1342
Jun'ichiro Muto
of cation contents, and of t h e presence of C02. These correlations suggest that the cracking, which involves acid sites, is enhanced by the acidic OH groups formed on COz adsorption. As usual for acid sites, these centers partly loose their activity upon aging. , In conclusion, this study indicates t h a t acidic OH groups can be created on CO2 adsorption a t t h e expense of basic hydroxyls groups in magnesium and calcium Y faujasites. T h e increase in cracking activity relates to the presence of these acidic OH groups. T h e reversibility of C02 action at 465 OC explains that a certain COz partial pressure is required to raise the catalytic activity.16 T h e improvement of catalytic properties by other acidic reactants (SOz, CS2,. . .)I5 might perhaps be interpreted on a similar basis. Also, zeolites exchanged with transition metal cations might exhibit similar activity enhancements. Acknowledgment. T h e authors thank the Laboratory of Chemical Analysis and Dr R. Beaumont for the gift of the Ca samples. This work was assisted by the DGRST (Contract No. 74.7.1 157).
References and Notes (1) A. Bielanski and J. Datka, J. Catal.. 32, 183 (1974); T. Yashima. H. Suzuki, and N. b r a , ibid., 33,486 (1974). (2) J. B. Peri, J. Phys. Chem., 79, 1582 (1975). (3) P. Jacobs and J. B. Uytterhoeven, J. Catal., 22, 193 (1971). (4) R . Beaumont, P. Pichat, D. Barthomeuf, and Y. Trambouze, Catal., Proc. lnt. Congr. 5th. 1972, 343 (1973). (5) J. B. Peri, Catal., Proc. lnt. Congr. Sth, 1972, 329 (1973). (6) P. A. Jacobs and J. B. Uytterhoeven, J. Chem. SOC., faraday Trans. 1, 69, 373 (1973). (7) J. Ward, J. Phys. Chem., 72, 2689, 4211 (1968).
(8) J. B. Uyttwbeven, R. Schoonheydt. B. V. Liengme. and W. K. Hall, J. Catal., 13, 425 (1969). (9) P. A. Jacobs, F. H. Van Cauwelaert. E. F. Vansant, and J. B. Uytterhoeven, J. Chem. Soc., 69, 1056 (1973). (IO) J. Scherzer and J. L. Bass, J. Phys. Chem., 79, 1200 (1975). (11) J. Ward, Adv. Chem. Ser., No. 101, 380 (1971). (12) See D. W. Breck, "Zeolite Molecular Sieves", Wlley, New Ywk, N.Y., 1974, p 484. (13) J. Scherzer and J. Bass, J. Catal.. 28, 101 (1973). (14) J. Ward, J. Catal., 9, 225 (1967). (15) Kh. M. Minachev and Ya. I. Isakov, Adv. Chem. Ser., No. 121,451 (1973). (16) Kh. M. Minachev. G. V. Isagulyants, Ya. I. Isakov. N. Ya. Usachev, and N. N. Rozhdestvenskaya, lzv. Akad. Nsuk SSSR,Ser. Khim., I, 42 (1974). (17) C. L. Angel1 and M. V. Howell, Can. J. Chem., 47, 3831 (1969). (18) M. V. Mathieu and P. Pichat, in "La catalyse au laboratoire et dans I'industrie", B. Claudel, Ed.. Masson et Cie. Paris, 1967, p 319. (19) J. Ward, J. Catal., 26, 451 (1972). (20) P. Pichat. R. Beaumont, and D. Barthomeuf. C. R. Acad. Sci., 272, 612 (1971); J. Chem. SOC., faraday Trans. 1, 70, 1402 (1974). (21) P. Gallezot, R. Beaumont, and D. Barthomeuf, J. Phys. Chem., 78, 1550 (1974). (22) C. V. McDaniel and P. K. Maher in "Molecular Sieves", Society Chemical Industry, London, 1968, p 186. (23) H. Bremer. W. M e , R. SchoM. and F. Vogt, Adv. Chem. Ser., No. 121, 249 (1973). (24) P. J. Anderson, R. F. Horlock, and J. F. Oliver, Trans. faraday SOC.,61, 2754 (1965). (25) L. H. Little, "Infrared Spectra of Adsorbed Species", Academic Press, London, 1966, pp 74-89. and references therein. (26) T. V. Evans and T. L. Whateley, Trans. faraday Soc., 63, 2769 (1967). (27) S.J. Gregg and J. D.Ramsay, J. Chem. SOC.A, 2784 (1970). (28) Y. Fukuda and K. Tanabe, Bull. Chem. Soc., Jpn., 46, 1616 (1973). (29) P. Pichat, J. Veron, B. Claudel. and M. V. Mathieu, J . Chim. Phys., 63, 1026 (1966); P. Pichat and G. Brau, ibid., 66, 724 (1969). (30) V. G. Amerikov and L. A. Kasatkina, Kinet., Katal.. 12, 165 (1971). (31) N. D. Parkyns, J. Phys. Chem.. 75, 526(1971). (32) C. L. Angel1 and P. C. Schaffer, J.phys. Chem., 69,3463 (1965); C. L. Angel1 and M. V. Howell, ibid., 74, 2737 (1970). (33) T. lizuka, H. Hattori. Y. Ohno, J. Sohma, and K. Tanabe, J. Catal., 22, 130 (1971). Faraday Trans. 1,69, (34) P. A. Jacobs and J. B. UytteWven. J. Chem. Soc., 359 (1973). (35) D. Barthomeuf and R. Beaumont, J. Catal.. 30,288 (1973).
Dimeric Properties of Rhodamine B in Glycerol, Ethylene Glycol, and Acetic Acid Jun'ichiro Muto Department of instrumentation Engineering, faculty of Engineering, Keio University, 832 Hiyoshi, Kohoku, Yokohama 223, Japan (Received October 6, 1975) Publication costs assisted by Keio University
Absorption spectra of rhodamine B are examined in solutions of glycerol, ethylene glycol, and acetic acid. Particular attention is given to the absorption spectra of monomers and dimers of rhodamine B in these solutions. Furthermore, dimeric structures of rhodamine B in these solutions are also investigated.
1. Introduction
Rhodamine B, one of the most widely used dye laser materials, has been found to lase in alcohol, water, and PMMA.'-4 T h e appearance of the rhodamine B organic dye laser in the various media mentioned above has stimulated a further study on spectroscopic and structural properties of this dye in relation to t h e lasing mechani~rn."'~ In aqueous solution, rhodamine B has a tendency to aggregate and form dimers with increasing concentration. Furthermore, the fluorescence quantum efficiency of aqueous rhodamine I3 was found to be greatly affected by dimerization, The Journal of Physical Chemistry, Vol. 80,No. 12, 1976
since dimers of rhodamine B in water were recognized to make little contribution t o fluorescence, though they were capable of optical a b ~ o r p t i o n . ~ In J ~ 'glycerol, however, rhodamine B was observed to have a fluorescence quantum efficiency of nearly unity.12 T h e absorption spectrum of rhodamine B is greatly influenced by addition of acid. On the other hand, the absorption spectrum of rhodamine B in acetic acid is quite similar to that in nonacidic solution such as alcohol or acetone.'O I n addition to the use for dye lasers as mentioned before, rhodamine B in ethylene glycol is commonly used as a quantum counter.15
Absorptlon Spectra
1343
of Rhodamine B In Various Media
Under these circumstances, it will be quite important to examine absorption spectra of rhodamine B in glycerol, ethylene glycol, and acetic acid in some detail with emphasis on determining the monomeric and dimeric absorption spectra in these solutions. Furthermore, dimeric configurations of rhodamine B in these solutions are also investigated. 2. Experimental Section Rhodamine B, used in our experiment, was obtained commercially (reagent grade, Tokyo Kasei Industries) and no further purification was done in our laboratory. TOeliminate the efficient light from the dye and to select only the transmitted light from the sample solution, absorption measurements a t various concentrations and solutions are done by using two monochromators. Details of the experimental apparatus are found in the literature.I0
3. Results and Discussions Absorption coefficients in the photon energy region between 2.1 and 2.5 eV are measured at room temperature in solutions of glycerol and ethylene glycol. The obtained results are shown in Figures l a and l b . As seen from the figures, the absorption spectra are somewhat different a t various concentrations and a t various solutions. However, the existence of a main absorption peak around 2.15 eV and a broad shoulder a t about 2.35 eV are clearly recognized a t all concentrations in the solutions examined. Similar spectral behaviors have already been reported in case of acetic acidic-rhodamine B.l0 Monomer and Dimer Spectra. In this section, we shall determine the absorption spectra of the monomer and dimer separately from the observed total absorption spectra, and then calculate the equilibrium constant of a monomer-dimer equilibrium state. First, we write for the relations of various absorption coefficients for the monomer-dimer equilibrium
a ( E )= a m ( E ) X
+ Lud(E)(1 - X )
(1)
where a ( E ) ,a,(E), and ad@) are the observed total molar absorption coefficient and those of monomers and dimers, respectively, a t photon energy E, and x denotes the fraction of monomer. From the law of mass action for the monomer-dimer equilibrium, we have
K = 2Cr2/(1 - X )
(2)
in which K expresses the equilibrium constant for the process [monomer monomer = dimer], and C, the total concentration. The free energy I G for the dissociation of the dimer is written as
+
AG = -RT In K
(3)
where R denotes the gas constant and T the absolute temperature. Now, integrating eq 1, we have
+
S a ( E ) d E = I ~1x 2
(4)
where I1 = S [ a m ( E )- a d @ ) ] d E , and 1 2 = S a d ( E )dE. In eq 4, the integrated molar absorption coefficient, I a ( E ) d E , increases with increasing x , if I1 is positive, while it decreases with increasing x , if I 1 is negative. As is found in eq 2, x (fraction of monomer) increases with a decrease of C (total dye concentration), since K (equilibrium constant) is positive and x is positive and less than unity.
A " B
21 21
C D
2.3 2.5 2.3 2.5 21 2.3 2.5 2s 2.3 2.5 Photon Energy (el!)
0" A
B C D
"
21
'
"
'
'
2.3 25 21 2.3 2.5 2.1 23 25 21 23 25 Photon Energy (eV)
Figure 1. Molar absorption spectra of rhodamine B in (a) glycerol (b) ethylene glycol. A indicates absorption spectrum at M, 6 at 5 X lod5 M, C a t M, and D at M.
Therefore, the integrated molar absorption coefficient in eq 2 decreases with an increase of concentration when I1 is positive and decreases when 11 is negative. Now, let us examine the results of the experiments. The concentration dependence of the integrated absorption coefficients in various solutions is shown in Figure 2. The results cf glycerol- and ethylene glycol-rhodamine B are graphically obtained from Figure 1, and those of water and acetic acid from previously reported results.1° In case of glycerol, ethylene glycol, and acetic acid, the obtained integrated absorption coefficients show close resemblance with each other in that they are found to increase with their increasing concentrations in the lower concentration range M), while they decrease in the higher concentration range (>lov4M). Taking into account the above mentioned discussions on the concentration dependence of integrated absorption coefficients, and also our experimental results described above, we conclude that the monomer-dimer equilibria are realized in these solutions at lower concentrations (S10-4 M). The decrease of integrated absorption coefficients at the The Journal of Physical Chemistry, Vol. 80, No. 12, 1976
1344
Jun'ichiro Muto
2
.r!
.-
L
x104 6p
"
'
" " '
'
' " ' I
"
. I -
-7 1
_.e.-
22
0
I
,
,
,
,
,
,
1o
,
,
,
-~
,
, , , , , , , ,
,
1o-4
1o
-~
Photon Energy (eY)
(mol/\) Figure 2. Concentration dependence of integrated molar absorption coefficient of rhodamine B in (GI) glycerol, (Ac)acetic acid, (Eg)ethylene glycol, and (Aq) water. Concentration
higher concentrations (>IO-* M) might be attributed t o t h e formation of trimers and/or higher polymers. In contrast with the experimental results of glycerol, ethylene glycol, and acetic acid, the integrated molar absorption coefficients of aqueous rhodamine B show slight changes with decreasing concentration (see Figure 2). Since t h e monomer-dimer equilibrium of aqueous rhodamine B is shown t o exist in t h e range of IOh5t o lo-" M,99'0J6 t h e area of t h e monomer spectrum becomes the same as the area of the dimer spectrum; that is, the two joined ions of rhodamine B in water have an absorption equal to that of the separate ions. Similar spectral behavior has been observed in case of aqueous Nafluorescein." Next, we shall calculate the monomer and dimer spectra of , 'First, ~ we assume rhodamine B in the following ~ a y : ~(1) various values of K in eq 2 and calculate the corresponding x for a concentration c. (2) Using these values, a,,,(E) and CYd(E) are obtained from eq 1, by the method of a least-squares fit. (3) Once we determine the best-fit monomer and dimer spectra for a given K , then we calculate the average standard deviations between the observed a ( E )and t h e best-fit a ( E ) for the data a t all concentrations. (4)Finally, we determine and &J(E)on t h e basis of t h e t h e best values of K , CY,,,(E), minimum average standard deviation. As discussed previously, t h e monomer-dimer equilibria of rhodamine B in glycerol, ethylene glycol, and acetic acid seem M). T h e n , the to exist a t t h e lower concentrations best-fit spectra of monomers and dimers in these solutions are calculated using the experimental results obtained in t h e above mentioned concentration range. In Figures 3a, 3b, and 3c, the absorption spectra of monomers have definite maxima a t about 2.2 eV, although details of t h e spectra are different for different solutions, whereas, dimeric absorption spectra show rather complicated structure as will be discussed below. In case of glycerol, one main band appears in the lower energy region, and a subsidiary band, in the higher energy region. In ethylene glycol, subsidiary bands of HIand Hz appear and the lower energy band is graphically found to result from t h e overlapping of the two bands with peak positions of 2.155 and 2.195 eV. While in acetic acid, two larger bands in addition to two smaller bands are clearly recognized.18 T h e obtained values of peak positions, in addition t o the dissociation constants and the free energies of rhodamine B in these solutions are presented in Table I. Structure of DirnPr. When two dye molecules come together The Journal of Physical Chemistry, Vol. 80, No. 12. 1976
x ~ ~ 5
a
2
d
P Photon Energy(eV1 x1~5
I
,
O 1 2:1 212 ' 213 ' 24 ' 2f5 Phot on Energy (eV) '
'
Figure 3. Molar absorption spectra of monomeric (solid curves) and dimeric (broken curves) rhodamine B in (a)glycerol, (b)ethylene glycol, (c) acetic acid.
to form a dimer unit, the monomer absorption peak splits into two peaks (see Figure 4):H band (higher energy band) and J band (lower energy band). T h e splitting is caused by a point dipole-dipole interaction between adjacent molecules in the dimer, T h e resulting splitting depends upon t h e spacing and the directional orientation of the adjacent molecules. The relative orientation of t h e two molecules also has an effect on t h e relative strength of the H band and the J band.'"-":' The simple exciton theory of the point dipole approximation mentioned above gives the expectation t h a t the H band is always larger than the J band in the parallel plane dimer configuration, while the J band is allowed to be larger than the
1345
Absorption Spectra of Rhodamine B in Various Media TABLE I : Results Obtained for Various Parameters for Rhodamine B Dimera
Solution
AG,
Dimer peak,eV
Glycerolc J H
cY,deg L , A
2.19 2.36
Ethylene J, 2.155 glycolc J, 2.195 H , 2.32
H, 2.39 Acetic acidc
J, 2.21 J, 2.26 H , 2.36
H, 2.43
}1 f
K,M
kcal/ mol
138
10.4
2.1 x 10-4
5.0
135 127
10.4 9.8
6.3 x 10-4
4.4
132 132
9.5 9.7
6.3 x 10-4
Figure 5. Schematic representations of simple dimer structures: (a) “parallel plane dimer”, (b) “oblique plane dimer”. The ovals in the figures correspond to the molecular profiles, the arrows the transition
dipole moments in the individual molecules at an angle 8 with each other, and L shows the distance between dimer molecules.
4.4
AE = 2DAfL:’
J 2.183a 520 8.3a 6.8 X 10-46 4.26 H 2.3750 0 Experimental results are from ref 6. 6 Experimental results are from ref 7. c Rhodamine B forms an oblique plane dimer. d Rhodamine B form a parallel plane dimer.
H,Od
where A is equal t o (3 - cos 8 ) / 2 a n d L indicates the spacing of t h e dimer molecule. T h e monomer dipole strength is related to the oscillator strength byI9 D = (3h2e2/8n2rn)(f~/E~)
-G
(8)
in which h is Planck’s constant, e t h e electron charge, and m indicates the electron mass. Then, the intermolecular distance L of the oblique plane dimer is solved as
Monomer
Dimer
Figure 4. Energy level diagram for monomer and dimer electronic energy levels and optical transitions. E and G indicate excited states and ground state, respectively. H means H band and J shows J band.
H band in t h e oblique plane dimer structure.22T h e parallel plane and t h e oblique plane dimers are schematically shown in Figure 5 . Aqueous rhodamine B, similarly to aqueous methylene blue, was proved to have a parallel plane configuration by dimerization, since the H band of these dyes in water was observed to be larger than t h e J band. In contrast to t h e results of aqueous rhodamine B and methylene blue mentioned above, the J band is found to be larger than t h e H band in glycerol-, ethylene glycol-, and acetic acidic-rhodamine B. Therefore, rhodamine B molecules in these solutions are expected to dimerize in order to have oblique plane configurations. Similar dimeric configuration of t h e oblique structure is observed in alcoholic ~ h l o r o p h y l l . ~ ~ It follows from the simple exciton t h e ~ r y l ~ -that ~ : j the dipole strength for the H band of t h e dimer is given by D ( l cos B ) , and t h a t for the J band of the dimer, D ( l - cos H ) , where 8 is the relative orientation and D, the monomer dipole strength. T h e angle 8 may be determined from t h e relative oscillator strength of the split bands
+
8 = 2 tan-’ (EHf.j/EejfH)”2
(5)
where E H , E,J,f ~and , f.J are the peak positions and oscillator strength of the H and J bands, respectively. The oscillator strength of the monomer f~ is determined by integrating the molar absorption coefficient a,(E) over the monomer band25
f~
(7)
= 1.51 X 10-‘J~a,(E) d E
(6)
in which the energy E is given in electron volts. The energy splitting AI3 of the oblique plane dimer is derived as
L = [3.29 X 1 0 2 f ~ A f ( E ~ A E ) ] 1 ’ 3 (A)
(9)
in which EMdenotes the monomer peak energy and AE t h e splitting in electron volts. T h e obtained dimeric configurations and spacings of rhodamine B in solutions are listed in Table I. So far t h e observed results were analyzed by t h e simple exciton theory and valuable information was obtained by it. However, some more complicated phenomena were also observed. As mentioned previously, dimers of rhodamine B in ethylene glycol and acetic acid have absorption spectra with two J bands and two H bands. The simple exciton theory does not seem t o explain the complicated character of t h e absorption curves. An explanation of these behaviors will need more detailed study concerning the mutual orientations of atoms in the dimer molecules and the role of molecular vibrations and the vibronic coupling between various states of the dimer.13 The simple exciton theory may be an oversimplified one, because it assumes a molecule as a point dipole and only t h e electronic states are taken into account. T h e “H” class dimer is thought to be nonfluorescent, while the “J” class dimer is fluorescent.‘ As a matter of fact, molecules of rhodamine B in water were proved to become nonfluorescent by dimerization, since these dimers show strong “H” a b s o r p t i ~ n ,whereas, ~ ~ ’ ~ our results show larger J-bands absorption in case of glycerol, ethylene glycol, and acetic acid. Similar spectral occurrence of strong “J” absorption has been observed for rhodamine B in EPA.7 As was described previously, the absorption spectra of rhodamine B in all solutions we examined are shown to be composed of those of fluorescent monomers and of fluorescent dimers. In the present work, glycerol-rhodamine B is seen to have the highest total absorption intensity (see Figure 2) and then, seems to show most efficient fluorescence. Actually, rhodamine B is said to show a fluorescence quantum yield of nearly unity in viscous solutions such as glycerol.’2 Details of the fluorescence processes will he clarified through studies on the fluorescence spectra, The Journal of Physical Chemistry, Vol. 80, No. 12, 1976
1346
D. W. James, R. F. Armishaw, and R. L. Frost
the relaxation times, and the quantum yields of both monomeric and dimeric rhodamine B in solution.
Acknowledgments. The author is indebted to K. Asakura and T. Yamaguchi for their technical assistances. The author also wish to thank Professor J. Yamashita for reading the manuscript. A part of this work is supported by financial aid from the Matsunaga Science Foundation. References and Notes ,
(1) E. P. Schaefer, W. Schmidt, and K. Marth, Phys. Lett., 24A, 280 (1967). (2) B. B. McFarland, Appl. Phys. Lett., 10, 208 (1967). (3) 0.G. Peterson and B. B. Snavely, Appl. Phys. Lett., 12, 238 (1968). (4) S. A . Tuccio and F. C. Strome, Jr., Appl. Opt., 11, 64 (1972). (5) R . W. Chambers and D. R. Kearns, J. Phys. Chem., 72, 4718 (1966). (6) K . K . Rohatgi, J. Mol. Spectrosc., 27, 545 (1968). (7) J. E. Selwyn and J. I. Steinfeld, J. Phys. Chem., 78, 762 (1972). (8) J. Ferguson and A. W. H. Mau, Chem. Phys. Lett., 17, 543 (1972).
i20) (21) (22) (23) (24)
J. Muto, Jpn. J. Appl. Phys., 11, 1217 (1972). J. Muto, Keio Engineering Report No. 25, 71, 1972. K. H. Drexhage, Laser Focus, 0, 35 (1973). K. H. Drexhage, "Dye Lasers", Vol. 4, "Structure and Properties of Laser Dyes", E. P. Schaefer, Ed., Springer-Verlag, New York, N.Y., 1973, p 114. R . W. Chambers, T . Kajiwara, and D. R . Kearns, J. Phys. Chem., 78, 380 (1974). M. M. Wong and 2 . A. Schelly, J. Phys. Chem., 78, 1691 (1974). W. H. Meihulsh, J. Opt. SOC.Am., 52, 1256 (1962). Th. Foerster and E. Koening, 2.Elektrochem., 81, 344 (1957). V. L. Levshin and K. V. Krotova, Opt. Spectrosc., 13, 457 (1962). Discussion In ref I O that the chemical reaction of monomers with the solvent is responsible for the spectral behaviors of acetic acidic-rhodamine B Is thought to be erroneous, A . R . Monahan and D. F. Blossey, J. Phys. Chem., 74, 4014 (1970). 0.S. Levinson, W. T . Simpson, and W. Curtis, J. Am. Chem. SOC.,79,4314 (1957). G. Hoijtink, 2.Elektrochem., 84, 156 (1960). M. Kasha, H.R. Rawls, and M. A. El-bayoumi, Pure Appl, Chem., 11, 37 1 (1965). S. S. Brody and M. Brody, Nature, 180, 547 (1964). Influence of the internal field, such as the Lorentz field or Onsager field, on the oscillator strength is not taken into account in eq 6.
Structure of Aqueous Solutions. Librational Band Studies of Hydrophobic and Hydrophilic Effects in Solutions of Electrolytes and Nonelectrolytes David W. James, Richard F. Armishaw,. Chemistry Department, University of Oueensland, St. Lucia, Oueensland 4067, Australia
and Ray L. Frost Chemistry Department. Oueensland Institute of Technology, Brisbane, Oueensland 4067, Australia (Received August 29, 1975)
The infrared librational spectrum of water has been examined for aqueous solutions of a series of alkyl substituted ureas and thioureas, formamide, acetamide, acetone, and a series of symmetrical tetraalkylammonium nitrates. The spectra enable separation of effects which can be described as structure making and structure breaking. The unique solution behavior of urea is shown to be dependent on its ability to hydrogen bond in a pseudo-tetrahedral pattern. The tetraalkylammonium salts are shown to decrease the tendency of water to hydrogen bond to four adjacent molecules but a t the same time increase the strength of the remaining hydrogen bonds.
There have been several studies of the vibrational spectra of solutions of nonelectrolytes most of which have examined the OH stretching region or overtone region in the infrared ~ p e c t r a . l -There ~ has been one Raman study6 and one recent infrared study7 of the librational region of aqueous solutions of sucrose and urea. I t is expected that a study of the librational band of water in aqueous solution will give direct information on the nature and extent of hydrogen bonded interactions. We report here a systematic study of the librational band in solutions of substituted ureas and some related solutes and the results are compared with those obtained for solutions of tetraalkylammonium nitrates. The properties of aqueous solutions of urea have been extensively studied. Although dielectric constant measurements indicate that there is an increase in hydrogen bonding in solutions,8 other studies indicate that urea destroys the long-range order of Thus it has been shown that urea raises the critical micelle concentration of dodeThe Journal of Physical Chemistry, Vol. 80, No. 12, 1976
cylpyridinium iodide and various dodecyl sulfatesg and this was attributed to a decrease in the ordering of the water. In an examination of the ability of urea to denature proteins it was noted that addition of urea t o solutions of serum albumin caused while alkyl ureas have the opposite effect to an extent which is dependent on the length of the alkyl chain and the number of alkyl substituents.'OJ1 Certain nonelectrolytes (dextrose and tetramethylurea) were found to be nondenaturants while acetamide and methylurea were found to be weak denaturants." Thermodynamic studies have confirmed the different behavior of urea and substituted ureas in solution.12 Aqueous solutions of symmetrical tetraalkylammonium salts have been extensively studied and the results have been recently reviewed.I3 I t is concluded that in dilute solution salts containing large cations, for example, tetrapropylammonium (Pr4N+)ortetrabutylammonium (n-Bu*N+) produce a structure enhancement, while at higher concentration the effect of ion-ion interaction becomes important.
