J. Phys. Chem. 1986, 90, 5870-5872
5870
eV for an indirect transition, and 1.24 eV for a direct transition are close to the reported single value of 1.21 eV (again, the mode was not assigned). The occurrence of both an indirect and direct transition is common for the CdIn2V14family of materials, and the difference between Ed and Ed has been reported to be 0.1 eV for CdIn2Te4.12 Our value of 0.08 eV is close to the 0.1 eV value. We also obtained a value of 0.16 eV for the difference between Eind and Ed for p-HgIn2Te4. It seems that the difference between Eindand Ed decreases from 0.22 eV for CdInzSe4to 0.08 eV for CdInzTe4” and increases from 0.08 eV for CdIn2Te4to 0.16 eV for HgIn2Te4. The Mott-Schottky plot intercept value of -0.48 V was close to that of the onset of photocurrent (-0.53 V, see Table I); therefore, the open circuit photovoltage is expected to be -0.6 V (Ercdox = -1.09 V). Considering a band gap value of 1.20 eV (0.9 eV for HgIn2Te4)and an analysis parallel to that given for p-HgIn2Te4above, the valence band would be located at --0.4 V (-0.6 V for HgIn2Te4) and the conduction band near -1.6 V (-1.5 V for HgInzTe4). The photocurrent onset and the maximum power efficiency in [Cr(III)EDTA]- with monochromatic light irradiation were close to that in Fe(II1)TEA complex solution, but the maximum power efficiency in [Cr(III)EDTA]- with polychromatic light irradiation was less than that in Fe(II1)TEA complex solution (the reason for this is given above). Therefore, the Fe(II1)TEA complex solution is a preferred one for both p-HgIn2Te4and p-CdInzTe4 electrodes. Considerable anodic dark current was observed at the potential more positive than -0.7 V, similar to that for p-HgIn2Te4in a polysulfide solution, but there was nearly no anodic dark current over this potential region in the blank solution for polysulfide. The explanation for this result would be similar to that given above for HgIn2Te4. For n-type CdIn2Te4 (the above discussion is for p-type), photoanodic current was found in a polysulfide couple solution (0.1 M S-1 M NazS-l M NaOH) under irradiation. With 600-nm monochromatic light (3.0 mW/cm2), the quantum efficiency of carrier collection at short circuit (-0.8 V) was 61%, which increased to almost 100% by -0.4 V. However, with multiple scans, photocorrosion was noted. No photocorrosion was found for ptype CdIn2Te4after PEC study in a polysulfide redox solution (2 M S-2 M Na2S-2 M KOH) in the cathodic region. It is obvious that the photocorrosion problem of the p-type material is less severe than that of the n-type material. In the iron hexaaquo
couple solution, the photocurrent onset was 0 V, and thus the open circuit photovoltage is predicted to be 0.4 V (Erdox= 0.4 V). The polychromatic power efficiency was 2%, higher than that in the polysulfide couple (1%). However, the photocorrosion problem also exists in the iron hexaaquo couple.
Conclusions We have successfully used two very different pH redox couples both with significantly negative redox potentials with p-type semiconductors: CrEDTA complex, pH 4.5, Eredox = -1.24 V, and FeTEA complex, pH 13, Erdox= -1.09 V. The photoelectrochemical behavior of p-HgIn2Te4 and pCdIn2Te4have been explored for the first time. The same is true for n-CdInzTe4 but to a much lesser extent. For p-HgIn2Te4,the quantum efficiencies of carrier collection are 590% even at short circuit in either redox couple. Monochromatic and polychromatic power efficiencies based on three electrode cell experments up to -10% and -3%, respectively, have been obtained. The flat-band potential is -0.70 V in either couple indicating no preferential absorption of H+ or OH- on the semiconductor surface. In the FeTEA complex couple, it appears that the semiconductor approaches 100% stability to photocorrosion. This semiconductor particularly with some increase in the fill factor would appear to be a useful small band gap semiconductor. For p-CdIn2Te4, the quantum efficiencies of carrier collection are 590% even at short circuit in either redox couple. Monochromatic and polychromatic power efficiencies up to 11% and 2%, respectively, have been obtained. It is clear that with an increase of the relatively low fill factor (0.24-0.3), the power efficiencies would show significant improvement. The flat band was near -0.55 V and largely independent of the nature of the two couples considered. Given the fact that there appears to be a direct gap transition at 1.24 eV, which is near the ideal desired value of 1.4 eV, and the foregoing characteristics, this semiconductor looks like a very promising one for photovoltaic cells whether liquid or solid junction. Acknowledgment. This work was supported by a grant from Phillips Petroleum Co. Registry No. [Cr(III)EDTA]-, 73690-75-2;CdIn2Te,, 12014-06-1; HgIn2Te4, 12383-65-2; HgTe, 12068-90-5; In2Te3, 1312-45-4; Fe(H20):+, 15365-81-8;Fe(H20)2+,15377-81-8;Na2S, 1313-82-2;KOH, 1310-58-3; NaOH, 1310-73-2; Cd, 7440-43-9; In, 7440-74-6; Te, 13494-80-9;Ag, 7440-22-4; S,7704-34-9.
Formamide, a Water Substitute. 12. Krafft Temperature and Micelle Formation of Ionic Surfactants in Formamide‘ I. Rico* and A. Lattes Laboratoire des IMRCP, UA CNRS No. 470, Universite Paul Sabatier, 31062 Toulouse Cedex. France (Received: March 10, 1986) Critical micelle concentrations (cmc) were determined from surface tension measurements for cetyltrimethylammoniumbromide (CTAB) and sodium dodecyl sulfate (SDS) at 60 OC in formamide. The cmc were considerably higher in formamide than in water, and the corresponding Krafft temperatures for these surfactants were also higher in formamide than in water. These results can be explained by the structure of formamide, which can be considered as a low-melting anhydrous fused salt.
Introduction ~ ~aggregation l in ~ nonaqueous ~ solvents l have ~ received relatively little attention compared to the number of investigations using water.
Since formamide has a relative permittivity of 109.5 and a surface tension of 58.5 “em-’ at 25 ‘c, it has sufficient cohesive ~ force to favor mOlecUlar aggregation. We recently demonstrated that microemulsions (defined as transparent dispersions of water, oil, and amphiphile)2 can be prepared using formamide instead
(1) Part XI: Escoula, B.; Rico, I.; Lattes, A. Tetrahedron Lett. 1986, 27(13), 1499.
(2) (a) Daniellson, I.; Lindmann, B. Colloids Surf.1981, 3, 381. (b) Friberg, S. E. Colloids Surf. 1982, 4 , 201.
0022-3654/86/2090-5870$01.50/0
0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 22, 1986 5871
Micelle Formation of Ionic Surfactants in Formamide
TS
m d .m . ' )
50
40
, 30 10-2
Figure
31
10-l
azz
1 IbJE (wlu
I-')
i. Surface tension (7)plotted against log Iconcentration1for SDS.
of water. Nonionic amphiphiles ( p l u r ~ n i c s )as~ well as ionic amphiphiles like cetyltrimethylammonium bromide (CTAB) or phosphonium salts (C6H5)3P+CfiwlX-.(X = Br, I)&' have been used as surfactants. The phosphonium salts with large polar heads are, however, insoluble in water and cannot be used to prepare aqueous microemulsions under normal experimental conditions. Contradictory results have been reported on the production of micelles in formamide using ionic surfactants.*-12 We therefore carried out a series of investigations using formamide as solvent. We demonstrated molecular aggregation in formamide of large ions which are insoluble in water.13 The long-chain phosphonium salts (C6H5)3P+C,H2,+11- and their fluorinated homologues (C6H5)3P+CH2CH2RFI-are insoluble in water even at high temperatures. However, they are soluble in formamide, and micellization was detected from measurements of surface tension and conductance. The Krafft temperatures in formamide for such derivatives are nevertheless quite high (around 60 "C), and the surface tension and conductance measurements were therefore carried out at 64 OC. There are no reports in the literature of measurements of Krafft temperatures for classical ionic surfactants (water-soluble surfactants) in formamide. Micellization only takes place above the Krafft temperature, which itself depends on the nature of the surfactant and the medium. The aim of this study was to determine the Krafft temperatures in formamide for two classical ionic surfactants: an anionic surfactant, sodium dodecyl sulfate (SDS), and a cationic surfactant, cetyltrimethylammonium bromide (CTAB). We also investigated the surfactant action of these compounds in formamide above the Krafft temperature. (3) Gautier, M.; Rico, I.; Ahmad-Zadeh Samii, A.; de Savignac, A.; Lattes, A. J . Colloid Interface Sci., in press. (4) (a) Rim, I.; Lattes, A.; Nouu. J. Chim. 19848 8(7), 429. (b) Rim, I.; Lattes, A. J. Colloid Interface Sci. 1984, 102, 285. (5) Rim, I.; Lattes, A. Surfactants in Solution; Ed. Mittal, K. L., Ed.; in press. (6) Rim, I.; Lattes, A. Proceedings of the 5th International Conference on Surface and Colloid Science; Rosano, H., Ed.: in press. (7) Ahmad-Zadeh Samii, A.; de Savignac, A.; Rico, I.; Lattes, A. Tetrahedron 1985,41, 3683. (8) Gopal, R.; Singh, J. R. J . Phys. Chem. 1973, 71, 554. (9) Gopal, R.; Singh, J. R. Kolloid Z.Z.Po1ym. 1970, 259, 699. (10) Singh, H. N.; Saleem, S. M.; Singh, R. P.; Birdi, K. S. J . Phys. Chem.
1980, 84, 219. (1 1) Couper, A.; Gladden, G. P.;Ingram, B. Faraday Discuss. Chem. Soc. 1975, 59, 63. (12) Almgren, M.; Swarup, S.; Lofroth, J. E. J. Phys. Chem. 1985, 89, 4621. (13) Esmula, B.; Hajjaji, N.; Rico, I.; Lattes, A. J . Chem. SOC.,Chem. Commun. 1984, 1233.
Figure 2. Surface tension (y) plotted against log )concentration1 for CTAB. TABLE I: Krafft Temperatures of SDS and CTAB in Formamide and Water Krafft temp, "C water formamide
SDS 16 55
CTAB 26 43
Experimental Section Materials. Sodium dodecyl sulfate (SDS) and cetyltrimethylammonium bromide (CTAB) (Merck, 99% min pur), formamide (Merck, 99.5% min pur) were used a supplied. Surface tension measurements were carried out at 60 OC by using a Prolabo tensiometer (Tensimat no. 3) by using the ring detachment method. Kraflt temperatures were determined by visual observation of test tubes containing 10 mL of formamide and varying amounts of surfactant. The formamide-surfactant mixtures were slowly heated with constant agitation, and the temperature at which all surfactant completely dissolved was recorded. The solutions under continuous vigorous agitation (surfusion can occur if the mixture is not shaken vigorously) were than slowly cooled, and the temperature at which the surfactant separated from the solution was also recorded. Results The change in surfactant solubility in formamide with temperature was found to be similar to that of such surfactants in water. The solubility becomes practically exponential at Krafft temperatures for CTAB and SDS of 43 and 55 OC, respectively. Surface tensions were therefore measured at 60 OC as a function of CTAB and SDS concentrations in formamide. The surface tension of formamide at 60 O C is 53 mN-m-' and after addition of surfactant drops to a final value of about 36-38 mN-m-', which does not change with further addition of surfactant. The cmc can be determined from plots of y vs. log concentration (Figures 1 and 2). We found values of 9 X mo1.L-' for CTAB and 2.2 X lo-' mo1.L-l for SDS. Discussion The Krafft temperatures of these surfactants are found t o be much higher in formamide than in water (Table I). Evans et al.14 found high Krafft temperatures for alkyltrimethylammonium bromides in ethylammonium nitrate (48 O C (14) Evans, D. F.; Yamaguchi, A.; Roman, R.; Casassa, E. Z . J . Colloid Interface Sci. 1982, 88, 89.
5812 The Journal of Physical Chemistry, Vol. 90, No. 22, 1986
for CTAB). Such high values can be explained in terms of the structure of the medium. The Krafft temperature can be defined as the temperature of fusion of solvated surfactant. For ethylammonium nitrate, the structure of the solvated surfactant is probably much more rigid than that of the same hydrated surfactant. In the former case, a type of mixed salt with electrostatic interactions is formed. The structure of the hydrated surfactant, on the other hand, is only held together by dipole type interactions, producing a more fragile entity. A similar explanation can be invoked for formamide. Many investigations using NMR, IR, Raman spectroscopy, etc.15-18have demonstrated an almostly totally ionic structure for formamide in the liquid state. It can be regarded as a planar molecule with a substantial contribution from a resonance structure (XI): H \N-c/H
+O
H'
I
-
t
>N=C
H,
'0-
H
I1
The C-N bond has considerable double bond character,I9 and the Arrhenius energy of activation for internal rotation has been calculated to be 75-79 kJ.mol-1.20,21 It has also been shown that solvation of ions by formamide involves the oxygen atom for cations and the nitrogen atom for anions, Le., mainly electrostatic interactions.22 The structures of liquid formamide and solvated salts in formamide therefore can be thought similar to the fused salt structures of compounds like ethylammonium nitrate. Electrolytes solvated by formamide can be thought of as mixed salts with higher fusion points than their hydrated homologues. The Krafft temperatures of ionic surfactants are therefore much higher in formamide than in water. (15) Costrain, C. C.; Dowling, J. M. J . Chem. Phys. 1960, 32, 158. Kurland, R. J.; Wilson, E. B. J . Chem. Phys. 1957, 27, 5 8 5 . Hirota, E.; Sugisaki, R.; Nielsen, C. J.; Sorerensen, G. 0. J . Mol. Spectrosc. 1974, 49, 251. (1 8) (a) Gardiner, D. J.; Lees,A. J.; Straughan, B. P. J. Mol. Struct. 1979, 53, 15. (b) Lees, A. J.; Straughan, B. P.; Gardiner, D. J. J. Mol. Struct. 1981, 71. - , 61.
(19) Drakenberg, T.; Forsen, S. J. Phys. Chem. 1970, 74, 1. (20) Summers, B.; Piette, L. H.; Schnieder, W. G. Can. J . Chem. 1960, 38, 681. (21) Kamei, H. Bull. Chem. SOC.Jpn. 1968, 41, 2269. (22) Lees, A. J..: Straughan, B. P.; Gardiner, D. J. J . Mol. Struct. 1979, 54, 37.
Rico and Lattes The cmc in formamide for both SDS and CTAB are much higher in formamide than in water (almost 100-fold). Couper et ai." and Ray23,24have obtained similar results for nonionic surfactants with similar differences between the two solvents. Taken together with our results, this would suggest that solvophobic interactions are weaker in formamide than in water, which would explain the higher cmc values in formamide. Similar increases in cmc have been observed between water and ethylammonium nitrate.14 We also compared our results with those of similar studies reported on the production of micelles in formamide using ionic surfactants.s-12 Some years ago, Couper et al." described molecular aggregation of some ionic surface-active agents (SDS and dodecyltrimethylammonium bromide) in formamide. In contrast to the nonionic surfactants, they found that adsorbed films of these ionic compounds did not display the discontinuity that is characteristic of micelle formation. This has recently been confirmed by Almgren12for SDS. These results are in agreement with our findings, since both Couper and Almgren carried out their investigations below the Krafft temperature (25 and 40 "C for the compounds studied by Couper; 20 "C for SDS). Gopal and Singh8s9and Singh et al.IOhave claimed that direct micelles for SDS and CTAB are formed in formamide at 20 "C and at low concentrations. The cmc values were obtained from conductance measurements. However, Almgren has pointed out that observation of a break point in conductance concentration plots is not necessarily an indication of micelle formation. The recent work of Almgren on self-diffusion, surface tension, and fluorescence quenching has demonstrated that SDS doe not form micelles in formamide at 20 OC.I2 In conclusion, in order to produce micelles of ionic surfactants in formamide one must work above the Krafft temperature. Under these conditions, the cmc values are higher in formamide than in water (for both ionic and nonionic surfactant^).']^^^^^^ Formamide would seem to confer properties of a low-melting anhydrous salt similar to ethylammonium nitrate.
Acknowledgment. We thank Greco Microemulsions and the U.S. Army for financial support. Registry No. SDS, 151-21-3; CTAB, 57-09-0; formamide, 75-12-7. (23) Ray, A. J . Am. Chem. SOC.1969, 91, 651. (24) Ray, A. Nature (London) 1971, 231, 313.
