydrolysis also occurs under highly basic conditions. Such hydrolysis would
yield both the required sulfate ions and the dodecanol according to the following
reaction:
Clearfield et al. [31] studied the effect of pH on the intercalation of SDS in Ni4Al–LDH–
Cl. It appears that they conducted their reactions at ambient conditions but reaction time
was varied from 4 to 24 h. They found that, as the pH was increased, so did the interlayer
spacing. However, surprisingly, the amount of dodecyl sulfate that was intercalated
decreased. Our suggestion provides a rationalization of the pH effect observed by
Clearfield et al. [31]. Hydrolysis of the surfactant releases sulfate ions that are
preferentially intercalated. This reduces the ability of the clay to absorb SDS. The cointercalation of the resultant dodecanol with the dodecyl sulfate explains the dramatic
increase in the basal spacing. Interestingly, at pH = 9 and pH = 10, they observed basal
spacing values of 36.6 and 42 nm, respectively. The latter value is just slightly larger than
that found by Kopka et al. [3] for the co-intercalation of dodecanol and dodecyl sulfate.
Zhao and Nagy [38] made similar observations with respect to a pH effect in Mg4Al–
LDH–DS and Mg5Al–LDH–DS prepared by co-precipitation. When prepared at pH = 10,
basal spacing values of 36.6 and 40.3 nm were found for these two compounds,
respectively. Zhao and Nagy [38] used SDS in stoichiometric excess (1.5 times) and long
reaction times at elevated temperatures (3 days at 65 °C). The degree of ion exchange
was estimated from total organic content. They observed a decrease in the apparent
degree of ion exchange with increase in the reaction pH. Again, we surmise that the high
basal spacing is caused by the partial hydrolysis of the SDS and the subsequent cointercalation of dodecanol with SDS. If this is indeed the case, the actual degree of ion
exchange would be even lower and attributable to concomitant sulfate intercalation.
Similar arguments may be relevant to explain the low degree of intercalation, as well as
the reported bilayer nature of SOS intercalation, reported by You et al. [37].
Co-intercalation of the sodium salts was found to accompany that of the surfactant anions
when the anion exchange or the direct precipitation methods are used to prepare LDH–
DS [10, 31] or LDH–DBS [10]. However, the amounts tend to be small, e.g.,
Na/Al < 0.065 [10]. Table 3 indicates that even lower amounts of sodium were found in
the present samples prepared using acetic acid as mediating agent. Quite the opposite
holds for the sodium content of the samples prepared by the regeneration method.
Considering Scheme I, this is not entirely surprising. Rehydration of the calcined clay
(LDO) oxide initially results in rapid formation LDH–OH [86] according to Scheme V.
Some magnesium (and aluminum) hydroxide will also dissolve. Both scenarios introduce
excess hydroxyl anions and the pH of the water phase increases. We found that the
hydration of the present LDO caused an initial rapid increase to pH = 10.7 with a slower
rise to pH = 12 over a 24-h period. By comparison, when the same amount of LDH–CO3
was suspended in distilled water, pH = 9.8 decreased to pH = 9.2 after 24 h.
Next these OH− anions may be exchanged with other anions that are present in the
mixture, e.g., intercalation of dodecyl sulfate ion according to Scheme VI:
Parker et al. [87] measured the amount of anion that was absorbed by freshly calcined
LDH after 24 h. They have found that the relative preference for anions follows the
sequence:
This matches the order of preferred affinity of anions in LDH (with x ≈ 0.3) reported by
Miyata [46]:
Note that Bontchev et al. [88], instead, found that Br− > Cl− for an LDH with x ≈ 0.25.
This suggests that the exact order of anion preference may depend to some extent on the
Mg/Al ratio. In either case it is clear that LDH has a high affinity for the divalent sulfate
ions and that they will easily replace hydroxyl ions present in the clay.
The mechanism of acid mediated decarbonation and intercalation
Bish [89] previously reported that the carbonate in LDH–CO3 is readily exchanged with
Cl−, NO3 −, and SO4 2− by treatment with dilute aqueous solutions of the corresponding
inorganic acids. Iyi et al. [90, 91] proposed a credible two-step process to explain the
decarbonation of LDH–CO3 in the presence of dilute acid or acid–sodium salt mixtures.
The first step is protonation of the carbonate and its conversion to the hydrocarbonate.
Simultaneously another anion is incorporated into the interlayer space to maintain overall
charge neutrality. Next the hydrocarbonate ion is removed via ion exchange with the
excess anion present in solution. A basic tenet of this plausible mechanism is that ion
exchange between the clay and the solution is much easier when monovalent ions are
involved. It could even be argued that the conversion of the interlayer carbonate ions into
monovalent ions is a prerequisite: It allows the interlayer spacing to increase to allow
accommodation of the much larger surfactant molecules. Iyi et al. [90] also observed that
a high degree of substitution requires the presence of a large excess of the counterion in
high concentration (>4 mol/L). The ease of substitution also decreases with increase in x,
the fractional aluminum substitution in the layers. These two factors may explain why
only partial replacement of carbonate was achieved in the present study.
Conclusions
Direct intercalation of surfactants in LDH–CO3 is difficult owing to the tenacity by which
the carbonate is held. The regeneration method has been used successfully to replace
carbonate with other guest ions. In this method the LDH–CO3 is heated and converted
into an essentially carbonate-free layered double oxide. This product is then suspended
and stirred in aqueous medium containing the desired anion. This study showed that,
while the method works for dodecylbenzenesulfonate (DBS), problems are encountered
when applying it to dodecyl sulfate (DS). The latter surfactant tends to hydrolyze in
highly basic medium to form sulfate ions and dodecanol. Thus, LDH–DS prepared by the
regeneration method contains LDH–SO4 as an impurity. Furthermore, the liberated
dodecanol may co-intercalate with DS in a bilayer format to give an additional impurity
phase. These insights provide explanations for the anomalous pH effect on the d-spacing
reported for LDH–DS by other investigations.
Iyi et al. [90, 91] showed that the conversion of LDH–CO3 into LDH-A, where A is
another inorganic anion, is facilitated by the presence of dilute acids as decarbonation
aids. It this study this approach was extended to the intercalation of LDH–CO3 with DS
and DBS. It was found that the intercalation proceeded smoothly under mild conditions
of pH and temperature when water soluble carboxylic acids were added to aqueous
suspensions LDH–CO3 and surfactant. Compared to the regeneration approach, wellcrystallized products with improved purity were obtained. However, the degree of
carbonate substitution that was achieved did not exceed 50%. It is well established that
much purer products can be obtained using (i) direct synthesis by coprecipitation or (ii)
ion exchange starting with LDH–Cl or LDH–, i.e., LDH precursors with more easily
exchangeable monovalent anions, e.g., Cl− or NO3 −. These latter methods should be
considered when high purity LDH–DS or LDH–DBS are sought.
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