Several methods have been reported in the literature for the synthesis of guanidinium salts [
20], including desulfurization-addition sequences, the addition of amines to electrophiles such as carbodiimides, cyanamides or chloramidinium salts [
21,
22], and the utilization of guanylation reagents [
23]. The synthesis route employed here made use of 1
H-Pyrazol-1-carboxamidine hydrochloride as guanylating reagent [
24]. This method allowed a nearly quantitative transformation of the starting material to the desired product under mild reaction conditions and afforded guanidinium salts in high purities and high yields [
25].
The synthesized guanidinium chlorides are capable of forming birefringent phases in water, as can be seen in
Figure 1. This behavior, although only evaluated qualitatively, indicates that guanidinium halides form anisotropic lyotropic phases. Additionally, the thermotropic behavior of these compounds was studied using differential scanning calorimetry. The formation of thermotropic mesophases is indicated by the C→LC and LC→I phase transitions of DDGC, TDGC and CGC (see
Table 1). In this context, DGC has a particular position, since, although forming lyotropic solutions in aqueous medium, it does not form thermotropic phases, as shown by the absence of LC→I phase transition in its DSC thermogram.
Figure 1.
Polarizing micrographs of DGC at 2.5× original magnification: (a) air phase; (b) isotropic phase; (c) anisotropic phase.
2.1. Krafft Temperature, Critical Micelle Concentration, and Micelle Dissociation
Figure 2 illustrates the variations of the specific conductance of micellar aqueous solutions as a function of the temperature for the guanidinium surfactants studied in the present paper. With the sole exception of DGC, these plots exhibit a break, which indicates a noticeable change in the surfactant solubility and may be used to determine the Krafft temperature,
TK [
26,
27]. The following values have been inferred from the temperature dependency of the specific conductance: 292 ± 1 K (19 ± 1 °C), DDGC; 304 ± 1 K (31 ± 1 °C), TDGC; 314 ± 1 K (41 ± 1 °C), CGC. For DDGC, the Krafft temperature is, within the margin of error, in a good agreement with the value reported by Miyake
et al. [
7].
Figure 2.
Plots of the specific conductance against temperature for micellar aqueous solutions of guanidinium cationic surfactants: (a) DGC (m = 31.72 mmol·kg−1); (b) DDGC (m = 8.01 mmol·kg−1); (c) TDGC (m = 2.55 mmol·kg−1); and (d) CGC (m = 0.7 mmol·kg−1).
Figure 2.
Plots of the specific conductance against temperature for micellar aqueous solutions of guanidinium cationic surfactants: (a) DGC (m = 31.72 mmol·kg−1); (b) DDGC (m = 8.01 mmol·kg−1); (c) TDGC (m = 2.55 mmol·kg−1); and (d) CGC (m = 0.7 mmol·kg−1).
Since no such break has been observed in
Figure 2a, the Krafft temperature of DGC should be lower than 276 K (3 °C),
i.e. the lowest temperature to be obtained in the measuring system used. As for classical ionic surfactants [
2], the Krafft temperature increases with the increase of the length of the alkyl chain. Compared to their alkyltrimethylammonium bromide or chloride homologues, the guanidinium type cationics appear to have much higher
TK values (e.g.,
TK = 293 K, CTAB;
TK = 284 K, CTAC [
28]). This also means that aqueous solutions of alkylguanidinium surfactants are more difficult to handle, especially in the case of CGC, where solvent evaporation may pose a serious problem.
In the next series of experiments, surface tension of surfactant solutions was determined at a constant temperature higher than the corresponding
TK value. The results are presented in
Figure 3. The appearance of a shallow minimum close to the CMC, especially for DGC, might be interpreted as being the result of some surface-active impurities in the solutions, although the uncertainty of the surface tension measurements should not be forgotten here. Note that the high purity of the surfactant samples was confirmed by the NMR and FT-IR studies.
Figure 3.
Plots of the surface tension against logarithm of the bulk molality for aqueous solutions of guanidinium cationic surfactants: (a) DGC at 298 K; (b) DDGC at 298 K; (c) TDGC at 306 K; and (d) CGC at 317 K.
Figure 3.
Plots of the surface tension against logarithm of the bulk molality for aqueous solutions of guanidinium cationic surfactants: (a) DGC at 298 K; (b) DDGC at 298 K; (c) TDGC at 306 K; and (d) CGC at 317 K.
When analyzing
Figure 3, the emphasis should be placed on the evaluation of the surfactant efficiency in lowering the surface tension of water. For example, DDGC decreases the surface tension down to about 24 ± 2 mN·m
−1 above the CMC at 298 K, whereas dodecyltrimetylammonium chloride (DTAC) shows much lower efficiency (
i.e., 42 mN·m
−1 at 298 K [
29]). This result is in agreement with those reported by other authors [
7,
19]. The propensity of guanidinium head-group to form hydrogen bonds may be considered to account for this reinforced surface activity of guanidinium cationics compared with their alkyltrimethylammonium homologues. The collective action of intermolecular hydrogen bonds makes water a highly structured liquid even at room temperature [
3]. When the surfactant units are adsorbed at the water-air interface in an oriented fashion, new hydrogen bonds between water molecules and guanidinium head-groups may form easily, thereby preventing water molecules from binding as tightly to one another and thus lowering the tension of the surface. In the case of TDGC and CGC, the temperature dependence of the surface tension should be additionally taken into account, since water loses some of its peculiar structure properties at higher temperatures [
30].
From plots of the surface tension
versus the logarithm of a solution molality, it is also possible to evaluate the critical micelle concentration (CMC) or the area (
Amin) per adsorbed surfactant unit at the solution–air interface based on the Gibbs equation [
2]. The following values have been obtained in the present work: CMC = 21 mmol·kg
−1 and
Amin = 1.4 nm
2, DGC; CMC = 5.3 mmol·kg
−1 and
Amin = 0.8 nm
2, DDGC; CMC = 1.8 mmol·kg
−1 and
Amin = 0.8 nm
2, TDGC; CMC = 0.5 mmol·kg
−1 and
Amin = 0.6 nm
2, CGC. In comparison with alkyltrimethylammonium chlorides containing the same alkyl chains and studied at the same temperature (
i.e., C
10 and C
12) [
31], the guanidinium-type surfactants self-assemble at relatively lower CMC values. On the other hand, the alkylguanidinium cations occupy more surface area per molecule at the solution-air interface. Different shapes of the polar head-groups should be taken into account in rationalizing this last observation. Indeed, the ammonium group can be described as a nearly spherical bowl, thus leading to a C
3v symmetry of the long-chain-substituted ammonium surfactants [
32]. On the contrary, the guanidinium group has a planar shape with a C
3h symmetry [
33]. The present results show that the lengthening of the alkyl chain by two methylene groups causes a steady decrease in the
Amin value, which thereby points out the increasingly enhanced adhesion of alkylguanidinium units within the adsorbed monolayer. This trend is at variance with the very high compactness of the surfactant monolayer already postulated by Miyake
et al. for DDGC based on the
Amin value of 0.37 nm
2 [
7]. Even though the small number of points in the molality region close to the break points in
Figure 3 makes the resulting CMC and
Amin values less precise here, there are no strong indications for unusually close packing of alkylguanidinium cations at the solution–air interface.
Figure 4.
Plots of the specific conductance against bulk concentration for aqueous solutions of guanidinium cationic surfactants: (a) DGC at 298 K; (b) DDGC at 298 K; and (c) TDGC at 306 K.
Figure 4.
Plots of the specific conductance against bulk concentration for aqueous solutions of guanidinium cationic surfactants: (a) DGC at 298 K; (b) DDGC at 298 K; and (c) TDGC at 306 K.
More reliable values of micellization parameters for guanidinium surfactants have been inferred from the measurements of the specific conductance of aqueous solutions. For comparative purposes, the conductivity measurements were carried above the Krafft temperature: at 298 K, DGC and DDGC, and at 306 for TDGC. Taking into account the high Krafft temperature of the C
16-homologue and given the difficulty in carrying out precise experiments at too high temperatures, further estimation of the main micellization parameters was restricted only to the first three members of the surfactant series. The conductivity
vs. concentration plots are given in
Figure 4 and the resulting values of CMC and counter-ion binding, β, to the micelle have been collected in
Table 2.
Table 2.
Critical micelle concentration, CMC, and degree of counter-ion binding to the micelle, β, for guanidinium cationic surfactants.
Table 2.
Critical micelle concentration, CMC, and degree of counter-ion binding to the micelle, β, for guanidinium cationic surfactants.
Surfactant | Temperature (K) | CMC (mmol·kg−1) | β | ΔmicG° (kJ·mol−1) |
---|
DGC | 298 | 26 ± 1 | 0.71 ± 0.01 | −22.3 ± 0.6 |
DDGC | 298 | 6.2 ± 0.3 | 0.74 ± 0.01 | −28.3 ± 0.9 |
TDGC | 306 | 1.8 ± 0.1 | 0.72 ± 0.01 | −33.6 ± 0.9 |
It should be noted here that the changes in the solution conductivity above the CMC observed in
Figure 2 and
Figure 4 are not very abrupt, which likely indicates that some partial Cl
− dissociation renders the micelles less effective charge carriers than the monomers. Of course, this hypothesis argues against the full counter-ion binding to the guanidinium micelles, at variance with the β values close to unity as reported previously [
7,
19]. The existence of globular micelles at smaller surfactant concentrations (still above the CMC), as revealed by the polarizing micrographs in
Figure 1, additionally corroborates the conclusion with respect to the β parameters drawn in the present work. The comparison between alkylguanidinium cationics and their alkyltrimethylammonium homologues [
7,
34,
35,
36] indicates that the former form less ionized micelles. Miyake
et al. have explained this increased counter-ion binding by a much closer packing of the DDGC units in micelles under the action of hydrogen bonds, thereby making the surface charge density higher than that of DTAC [
7].
The generally used rule for ionic surfactants [
2] that the CMC is divided by 4 on the addition of two methylene groups still holds approximately in this case. On the contrary, the β parameter seems to be much less sensitive to the alkyl chain length and temperature [
34,
35].
With the estimates of the two parameters, it was possible to determine the standard Gibbs free energy of micellization (Δ
micG°) of the surfactant under given conditions, using the following expression [
2,
37]:
where
XCMC is the mole fraction of the surfactant in the aqueous solution at the CMC. In this formulation, Δ
micG° represents the Gibbs free energy of transfer of 1 mole of the surfactant solute from the aqueous phase to the micellar pseudo-phase. The resulting Δ
micG° values have been added to
Table 2. As for all types of single-chain surfactants, the micellization of guanidinium cationics is a spontaneous phenomenon that is accompanied by a decrease in the Gibbs free energy. The Δ
micG° change involved in the transfer of a methylene unit of the hydrophobic group from an aqueous environment to the interior of the micelle at 298 K can be calculated from the two values obtained for DGC and DDGC. An increase in the length of the hydrophobic group makes Δ
micG° more negative by about 3.0 kJ per methylene group, which corresponds well to values usually reported for conventional surfactants [
2,
37,
38].
2.2. Thermal Effects of Micelle Formation in Various Aqueous Media
Isothermal titration calorimetry (ITC) can provide useful information about the outcome of molecular interactions involved in the micellization phenomenon [
39,
40]. Given the high
TK temperatures of alkylguanidinium surfactants, the ITC equipment offers a better possibility of controlling the loss of water due to evaporation. An example of thermograms recorded during dilution calorimetry measurements is given in
Figure 5a. The integration of peaks appearing in this thermogram leads to the determination of the total enthalpy changes during successive injections, Δ
injHi [
39,
40]. The resulting thermal effects may be summed up to obtain the cumulative enthalpy of dilution per mole of the surfactant, Δ
dilHcum, which can be expressed in the following manner [
39,
40]:
where
i is the number of successive injections,
is the number of moles of the surfactant solute (say component 2) injected into the calorimetric ampoule during each injection;
represents the apparent molal enthalpy of the surfactant corresponding to a given molality
,
is the molality of the stock solution. Taking into account the relationship between the apparent molal enthalpy,
, and the partial molal enthalpy,
, linear dependence of Δ
dilHcum against surfactant molality in
Figure 5b indicates that
is a constant function of the molality in the pre-micellar and post-micellar regions [
39]. Therefore, the intersection of the two linear portions provides estimate of the CMC for the surfactant in a given environment. In addition, the plot of Δ
dilHcum as a function of the injection number in
Figure 5c is also represented as two straight lines intersecting at the CMC. In accordance with Equation (2), the standard enthalpy of micelle formation per mole of surfactant, Δ
micH°, is determined easily from the difference between the slopes of the two linear regression segments [
39].
Figure 5.
Results of calorimetry measurements of the cumulative enthalpy of dilution obtained by injecting a 60 mmol·kg−1 aqueous solution of DDGC into deionized water at 298 K: (a) records of 40 successive injections of 5 μL aliquots into a 1 mL glass ampoule containing initially 800 μL of deionized water (the equilibration time between 2 injections was set at 45 min); (b) cumulative enthalpy of dilution as a function of the equilibrium surfactant molality; (c) cumulative enthalpy of dilution as a function of the injection number.
Figure 5.
Results of calorimetry measurements of the cumulative enthalpy of dilution obtained by injecting a 60 mmol·kg−1 aqueous solution of DDGC into deionized water at 298 K: (a) records of 40 successive injections of 5 μL aliquots into a 1 mL glass ampoule containing initially 800 μL of deionized water (the equilibration time between 2 injections was set at 45 min); (b) cumulative enthalpy of dilution as a function of the equilibrium surfactant molality; (c) cumulative enthalpy of dilution as a function of the injection number.
Figure 6 shows that the linearity of the Δ
dilHcum vs. injection number plots is preserved for DGC and DDGC in presence of the background NaCl electrolyte. Similar curves were obtained for all alkylguanidinium and alkyltrimethylammonium surfactants studied here. The values of CMC and Δ
micH° were inferred accordingly from such data. These parameters have been collected in
Table 3.
Figure 6.
Effect of NaCl addition on the cumulative enthalpy of dilution per mole of DGC (a) and DDGC (b) determined by Isothermal Titration Calorimetry at 298 K: single-solute solution (circles), 0.001 NaCl solution (triangles), 0.01 NaCl solution (squares), 0.1 NaCl solution (stars). The enthalpy of dilution has been plotted as a function of the injection number to illustrate the direct calculation of the enthalpy of surfactant micellization.
Figure 6.
Effect of NaCl addition on the cumulative enthalpy of dilution per mole of DGC (a) and DDGC (b) determined by Isothermal Titration Calorimetry at 298 K: single-solute solution (circles), 0.001 NaCl solution (triangles), 0.01 NaCl solution (squares), 0.1 NaCl solution (stars). The enthalpy of dilution has been plotted as a function of the injection number to illustrate the direct calculation of the enthalpy of surfactant micellization.
The discrepancies between the corresponding CMC values obtained in conductimetry and calorimetry experiments are obviously due to the differences in the sensitivity and reliability of the two methods and further data processing (e.g., the sharpness of the break in the concentration dependence of a given physical property [
41]).
With the knowledge of such thermodynamic functions as Δ
micG° (
Table 2) and Δ
micH° (
Table 3), it is possible to calculate the standard molar entropy of micellization, Δ
micS°, for DGC, DDGC, and TDGC in single-component solutions. The following values have been obtained:
T·Δ
micS° = 17.8 ± 0.9 kJ·mol
−1, DGC at 298 K;
T·Δ
micS° = 20.5 ± 2 kJ·mol
−1, DDGC at 298 K;
T·Δ
micS° = 16.1 ± 0.8 kJ·mol
−1, TDGC at 306 K. When there is less than 14 carbon atoms in the alkyl chain, the negative values of Δ
micG° are mainly due to the positive values of Δ
micS°, Δ
micH°, even negative, being smaller than the value of
T·Δ
micS°. As for other types of surfactants, micellization of guanidinium cationics is an entropy-driven process owing to the transfer of the hydrophobic tails from the water environment to the micelle core [
3,
42,
43]. Therefore, the downward trends in the CMC with lengthening the alkyl chain, increasing the salt content, and decreasing the temperature, as observed qualitatively for DTAC, TTAC, DGC, and DDGC in
Table 3, may be considered as consistent with the reinforcement of the hydrophobic effect in relation with changes in the peculiar solvent structure or in the micelle size and shape. In the case of TDGC, the enthalpy and entropy contributions to Δ
micG° are almost equal (but still opposite in sign) and the phenomenon also becomes enthalpy-driven. When analyzing the variations of Δ
micH° in
Table 3, it should always be remembered that the thermal effect of micellization is a result of the interplay among various competitive or co-operative interactions; besides the hydrophobic interaction, the outcome of forces operating among the ionized head-groups, counter-ions and water molecules constitutes the most important contribution.
Table 3.
Critical micelle concentration, CMC, and standard enthalpy of micelle formation per mole of the surfactant, ΔmicH°, for guanidinium cationic surfactants.
Table 3.
Critical micelle concentration, CMC, and standard enthalpy of micelle formation per mole of the surfactant, ΔmicH°, for guanidinium cationic surfactants.
Surfactant | Solvent | Temperature (K) | CMC (mmol·kg−1) | ΔmicH° (kJ·mol−1) |
---|
DTAC | H2O | 298 | 21.5 ± 0.1 | 5.1 ± 0.2 |
DTAC | 0.01 M NaCl | 298 | 18.1 ± 0.1 | 4.2 ± 0.1 |
DTAC | 0.1 M NaCl | 298 | 8.6 ± 0.1 | 3.6 ± 0.1 |
TTAC | H2O | 298 | 5.0 ± 0.1 | 2.2 ± 0.2 |
TTAC | 0.1 M NaCl | 298 | 3.3 ± 0.4 | 1.1 ± 0.2 |
TTAC | H2O | 318 | 6.0 ± 0.4 | −8.9 ± 0.2 |
DGC | H2O | 298 | 28.4 ± 0.1 | −4.5 ± 0.1 |
DGC | 0.001 M NaCl | 298 | 25.5 ± 0.6 | −4.2 ± 0.3 |
DGC | 0.01 M NaCl | 298 | 22.7 ± 0.2 | −4.0 ± 0.1 |
DGC | 0.1 M NaCl | 298 | 9.6 ± 0.3 | −6.4 ± 0.1 |
DDGC | H2O | 298 | 6.2 ± 0.2 | −7.8 ± 0.5 |
DDGC | 0.001 M NaCl | 298 | 5.7 ± 0.2 | −9.0 ± 0.1 |
DDGC | 0.01 M NaCl | 298 | 3.5 ± 0.1 | −9.5 ± 0.5 |
DDGC | H2O | 306 | 6.6 ± 0.4 | −12.3 ± 0.9 |
DDGC | H2O | 310 | 7.3 ± 0.4 | −13.0 ± 0.4 |
TDGC | H2O | 306 | 1.5 ± 0.3 | −17.5 ± 0.3 |
By far the most important conclusion drawn from the analysis of
Table 3 is that the micellization of alkylguanidinium chlorides is an exothermic phenomenon, irrespective of the experimental conditions used. From this point of view, this micellization behavior is at variance with the endothermic micelle formation observed here for their alkyltrimethylammonium homologues. A more detailed comparison with the thermodynamic parameters reported in the literature for various ionic surfactant indicates that this behavior is rather similar to that exhibited by alkyltrimethylammonium bromides with the same hydrophobic tails, even though the micellization of the guanidinium cationics studied occurs at lower CMC values [
44]. The latter conclusion may be easily rationalized since the degrees of counter-ion binding to the micelle are similar in both cases, and Cl
− is known to be more effective than Br
− in salting out the surfactant cation and thus depressing the CMC [
2]. In accordance with some previously reported results [
35,
36], a passage from endothermic to exothermic micellization for alkyltrimethylammonium chlorides is favored with increasing the temperature and NaCl addition to the aqueous phase (see the results reported for DTAC and TTAC in
Table 3). From the viewpoint of enthalpy, the present alkylguanidinium surfactants thus behave as their alkyltrimethylammonium homologues at higher temperatures and salt contents. This clearly points towards the guanidinium head-group as being responsible for the particularly exothermic character of micellization.
When the DDGC units self-assemble to form micelles in the presence of background electrolyte, only the first addition of NaCl to the aqueous phase makes Δ
micH° much less negative, and the thermal effect then seems to level off. This likely means that the DDGC micelles are fairly compact already in single-component solutions, but not to the degree proposed by Miyake
et al. [
7]. They then attain their optimal shape and size (e.g., maximum aggregation number) at low ionic strengths. For the same tail lengths, the fractions of surfactant units in the micellized state should be thus much greater than those characterizing the cationic micelles of alkyltrimethylammonium chlorides. For DGC in the presence of NaCl content up to 0.01 mmol·kg
−1, a small decrease in the CMC is paralleled by a near constant value of Δ
micH°. Here, the salt addition to the aqueous phase does not cause much change in the self-aggregation behavior of the guanidinium type surfactants. High NaCl contents (of the order of 0.1 mmol·kg
−1) render the phenomenon again more exothermic with the concomitant decrease in the CMC. Since the pre-micellar and post-micellar parts of the Δ
dilHcum vs. injection number plots preserve their linear character (see
Figure 6a), it may be argued that the DGC micelles formed at high ionic strengths have a different shape from what is observed in single-solute solutions or at lower NaCl contents.
A stronger counter-ion binding in the Stern layer and hydrogen bonding with water molecules are also to be taken into account. Stronger interactions in the polar mantle of the micelle likely account for the enthalpy gain when the cationic head-groups are transferred from the aqueous environment to the micelle mantle. The increased magnitude of the exothermic micellization for longer hydrophobic tails may still be explained on the same basis when postulating the concomitant increase in the micelle aggregation number.