Dyes as biologic probes, color and dye microenvironments

•

.

1,2

Dyes as Biologic Probes, Color and Dye Microenv,ronmenrs C H A R L E S E. W I L L I A M S O N A_ND A L S O P H H. C O R W I N Department of Chemistry, The Johns Hopl~ins University, Baltimore, Maryland 21218

Received January 25, 1971; accepted June 8, 1971 The color changes of several dyes were studied in the presence of detergent micelles. Changes in the absorption maxima of the dyes and Mterations in their equilibrium constants were measured in solvents of varying polarity. Similar effects were measin detergent mieelles as a result of the hydrophobic microenvironment experienced by the micelle-bound dye. In addition, microenvironmental electrostatic effects were observed. The latter cause large changes in the equilibrium constants of the dyes as a result of the juxtaposition of charged functional groups in the detergent micelles. (4) reported their ability to shift indicating dye equilibria. Until recently, however, the mechanisms have been unclear. Changes in reactions induced by cationic, anionic or neutral mieelles formed from detergents such as lauryltrimethylammonium bromide (LTAB) or sodium lauryl sulfate (SLS) have been examined by several investigators (5-12) and recently reviewed by Fendler and Fendler (13). The kinetics of such reactions as the acid or base-catalyzed hydrolyses of esters (7, 8, 13), Shift bases (11), or the hydroxylation of certain dyes in aqueous solutions are altered in the presence of detergents. Although these effects are normally observed above the critical micelle concentration, it has been shown that substrates containing relatively nonpolar moeities such as the phenyl group in methyl orthobenzoate (7) can induce micelle formation below the critical level, thus incorporating the substrate and locating it in the relatively apolar micelle environment. Duystee and Grunwald (6) consider van der Waal interactions of the London dis1 This work was supported in part by Grant 5 persion type the outstanding force of asR01 HE07496-07 from the U. S. Public I-IeMth Service, National Heart Institute, Bethesda, Md. sociation in these systems and emphasize interactions of short range. Thus, these 20014. Taken in part from a dissertation submitted workers consider such organic compounds by C. E. W. to this Department in partial fulfill- in aqueous solution as solvated b y other organic molecules in preference to water. ment of the requirements for the PhD Degree. Suggested as a general term for such comThe observed alterations in reaction pounds by Edelman and McClure (1). ldnetics can be explained by invoking When employed as biologic stains, certain dyes respond to their environment in tissue with a color change. Although the mechanism of this phenomenon (metaehromasy) is vague, an interpretation of the forces that act on the dyes can produce information to describe the chemical nature of the binding sites of proteins and other maeromoleeules. T h e use of dyes as chemo-optieal a probes for the study of maeromolecules offers advantages over fluorescent probes such as 1anilinonaphthalene-7-sulfonate and 2-toluidinylnaphthalene-6-sulfonate which have been used by a number of workers to detect the hydrophobie nature of protein binding sites (1-3). Dyes may be chosen to investigate the polarity as well as the electrostatic nature of biologic binding sites. In this investigation detergent micelles were employed as models to study the areas of macromolecules to which dyes bind. The marked effects of mieelle-forming detergents on the kinetics of certain reactions in aqueous solutions have been recognized since I-Iartley

Copyright © 1972 b y Academic Press, Ine.

Journal of Colloid and Interface Science, Vol. 38, No. 3, March 1972

567

568

WILLIAMSON AND CORWIN

Hartley's rules (4) which assume a cationic environment surrounding an anionic micelle and an anionic environment surrounding a cationic micelle. The reactions that ensue can be interpreted in terms of physical association, and attraction or repulsion of ionic reactants. The extent of the attraction of substrates to apolar areas is quite important in terms of the rates (and therefore type) of reaction to be expected in aqueous medium. Such effects as these can be of importance in tissues where apolar areas are common and hydrophobic sites on proteins widely reported. In this study, we have examined a number of dyes in an effort to interpret the mechanisms by which color changes occur in the presence of model systems. EXPERIMENTAL PROCEDURE

Materials Bromthymol blue, neutral red, methyl orange, janus green B, phenolphthalein, and the detergents sodium lauryl sulfate and lauryltrimethylammonium bromide were commercial products of reagent grade. The nonionic detergent, polyoxyethylene (20) sorbitan monopalmitate (Tween 40), Lot 304, was obtained as a sample from Atlas Chemical Industries, Inc., Wilmington, Del. All solvents were of commerical reagent quality or better and used without further purification.

Solutions and Measurements Most solutions wre prepared from a 0.l M, pH 7.4 phosphate buffer prepared from reagent-quality sodium hydroxide and 85 % phosphoric acid. All pH measurements were performed on a Beckman expanded scale pH meter (Beckman Instruments Inc., Fullerton, Calif.). Dyes were prepared in 0.1 M, pH 7.4 phosphate buffer as stock solutions (usually 10-3 M). Aliquots were diluted to the experimental concentration. The detergent solutions were prepared daily to the desired concentrations. Absorbance measurements were made with a Beckman DK 2 spectrophotometer or with a Bausch and Lomb Spectronie 600 spectrophotometer (Bausch and Lomb, Rochester, Journal of Colloid and £nterface Science, Vol. 38, No. 3, March 1972

New York 14602). Wavelengths were monitored with the aid of a Holmium oxide glass (H-98) standard (National Bureau of Standards, Washington, D. C. 20234). All aqueous solvent mixtures were percentage by volume. RESULTS AND DISCUSSION In these experiments, color changes in t h e dyes studied can be attributed to (1) bathochromic or hypsochromic shifts due to a change in the polarity (empirically deterlnined) of the solvent or the micellar microenvironment; (2) a shift in the equilibrium constant of an indicating dye due to a change in the polarity of the solvent or other local mieroenvironment; or (3) a shift in the equilibrium constant due to specific electrostatic effects of the environment that can be described in attractive or repulsive Coulombie terms. Since the visual color of a dye can be altered by one or a combination of these effects, interpretation of the color changes can serve to describe the chemical effects that are exerted on the dye by its environ~ ment. Although separation of these effects can be difficult, it is simplified by interpreting the color change of several dyes in the same environment. Several dyes have been used to illustrate effects 1, 2, and 3. Effects 1 and 2 were demonstrated in organic aqueous solvent systems, and effect 3 in detergent micelles.

Bathochromic and Hypsochromic Shifts due to Polar Environmental Effects The red or blue shifts in the spectra of dyes that accompany changes in the polarity of the solvent can be interpreted as one of the changes observed in metachromasy. These shifts were measured for several dyes in aqueous solvent systems and were related to the ionizing power scale Y (14). Certain dyes characteristically exhibit large shifts in ~ma~with a change in polarity of the solvent whereas others show little change. The former can be employed as probes to determine the polarity of the binding site of the dye, while the latter are more valuable in interpreting color changes due to other effects. In the past, such shifts have been correlated with physical char-

DYES AS BIOLOGIC PROBES

569

are the regions of polarity that appear to be most important in the study of proteinbinding sites (2). The same data plotted against the Z scale is sho~m in Fig. 2. The transition frequencies of janus green B, on the other hand, were not linear with any scMe employed, exhibiting the highest absorption frequencies in pure water or pure organic solvent and a minimum in approximately 50% water-solvent mixtures (Fig. 3). This diagram probably represents a composite of two curves; one describing energy changes due to a transition which is more difficult in aqueous than nonpolar media and the other describing energy changes due to solvation of the excited state possessing a

aeteristics of the solvent such as dielectric constant, refractive index or various empirical polarity scales (14-16). Although not well understood, this phenomenon has been employed to study the polarity of a dye's environment in micellar (4, 6, 9) and biologic studies (4, 17, 18). Parallel studies have also been performed employing fluorescent probes instead of dyes (1-3). Janus green B (I) and methyl orange (II), were selected for study because of the sensitivity of their Xm~ to rather small deviations in polarity of the solvent. When the polarity is lowered by the addition of a less polar solvent, a red shift is observed for janus green B and a blue shift for methyl orange. Tables I and II give the absorption maxima for these dyes in various solvents and the corresponding data relating to solvent polarity. Employing reported (2) procedures the transition frequencies of these dyes in the various solvents were plotted as a function of the dielectric constant, the empirical polarity Z scale (15) and as a function of the ionization power of the solvent (Y seale) (14). The best fit for methyl orange was obtained when its transition frequencies were plotted against the Y scale as shown in Fig. 1. For the solvents tested, the absorption frequencies of methyl orange are linear with Y from 3.5 to approximately - I . 0 . These

N~+

-03S~ . - N

N-~N (CH3)

=N ~

N(CH3)z

T-r SCheME A

TABLE I ABSORPTION MAXIMA OF METHYL ORANGEa IN SOLVENTS O~ VARYING POLARITY, 25°C Solvent (%)

Xmax (m~)

v(cm-1)

D

Zb

H20 Me0H (50) MeOH (80) MeOH (100) EtOH (100) n-PrOH (100) n-BuOH Dioxane (50) Dioxane (80) Dioxane (100) Acetone (100) t-BuOH (100)

465 456 432 419 417

21,505 21,930 23,148 23,866 23,981

415 444 422 416 412 413

24,096 22,523 23,697 24,083 24,272 24,213

78.5 58 42 33 24 20 17 32 10 2 21

94.6 90.9 87.3 83.6 79.6 78.3 77.7 86.7 80.2

Methyl orange concn= b R e f e r e n c e (19). R e f e r e n c e (14).

yc

3. 493 1.972 0.381 --1.090 --1.974 1.361 -0.835

65.7 71.3

1.5 X 10-5 M.

Journa~ of Colloid and Interface Science, Vol. 88, No. 3, March 1972

570

WILLIANISON A N D CORWIN TABLE II ABSORPTION MAXIMA OF JANUS GREEN B ~ IN SOLVENTS OF V-&RYING POLARITY, 25°C Solvent (%)

Xm~,x

tt20 MeOI-I (50) MeOH (80) MeOIt (100) E t O H (10) E t O H (30) E t O H (40) E t O H (50) E t O t t (100) Dioxane (10) Dioxane (20) Dioxane (30) Dioxane (50) Dioxane (60) Dioxane (80) Dioxane (90) Dioxane (95) Dioxane (100)

(m#)

603 654 654 645 618 641 658 660 651 623 638 657 663 662 662 663 634 612

Janus green B COhen 1.5 b Reference (19). Reference (14).

~, (cm-t)

D

Zb

yc

16,584 15,291 15,291 15,504 16,181 15,606 15,198 15,152 15,361 16,051 15,674 15,221 15,083 15,106 15,106 15,083 15,773 16,340

80 58 42 33 74 64 58 52 24 76 64 50 32 25 10 5

94.6 90.9 87.3 83.6 93.6 91.6 90.5 89.2 79.6 93.0 91.4 89.9 86.7 85.0 80.2 76.7

3.493 1.972 0.381 -1.090 3.312 2.721 2.196 1.655 - 1.974 3.217 2.877 2.455 1.361 0.715 --0.833 -2.030

2

X 10 -5 M .

2.4 -o E t 0 H oMeOH o 80%

Dioxane

I

o D X

o 80%

'T E v

MeOH

t..)

o 50%

D[oxane

u~ 2.2

o 50%

MeOH

o H20

2.1

I -I

t 0

r /

t 2

i 3

i 4

Y

FIG. 1. Plot of absorption frequencies of methyl orange in various solvents vs the polarity scale, Y (14). Journal of Colloid and Interface Science, Vol. 38, No. 3, March 1972

DYES AS BIOLOGIC PROBES

57t

o Acetone o t- BuOH o n-BuOH o MeOH o 80%

Dioxane

,¢

5 x

o 80%

MeOH

25

i

(D o 50%

Dioxone

c aD E;eu LL

2.2 o 50 % MeOH

o H20

2. I

i

i

70

8tO

75

8/

5

,

p

90

95

z

Fro. 2. Plot of absorption frequencies of methyl orange in various solvents vs the polarity scale, Z (15). 1.7

O-HzO

o-

)ioxene

o-lO%EtOH o - I 0 % Diox.

1.6

o-20%

E"

b x T E v

o-20% Diox. °-50 % EtOH

o-MeOH ,-EfOH

EtOH

o-80%MeOH 0-80%

Diox.

1.5 ~ 0 ° I o D i o x .

0,50% MeOH 50%EtOH o 0-30% Diox. ;¢0°/o E t O H o o -o 6 0 % ' O i o x . "50oA, O i o x .

LL

1.4

I

I

I

i

I

I

-I

0

I

2

5

4

Y FIG. 3, Plot of absorption frequencies of janus green B in various solvents vs the solvent polarity scale, Y (14). Journal of Colloid and Interface Science, Vol. 38, No. 3, March 1972

572

WILLIAMSON AND CORWIN 0.0 OI 0.2

~ Z 0

0.4 0.5 0.6 0.7 0.8

09

,o 1.1 1.5 [.4 1.5 550

', \ , \

91

,\.I \ I k~/ I 400 450

I 500

550

600

W A V E L E N G T H , nm

F r o . 4. T h e a b s o r p t i o n s p e c t r a of 5 X 10-~ M m e t h y l o r a n g e in w a t e r ( ) a n d in 80% m e t h a nol (---). 0.0 0. I

0.3 w

//

O.5 0.7

o.~' ,<

\\

r.o: ~.t

1'2 f 1.4 1.5 500

g/

/

~

g/ \

/

/

\",. T 550

I 600

/

[ 650

[ 700

750

WAVELENGTH, nm

FIG. 5. T h e a b s o r p t i o n s p e c t r a of 3 × 10.5 M j a n u s green B in w a t e r ( ) and in 80% m e t h a nol (---).

relatively large dipole moment. The latter should be a lower energy process in aqueous than nonpolar media. The absorption frequencies in Fig. 3 do, however, appear to be a function of the Y scale in the more polar regions of greatest interest. Although methyl orange is suited as both a qualitative and quantitative probe for polarity, janus green B may only be used qualitatively. Both were employed to probe the apolar areas of micelles and proteins. In the presence of low concentrations of micelles or proteins, shifts in the absorption frequencies of the dyes were assumed to be the result of their binding, responding to the lower polarity. The absorption spectra for methyl orange and janus green B in water and in 50 % ethanol arc shown in Figs. 4 and 5. Although many dyes aggregate at low concentrations, it was shown by conductivity experiments that methyl orange does not at the concentrations employed herein (19). When aggregation occurs in metachromasy (20) new absorption bands appear. This phenomenon was not observed in the case of janus green B. The absorption maxima of these dyes in various concentrations of detergents are tabulated in Table III. As might be expected, all dyes do not bind to all detergents at these concentrations. Employing detergent concentrations in the region of the critical micelle concentration, (0.008M for SLS (21) and 0.0016 M for LTAB (22)) positively charged dyes such as janus green B (I) were observed to bind to negatively charged

TABLE III ABSORPTION MAXIMA OF DYESa IN 0.1 M, p H 7.4 PHOSPHATE BUFFER CONTAINING VARYING CONCENTRATIONS O F DETERGENTS, 25°C Xraax

Detergent

Sodium lauryl sulfate Tween 40 L a u r y l t r i m e t h y l a m m o n i u m bromide Ethano~ No a d d i t i v e (buffer only)

Janus green B

Methyl orange

1.0

0.1

0.01

1.0

0.1

0.01

666 662 660

665 615 593 651 594

617 594 594

454 426 431

462 460 429 417 463

463 463 463

a M e t h y l orange and j a n u s green B c o n c e n t r a t i o n = 1.5 X 10 -s M . Journal of Colloid and Tnterface Science, Vol. 38, No. 3, March 1972

DYES AS BIOLOGIC PROBES

573

Under these conditions, the selective redetergents such as sodium lauryl sulfate, but not positively charged detergents such as moval of charged or uncharged dye species lauryltrimethylammonium bromide (Table from solution by mieelles can be expected to III). Negatively charged hydrophobic dyes cause equilibrium shifts. At concentrations such as methyl orange (II) bind to positively well above the critical mieelle concentration but not negatively charged detergent both forms are usually bound. We have micelles. Both positively and negatively studied the effect of polar and electrostatic charged dyes bind to the nonionie detergent forces on changes in indicator equilibria Tween 40, but more weakly. under these conditions. Dyes containing amino functions of low Shifts in Indicating Dye Equilibria by pK~, such as pyridine or aniline groups are Polar and Electrostatic Forces essentially electrically neutral at pH 7.4 and bind well to both positively and negatively The acidity constant of an acid: charged mieelles. An example is the binding HA ~- H + + A[1] of neutral red to both sodium lauryl sulfate and lauryltrimethylammonium bromide. may be written (23) as: The shifts shown in Table III illustrate K~A CAO~A7HA the extent of dye-mieelle interactions. It is = - 0LH+ ~ . . noted that the shift of Xm~ WaS always in the CH~ a~ 7Adirection of an apolar solvent effect. These and the conjugate acid of an amine: shifts presumably reflect the apolar environ÷ ment of the bound dye and serve as a basis RNHa ~- RNH2 ÷ H + [2] for interpreting the nature of dye-binding sites on proteins. aS: At higher detergent concentrations, association between the positive janus green B + CRNtt 2 K R N H 3 -~ and the positive lauryltrimethyl ammonium C+ RNH 3 bromide was observed. The equivalent + binding of the negative methyl orange to __ 0~RNH2 VRNH 3 -~ " ~H + • - - , sodium lauryl sulfate was also observed, but OLRIgH3 "YRNH2 somewhat to a smaller extent. The data in Table III can be readily where C is the concentration, a the activity, interpreted by considering the formation of 7 the activity coefficient, and a~2/aRNH `+ mixed mieelles by the detergent and dye •a~+ is the thermodynamic constant indemolecules. The formation of mixed mieelles pendent of the medium. in the region of the critical mieelle concenIf both are dissolved in a solvent of lower tration has been studied by Behme et al. (7). polarity, the charged activity coefficients Through the forces of hydrophobie-eleetro- will increase thereby producing a smaller static interaction the dye molecule of op- value for K~A and a larger value for K R N+H 3 • posite charge to the detergent is most These are the effects generally observed likely to participate in mixed micelle forma- when such equilibria are measured in ention, although a certain number of positive vironments of lower polarity either by the dye molecules will also be loeated within a addition of a solvent or through binding to positive mieelle. mieelles or macromoleeules. Shifts in the At higher detergent eoneentration, es- equilibria of the dyes neutral red and sentially all of the dye moleeules, regardless bromthymol blue are illustrated in Fig. 6 of charge type, will be incorporated into where the pH necessary to maintain equilibmieelles and specificity will be lost. Thus, it rium is plotted against the percentage of can be expected that the specificity of dyes ethanol in the aqueous solvent. for mieelles will only be observed at lower Thus, when an indicating dye is used as a detergent concentrations, and it is in this probe in the pH region of color change, region that the parameters of dye-mieelle polar effects can cause a shift in the dye interactions can be best studied. equilibrium in addition to the bathochromic .

_

O ~ H +

_

.

•

_

_

a H +

Journal of Colloid and Interface Sciencs, Vol. 38, iNTO.3, March 1972

574

WILLIAMSON AND CORWIN

/

I00

90

80

70

6O c

/

50 LO

o3

40

/

50

20

/

/

/

/

/

/

÷

though readily apparent visually, is not reflected in Table IV because the addition of acid or base can correct the color of the experimental sample for equilibrium but not bathochromic shifts. In the presence of charged micelles, large deviations in color were observed that can be attributed to the specific electrostatic effects described by Hartley (4) and can be easily rationalized by the Bjerrum (24) effects. Shifts in Dye Equilibria by Electrostatic and Polar Effects

The electrical work done under the influence of the potential in removing a proton from the acidic group at rl to infinity is (per mole):

+

A W = N¢~ = N Z ¢ ~ / D r l ,

where N is Avogadro's number, e the protonie charge, ¢ the electric potential, Z the valence, and D the dielectric constant of the medium (24). When the carboxylic function is bound to an anionic micelle the attractive electric potential exerted on the hydrogen atom will be approximately:

+

+

¢ = (¢/Drl) + (e/Dr2) + (e/Dr~) . . . , O

I 6

i 7 pH to

I 8

[3]

I 9

r

I0

Maintain Equilibrium

FIG. 6. Plot of pH necessary to maintain equilibrium vs the percentage of ethanol in the aqueous solvent; neutral red, © ; bromthymol blue, X. or hypsochromic shifts described above. These shifts can complement or oppose one another, giving rise to a color change that can only be interpreted by separation into its components. In this study, they were separated simply by emoloying several dyes that produce known bathochromic or hypsochromic shifts. These effects are illustrated in detergent micelle experiments where indicating dyes were allowed to interact with the nonionic Tween 40 (Table IV). It is noted that, in each instance, the equilibrium was shifted in a direction consistent with an apolar solvent effect. The bathochromic shift for indicators such as bromthymol blue, alJournal of Colloid and Interface Science, V o h 38, :No. 3, M a r c h 1972

[4]

where rl is the CO0-H bond distance and r2, r3, etc. are the distances from the micelle's negative charges to the departing hydrogen atom. If the local environment is of low dielectric constant, ¢ will increase further and the ionization process will become even more difficult. When the carboxylic function is bound to a cationic micelle the expression becomes: = (e/Dr~) -- (e/Dr2) -

(e/Dr3)...

[5]

and less energy is required for ionization. Similar equations can be written for other reactions involving a charge entering or leaving the charged micelle environment. The equilibria of neutral red and bromthymol blue are depicted in Eqs. [6] and [7]. The extent of the equilibrium shifts of these equations in the presence of detergents is given quantitatively in Table IV. A standard color was prepared by dissolving a dye in 0.1 M, pH 7.4 sodium phosphate buffer and adjusting to the test pH. The color obtained

575

DYES AS BIOLOGIC PROBES TABLE IV APPARENT p~-~ S H I F T OF D Y E S IN THE PB, ESENCE OF D E T E R G E N T S a Dye

pH of standard color

pH of dye-detergent mixture to reproduce standard color

Direction of equilibrlmn shift

Bromthymol blue Bromthymol blue Phenolphthalein Methyl orange Neutral red

7.25 6.5 8.85 3.9 7.0

7.15 6.15 8.95 1.2 5.45

Pos Pos Hyd-Neg Hyd-Pos Hyd-Pos

Bromthymol blue Bromthymol blue Phenolphthalein Methyl orange Neutral red

Sodium lauryl sulfate 7.4 9.45 6.5 8.5 8.7 9.3 3.9 4.4 7.0 9.5

Hyd-Neg I-Iyd-Neg Hyd-Neg Neg Neg

Bromthymol blue Bromthymol blue Phenolphthalein Methyl orange Neutral red

7.4 6.5 8.7 3.9 7.0

Tween 40

Detergent conch phosphate, 25°C.

=

0.03 M, dye conch

9.6 8.5 9.9 2.7 6.0 =

1.5

X

Hyd-Neg ttyd-Neg Hyd-Neg Hyd-Pos ttyd-Pos

10-5 M; all determinations conducted in 0.1 M,

H ( C H 3 ) z N ' ~ N/ - ~ N H+2 "

-.~ -N~"-.~ CH3

~

Red

N//''~#"C H 3

Eq. 6

Yel low

O/~H(CH3)2

O/~C H(CH3)2

Br H 3 ~ C ~ C

.

C H3~/I~

Br CH3

C

_ +H + Eq, 7

CH3~I

Br'"-.~"CH (CH3 )2 OH

Br''~/'CH (CH 3)2 0

Yellow

Blue E Q S , 6 AND 7

was t h e n visually reproduced in the presence of a detergent and the pI-I determined. These differences reflect the extent of equilibrium shift accountable to dye-detergent interactions. I n the presence of 0.03 M Tween 40, Eq. [6] was shifted to the right b y 1.0 p H unit (Tsble IV). This is presumably ~ solvent

effect on the equilibrium constant caused b y the hydrophobic microenvironment of the bound dye. A positively charged environment would be expected to potentiate this shift because of the electrical effect described in Eq. [4]. This is illustrated b y neutrsl red in 0.03 M L T A B where a p H change of 1.55 was observed (T~ble IV). This change preJournal of Colloid and Interface Science, Vol. 38, No. 3, March I972

576

WILLIAMSON AND CORWIN

sumably reflects the added effects of hydrophobic and positive electrostatic forces acting upon the dye. The data describing this effect in Table IV is noted "hydrophobicpositive". When this dye is bound to SLS the hydrophobic and negative electrostatic effects oppose one another and the observed changes in Eq. [6] should be substantially smaller. In these experiments, however, this was not observed. The negative appeared considerably stronger than the positive electrostatic effects in the neutral pH region. In the presence of 0.03 M SLS Eq. [6] was shifted strongly to the left and a pH change of 2.5 was observed. This effect is noted "negative" in Table IV since it is in an opposite direction to hydrophobic or positive electrostatic effects. In interpreting the results of Table IV it should be noted that Tween 40 will produce a larger micelle than LTAB or SLS and also that the hydrophobic effect on reactions such as Eq. [7] is greater than that on Eq. [6] (Fig. 6). Accordingly, Eq. [7] was shifted to the left by 2.0 pH units at neutral pH in both Tween 40 and SLS. In the presence of LTAB it was shited to the right but to a smaller extent due to opposing hydrophobie and positive electrostatic effects. Each dye in Table IV canbe interpreted in this manner. It is thus clear that either effect can predominate depending on the nature of the dye, the binding site, and the magnitude of hydrophobic and electrostatic forces. It is important in these studies that the nature of the micelle binding surface can be described by interpreting the color changes that occur when the dye interacts. This interpretation is employed in the following report to estimate the hydrophobie and electrostatic nature of some protein binding sites from color changes observed upon using dyes as biologic probes.

by the micelle-bound dye. In addition microenvironmental electrostatic effects were observed. The latter cause large changes in the equilibrium constants of the dyes as a result of the juxtaposition of charged 5mctional groups in the detergent micelles. REFERENCES 1. EDELMAN, G. M., AND McCLuRE, W . 0 . , Account~ Chem. Res. 1,65 (1968). 2. TURNER, D. C., AND BRAND, L., Biochemistry

7, 3381 (1968). 3. RUBALCAVA,B., MARTINEZ, D., AND GITLER, C. Biochemistry 8, 2742 (1969). 4. HAaTLE¥, G. S., Trans. Faraday Soc. 30,444

(1934). 5. BUNTON, C. A., AND ROBINSON, L., J. Org. Chem. 34,773 (1969). 6. DUYNSTEE, E. F. J., AND GRUNWALD,E., Or. Amer. Chem. Soc. 81, 4540 (1959). 7. BEItME, M. T. A., FULLINGTON,Z. G., NOEL, R., aND CORDES,E. H., J. Amer. Chem. Soc. 87, 266 (1965). 8. DUYNSTEE, E. F. J., AND GRUNWALD, E., Tetrahedron 21, 2401 (1965). 9. TONG,L. K. J., REEVES, R. L., AND ANDRUS, R. W., J. Phys. Chem. 69, 2357 (1965). 10. BRUICE, T. C., KATZHENDLER,J., AND FEDOR, L. R., J. Amer. Chem. Soc. 90, 1333 (1968). 11. BEHME,M. W. A., AND COI~DES,E. ~-~.,J'. Amer. Chem. Soc. 87,260 (1965). 12. WINTERS,L. J., AND GRUNWALD,E., J. Amer. Chem. Soc. 87, 4608 (1965). 13. FENDLER, E. J., AND FENDLEE, J. H., Advan. Phys. Org. Chem. 8,271 (1970). 14. GRUNWALD,E., ANDWINSTEIN, S. J., J . Amer. Chem. Soc. 70,846 (1948). 15. KosowER, E. M., J. Amer. Chem. Soc. 80,

SUMMARY The color changes of several dyes have been studied in the presence of detergent mieelles. Changes in the absorption maxima of the dyes and alterations in their equilibrium constants were measured in solvents of varying polarity. Similar effects were measured in detergent micelles as a result of the hydrophobic microenvironment experienced Journal o] Collqidand fnterface Science, Vol.38. No. 3, March 1972

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