3872
MICHAELE. LAMMAND DAVIDM. NEVILLE,JR.
films provided the room-temperature points on the solvus. The solubility of barium in sodium is of the order 0.5 atomic % barium at room temperature and a maximum solubility of 3 atomic % at 6 5 O , the peritecOoid temperature. The solubility of sodium in barium was found to be 3 atomic % sodium at room tempera-
ture and a maximum solubility of 5 atomic % a t 197O the peritectic temperature.
Acknowledgment. The authors wish to express their gratitude to the Atomic Energy Commission for the financial support which made this study possible.
The Dimer Spectrum of Acridine Orange Hydrochloride
by Michael E. Lamm* and David M. Neville, Jr. Laboratory of Neurochemistry, National Institute of Mental Health, Bethesda, Maryland (Received M a y 88, 1966)
The absorption spectra of the cationic dye acridine orange hydrochloride have been deterM . Over this range mined at eight different concentrations in water from loP6to the spectra change continuously. An isosbestic point is observed at 470 mk. The spectral data are interpreted in terms of monomer-dimer equilibria. Agreement of the data with two different equilibria has been checked by means of a computer program employing a reiterative procedure which varies the equilibrium association constant, K , as an arbitrary parameter until a value of K is found which gives the smallest root-mean-square deviation of the optical density data with the equilibrium model. The two models, (1) involving dye cations only, 2D+ = Dz2+,and (2) anions as well, 2D+ A- = D2A+, fit the data equally well with K = 1.05 X lo4 I. mole-' and K = 4.7 X lo8 respectively. Dimer spectra obtained by extrapolating the data with the above models differ. Dimer spectrum 1 shows a symmetric splitting in relation to the monomer peak while dimer spectrum 2 exhibits a redistribution of intensity between the monomer peak band and shoulder band. The significance of the dimer spectrum in relation to current theories of metachromasy is discussed.
+
Introduction The pronounced effect of concentration on the color of aqueous cationic dye solutions has been known for over 50 years. Various physical descriptions of this phenomenon have been proposed, yet to date the superiority of any one model remains to be demonstrated. The spectral changes are known to be associated with reversibIe poIymerization of dye, and at low concentrations the predominant aggregate species is believed to be a dimer.2 A knowledge of the absorption spectrum of the dimer is crucial to the understanding of the physical process involved in the spectral The JOUTTMZ~ of Physical Chemistry
changes. Unfortunately, the dimer never exists alone and its spectrum is obscured by contributions from monomers and higher aggregates. To date, the dimer spectra of at least four different dyes have been derived utilizing certain critical assumptions.a-6 In each case the reaction has followed the scheme ~
*Department of Pathology, New York University School of Medicine, New York, N. Y. (1) S. E.Sheppard, Proc. Roy. SOC.(London), A82, 250 (1909). (2) V. Zanker, 2.phg8ik. Chem., 199, 226 (1952).
THEDIMERSPECTRUM OF ACRIDINE ORANGE HYDROCHLORIDE
2D+ = D2Z+
(14
or
60
2D = D2
3873
Ob)
xx 00 AI
The ability of these models to fit the data has reduced interest in more complex models. Recently, however, Haugen and Hardwick' have presented evidence for a model which includes counterion participation in the equilibrium
2D+
+ A-
no 50
I1 fi BB
tt 40
=
D2A+
(2)
I n the present study we have investigated the nature of the dimer spectrum of acridine orange hydrochloride by employing a computer program to determine the best fit of a large quantity of spectral data to two different dimerization equilibria.
I i 0
X Y
30
20
Methods Purification of AO. Twenty-five grams of commercial acridine orange (Chroma) was dissolved in 250 ml. of ice-cold water and titrated with 0.1 N NaOH to pH 10.5, The resulting suspension of free base was filtered in the cold, washed with water to remove residual salt, dried at loo", and dissolved in methanol. The pH was lowered to 5.5 by the addition of an acid-methanol solution (1 vol. of concentrated HCl and 9 vol. of methanol), and the hydrochloride, AO. HC1, was precipitated by the addition of 10 vol. of diethyl ether. The dye was collected and dried, and the whole process was repeated once more. One gram of AO.HC1 was dissolved in absolute ethanol and passed through an 8 X 40-em. acid-washed alumina column run in ethanol. Most of the dye emerged in the first peak leaving a pink trailing band and a pink band at the origin. The eluate comprising the main fraction was filtered to remove any alumina particles present. When rerun on both a small alumina column and on a thin layer of silica gel in CHC13-CH30H, 3 : 1, an aliquot of this solution moved as a single band. The AO.HC1 was precipitated from the eluate by t'he addition of ether, collected by filtration, and dried at 100". To obtain a sample of dye for determination of the molar extinction coefficient, 200 mg. of the purified AO-HC1 was converted to the free base as described above. The insoluble base was removed from the aqueous suspension and dried in vacuo overnight at 100". The free base was dissolved in chloroform, which can be removed quantitatively.2 The chloroform was evaporated off and the dye was dried in vacuo at 100" overnight (yield 120 mg.). Less than 0.25% total chlorine was found on analysis of this final sample.
10
410
430
450
470 A,
490
510
530
mu.
Figure 1. Absorption spectra of acridine orange hydrochloride in water are shown in units of molar extinction coefficients over a concentration range of 2.5 X t o 1.8 X 10-4 M .
Determination of Molar Extinction Coefiients. The main problem that had to be overcome in order to determine accurate and reproducible molar extinction coefficients was adsorption of dye from solution. The following surfaces were explored : Pyrex glass, siliconed glass, glass exposed to hydrofluoric acid, Teflon, polypropylene, and polyethylene. Polyethylene bottles (Nalge Gorp., Rochester, N. Y.) were found to be most suitable. During the time required to take the absorption spectrum, absorption could be held to about 1% of the dye content per dilution. No adsorption occurred during the first few minutes after addition of dye (3) G.S. Levinson, W. T. Simpson, and W. Curtis, J. A m . Chem. SOC., 79, 4314 (1957). (4) K.L. Arvan and N. E. Zaitseva, Opt. Spectry., 11, 3 8 (1961). (5) K. Bergmann and C. T. O'Konski, J. Phy8. Chem., 67, 2169 (1963). (6) G.R. Haugen and W. H. Melhuish, Trans.Faraday SOC.,60, 386 (1964). (7) G. R. Haugen and E. R. Hardwick, J . Phye. Chem., 67, 726 (1963).
Volume 69,Number 11
November 1966
3874
solution to quartz cuvettes which had been rinsed with water, hydrochloric acid in ethanol, water, and ethanol, provided the dye solution was not stirred or mixed within the cuvette and providing the dye conM . As the dye concentracentration was above tion was reduced below M , optical density readings a t a given dilution showed increased scatter, especially below 470 mp. Only part of this error could be accounted for by the increase in dye adsorption a t low concentrations. Stock solutions of dye were prepared by placing a weighed sample of purified base in a dry polyethylene bottle and adding 1.05 equivalents of HC1 converting all of the dye to the monovalent cation. Distilled water was added to make the dye concentration in the range 2-7 X 114. Stock solution pH ranged between 5.4 and 6.0. Aliquots of stock solutions were poured into dry polyethylene bottles, diluted with distilled water, mixed by swirling, and poured into 1- or 2.5-cm. path length cuvettes which had been cleaned as described above. Dye solutions were added to 0.1-cm. cuvettes by means of capillary pipets which had been rinsed in the dye solution. Optical densities were read on a Cary Model 14 spectrophotometer. For very dilute solutions an expanded-scale slide wire was used. All additions and dilutions were made by weight, and all determinations were done at 25 f 0.5". Molarities were calculated by weights using a formula weight of 265 for A 0 base. The assumed identity of the formula weight and the molecular weight was based on the chromatography results and the demonstrated elimination of solvent of crystallization. All calculated molarities were corrected for dye adsorption by reducing the initially calculated value 1% per dilution. From the calculated molarity and the observed optical density a t each dilution, the molar extinction coefficient was determined at that concentration. Multiple determinations were made from three separate weighed samples of solid A 0 base.
Results The visible absorption spectra in units of molar extinction of acridine orange hydrochloride over a 100fold range in concentration are shown in Figure 1. The dye does not obey Beer's law. I n dilute solutions the absorption maximum occurs a t 492 mp and a shoulder is present at -470 mp. As the concentration is increased, the shoulder becomes more prominent as the main band decreases in intensity. This phenomenon is characteristic of the metachromatic dyes. At 470 mp an isosbestic point is present, E 43,000, indicating that only two spectral species are present over M . The presence the concentration range to The JOUTTW,~ of Physical Chemistry
MICHAELE. LAMMAND DAVIDM. NEVILLE,JR.
of only two spectral species permits us to represent the total dye concentration C in units of moles per liter and the total molar extinction coefficient E as the sum of contributions from two species, monomer and dimer.
c E
=
e
(Y(€M)
CM
+ 2cD
+ (1 - a)ED/2
(3) (4)
Where a is the fraction of dyes present as monomers +) a=------
(5)
C
Assuming that CM and ED are independent of concentration, a plot of t vs. a a t any wave length should yield a straight line with intercepts a t E M for CY = 1 and ED/2 for a = 0. The value of a can be computed from the relation between a and the association constant. The association constant for equilibrium 1, KI, is 17
The association constant for equilibrium 2, K2, is in the form of a cubic equation, the real roots of which were found by the usual methods
(7) The value of K was determined by a reiteration procedure. I n this study we used a computer program developed in this laboratory by Dr. H. DeVoe to perform the reiteration. An arbitrary K is chosen and the corresponding values of a for each of several solutions with different concentrations are calculated using eq. 6 or 7. For eachwave length the best straightline fit of E vs. a is found by weighted least squares. The procedure is repeated with different values of K until a minimum is reached in the root-mean-square deviation of the experimental O.D. values from the computed best straight-line values. The deviations in O.D. were weighted to be proportional to C. When the data shown in Figure 1 were used as input (involving eight solutions and thirteen wave lengths), the computer found a single minimum for both equilibria. Figures 2 and 3 show a plot of root-meansquare deviation in units of O.D. vs. K for equilibria 6 and 7. For eq. 6 the minimum corresponds to K = 1.05 X lo4 1. mole-' while for eq. 7 the best value is K = 4.7 X lo8 L2 mole-2. At the minimum of each curve the root-mean-square deviations are essentially equal. Consequently, the data fit both equilibria
THE DIMERSPECTRUM OF ACRIDINE ORANQE HYDROCHLORIDE
3875
0.05
s
.I
? d
0.04
0
;0.03
d
;0.02
d 0
0
0.01
00
i Amur
5
4
3
Log K.
0
0.1
0.2
0
0.3 0.4
0.5
0.6
0.7
0.8
0.9
1.0
a.
Figure 2. The average root-mean-square deviations in units of O.D. for plots of e us. Q a t 13 different wave lengths and eight concentrations are plotted against K . e is taken frum the data of Figure 1. Q is computed from eq. 6 a t each arbitrary value of K . The minimum deviation, 0.006, occulg at K = 1.05 X lo4
Figure 4. Molar extinction coefficients a t the indicated wave lengths and at eight different concentrations are plotted us. a,the fraction of monomers. The value of Q is computed from eq. 6 using the value of K indicated in Figure 2 and the known dye concentration. T
0.05 4
'a
0.04
6
0- 0.03
2' .c1
4 0.02
e d
0.01
8
Log
K.
9
10
0
0.1
0.2
0.3
0.5
0.4
0.6
0.7
0.8
0.9
1.0
a.
Figure 3. Same plot as in Figure 2 except that cy is computed from eq. 7. The minimum deviation, 0.007, occurs at K = 4.7 X 108 M - 2 .
Figure 5. Same plot as in Figure 4 except the value of cy is computed from eq. 7 using the value of K indicated in Figure 3.
equally well. The agreement between the data and the two models can be better visualized in Figures 4 and 5, where E vs. a is plotted at several wave lengths for both models using the best value of K for each case. The deviation of each point from the line represents the deviation of the data from the model. The derived dimer and monomer spectra shown in Figure 6 were obtained from the intercepts of the straight lines at CY = 0 and a = 1 and include data obtained from additional solutions and wave lengths not used to determine the equilibrium constants.
between the minimum root-mean-square deviations for each model, 0.001, is not significant. I n general, both models fit the data equally well at all wave lengths except at the lowest concentrations where the points below 470 mp show larger deviations. Previous studies have reported dye dimer spectra derived on the basis of eq. 1. In each case a fit of the experimental data to the single model was taken as the major evidence in support of the model. It is obvious that the ability of a model to fit data is not sufficient cause for acceptance of the model, and, unless predictability is a major feature of a model, all other possible models must be eliminated by some means. I n the present study we have tested what we consider to be the two most likely dimerization models. Since
Discussion The surprising result of this study is the finding that both dimerization models fit the data. The difference
Volume 69, Number 11
November 1966
MICHAELE. LAMMAND DAVIDM. NEVILLE,JR.
3876
60
40
88
> 2 X
20
540
500
x, mrr.
460
420
Figure 6. The extrapolated monomer and dimer spectra in units of molar extinction coefficient per monomer are shown for both equilibria.
the data presented here do not provide a basis for choosing between the models, and since the literature contains conflicting data on the question of anion participation in monomer-dimer e q ~ i l i b r i a , ~ *we ~J will discuss the significance of both sets of spectra, recognizing that at least one of them is incorrect. The monomer spectra derived from the different dimerization equilibria are similar and closely resemble the spectrum of unassociated dye in organic solvents. Both spectra display a shoulder at shorter wave lengths (-470 mp) characteristic of monomeric cationic dyes. Sheppards has proposed that the shoulder is derived from vibrational coupling with a single electronic transition. I n support of this thesis, Zankere has demonstrated a mirror image relationship between the monomer absorption and fluorescence spectra of acridine orange. From Zanker’s data it appears that the 492-mp peak and the 470-mp shoulder represent the 0 0 and 0 + 1 transition of a major vibrational progression. The significance of the present study in relation to these data can be seen by examining the dimer spectrum derived on the basis of anion inclusion. The absorption maximum occurs a t 470 mM, coincident
with the 0 + 1 transition while a prominent shoulder exists at -495 mp coincident with the 0 -t 0 transition. If this represents the correct spectrum, it provides strong support for the idea of Sheppard and Zanker that dimerization induces a redistribution of the transition probabilities between the 0 + 0 and 0 -+ 1 states. In Zanker’s scheme this is accomplished by a differential shift in the ground and excited state potential energy curves of each monomer, thus altering the Franck-Condon vibronic intensity factors. More recently, DeVoe’O has shown that a similar redistribution of intensity within the band is predicted when the dyes are oriented parallel to one another in the dimer and the oscillators in the dyes are coupled by coulombic forces. The dimer spectrum derived from eq. 6 fails to show a band in the region of the monomer 0 --t 0 transition. Instead, there are two bands roughly symmetrically
-+
The Journal of Physical Chembtry
(8) 8. E. Sheppard, Rev. Mod. Phys., 14, 303 (1942). (9) V. Zanker, M. Held, and H. Rammensee, 2. Naturforsch., 14b, 789 (1969). (10) H. DeVoe, J . Chem. Phys., 41,393 (1964).
THEDIMERSPECTRUM OF ACRIDINE ORANGE HYDROCHLORIDE
distributed above and below the 0 + 0 band, a strong band at -470 mp, and a weak shoulder at -515 mp. This appears to be a property of dimer spectra derived from eq. 6 and has been noted for pyrid~cyanine,~ rhodamine14 methylene blue,6 and proflavine.6 This type of spectral splitting is predicted for dyes which are oriented at a small angle to one another in the dimer, by the strong coupling exciton theory of Levinson, Simpson, and cur ti^.^ Their theory states that the degeneracy of the energy levels within the dimer results in a splitting of these levels and an unequal distribution of intensity between them. We see then that depending on which equilibrium model is used we derive different dimer spectra, each of which tends to support different physical explanations of the spectral shifts. The conflicting evidence within the literature on the question of anion partticpation in dye dimerization equilibria points up the difficulty in designing experiments to test these models. Part of the problem is due to the narrow concentration range over which dye solutions are free from higher aggregates. The addition of electrolyte narrows this range even further, so that it becomes
3877
difficult to distinguish between ionic strength effects and specific anion effects. To illustrate the complexity of the problem, we have noted that certain organic electrolytes such as guanidine hydrochloride can either promote association or dissociation depending on whether they are present in low or high concentrations.ll This is undoubtedly due to the unique properties of the guanadinium ion, but it illustrates the hazards of drawing specific conclusions about ion participation in a given equilibrium when the ions in question may also alter the solvent structure.12 Perhaps the least ambiguous experiments would be the determination of the association constants for a series of dye salts in the absence of added electrolyte. Anion participation should be reflected by variations in K among the various salts.
Acknowledgment. The authors are greatly indebted to Dr. H. DeVoe who made his computer program available for this study. The authors also wish to thank Dr. D. F. Bradley for many helpful discussions. (11) D. M. Neville, Jr., unpublished data. (12) W. P. Jencks, Fed. Proc., Suppl. 16, 5-50 (1966).
Volume 69, Number 11 November 1966
