1216
JAMESE. LUVALLEAND JUANITA M. JACKSON
Vol. 61
I n spite of the success obtained with the simple solubilities of hydrocarbons in anhydrous liquid eq. 4 and 8 in systems Containing only hydrocar- ammonia indicate that it is possible t o obtain quane n t stitative expressions for the activity of the hydrocarbonsi it is apparent that further r ~ f i ~ ~ ~must be made in the theory before azeotropes containing bons in such solutions when terms for the can be treated with the same success. Calculations just com- tion forces are included in the theoretical equations. data on the Further studies along this line may give relationpleted6 with the aid of accurate calculations for ‘Ys(6) T. E. Taylor, M. R. Fenskeand T. ILI. Reed, I l l , THIS JOURNAL, ships which in preaa. tems including polar compounds.
THE ZETA POTENTIAL OF SILVER BROMIDE I N THE PRESENCE OF VARIOUS SUBSTANCES OF PHOTOGRAPHIC INTEREST’v2 BY JAMESE. LUVALLE AND JUANITA M. JACKSON Technical Operations, Inc., Arlington Laboratory, -4 rlington, Mass. Received May 16, 1067
The zeta potentials of silver hromide sols have been determined in the presence of some compounds commonly found in photographic processing solutions. It has been shown that conclusions regarding the zeta potential a t the silver bromide surface are open to considerable question when gelatin, sensitizing dyes, quaternary salts, etc., are simultaneously present. The data clearly indicate the importance of adsorption processes in photographic development. The zeta potentialconcentration curves (for different substances) are approximately fitted by a Langmuir adsorption isotherm 0 = kc/(l kc) where k = 3 X lo8 cm,a/mole arid 0 = ( f i rm)/(r~ fm) where fo, f i and fm are the zeta potentials in the absence of the adsorbed compound, the zeta potential at concentration c i of the adsorbed compound and the seta potential when the surface is completely covered by adsorbed compound, respectively.
-
+
-
Introduction As part of a program on the investigation of the chemistry of the photographic development process, the electrokinetic zeta potentials of silver halide sols have been determined: (1) in the absence of all agents other than the potassium nitrate produced in the reaction of potassium bromide and silver nitrate; (2) in the presence of Ag+, Br-, HaO+, OH-, dyes, quaternary salts, thiourea, photographic developing agents and gelatin. The effect of charged ions adsorbed to the silver halide grain (which give rise to the zeta potential) upon the photographic process has been the subject of numerous papers which have been adequately summarized by Meese8 The problem involved is as follows: an anionic developer ion must be adsorbed to the silver bromide grain for development to take place. When a negative charge is on the grain, a coulombic repulsion is exerted on the anion. The repulsion may be computed by considering the energy necessary to move a point charge of one, two or three electron units up to the surface against a negative potential. The energy required in electron volts is just the product of the retarding potential and the anionic charge. This energy is converted to kcal./mole by multiplying by 23.05. Computation of the energy (1) This research was supported by the United States Air Force through the Air Force Office of Scientific Reseprch of the Air Researah and Development Command under Contract No. AF 18(600)-371. Reproduction in whole or in part is permitted for any purpose of the United States Government. (2) A more detailed form of this paper has been deposited as Document Number 5191 with the AD1 Auxiliary Publications Project, Photoduplication Service, Library of Congress, Washington 25, D. C. .4 copy may be secured by citing the Document Number and by remitting $2.50 for photoprints, or $1.75 for 35 mm. microfilm in advanae by aheck or money order payable to: Chief, Photoduplication Service, Library of Congress. (3) C. E. K. Mees, “The Theory of the Photographic Process,” The Maomillan Co., New York, N. Y., 1954, Reviaed Edition, Chapters 13 to 16 incluaive.
necessary to surmount the barrier showed that the energy requirement became quite appreciable for negative zeta potentials of even a few millivolts. Although extensive dat8aexist on the zeta potentials of silver iodide very few data are available on the zeta potentials of silver bromide sols. Davies and Halliday have adequately summarized the available data on silver bromide in their papers.6 The object of this investigation was to determine the effect of compounds commonly present in photographic emulsions or photographic processing solutions on the zeta potential of undialyzed sols.
Experimental and Calculations The apparatus consisted of a Northrup-Kunitz electrohoresis cell obtained from the A. H. Thomas Company. &he electrode units were similar to those described by Hartman, Bateman and Lauffer.6 Stationary levels in the cell were computed after the method of Abramson, Moyer and Gorin.’ Velocity profiles for the cell were determined and proved satisfactory. A11 measurements of the migration velocity of the silver bromide particles were made at the stationary levels. On the recommendation of Davies and Halliday,’ all glassware, beakers, pi ets, etc., and the electrophoresis cell were treated with shgfitly acid fluoride for at least one-half hour just prior to use t o remove silicic acid adsorbed to the walls of the vessels. Comparative experiments showed that properly aged polyethylene containers could be used to store sols and solutions. Distilled water was used for the reparation of all solutions. All measurements and all andling of the sols were done under Series 1A Red Safe Lights. The microscope illuminator was used with a red
R
(4) (a) E. J. W. Verwey, Chem. Reus., 16, 363 (1935); (b) E. J. W. Verwey and J. Th. G. Overbeek, “Theory of the Stability of Lyophobia Colloids,” Elsevier, Amsterdam, 1948. (5) (a) K . N. Davies and A. K. Halliday, Trans. Faraday Soc., 48, 1061 (1952); (b) 48, 1066 (1952). ( 6 ) R. S. Hartman, J. B. Bateman and M. A. Lauffer, Arch. Bioohem. Biophye., 89, 56 (1952). (7) H. A. Abramson, L. 8.Moyer and M. H.Gorin. “Electrophoresis of Proteina and the Chemistry of Cell Surfaces,” Reinhold Publ. Corp., New York, N. Y,, 1942, Chapter 3.
.
THEZETA POTENTIAL OF SILVER BROMIDE
Sept., 1957
Nter on it at all times when sols were being observed. All sols in this work were 4.00 X 10-4 M in silver bromide. Attempts to utilize 1.OO X 10-6 M silver bromide sols were made impractical by the failure of these dilute sols to exhibit an isoelectric point. At least two sols were prepared for each set of experimental conditions. No more than two velocity determinations were made with each filling of the cell. At least ten velocity determinations were made for each sol. The mean value of the ten determinations was used as the electrokinetic velocity for that sol in the computation of the zeta potential. The electrophoresis cell was mounted in an aluminum block which was in thermal equilibrium with the room. The room temperature ranged between 20 and 30” but was 25” for most of these experiments; hence, the calculations were based on a temperature of 25”. Measurements of pH were made with a Beckman Model G pH Meter. All H adjustments were made with nitric acid or potassium droxide, as buffers could not be utilized in this work. The pAg determinations were made with silver-silver chloride and silver-silver bromide electrodes prepared in these laboratories. The silver-silver chloride electrode was used as the reference electrode and the silver-silver bromide electrode as the probe electrode. The surfaces of these electrodes were renewed weekly and the potential of the cell I Ag-AgC111 N KCI I I 1N KBr IAgBr-Ag I determined. Tl& electrodes were considered satisfactory when the potential of the cell was within one millivolt of 151 mv., the theoretical value. Overbeek has discussed the computation of the zeta potential from electrokinetic data.’ For symmetrical electrolytes, the equation for the electrophoretic velocity is
1217
I I
$-
-
where n is the ratio of the refractive index of silver bromide t o that of water. The value of 1.69 was used for n. Figure 1 shows that the theoretical Rayleigh curve is as good as the experimental scattering curves for K / a = 1 and (Y = 1.4. Hence the theoretical curve was utilized. The absorbance, A , is related to the scattering coefficient by A K = 0.435nhra
where When the assumptions of Davies and Halliday regarding D and are used in equation 1, it becomes for 1-1 symmetrical electrolytes
n=-
p*,
0.184
(8)’f,
(KP)]
(2)
where the fl(m) are the tabulated Overbeek relaxation functions,’ u is the observed electrophoretic velocity, D is the dielectric constant of water, E is the measured field strength in volts per centimeter, q is the viscosity of water, f is the zeta or double layer potential, e is the charge on the electron, k is Boltzman’s constant, T is the absolute temperature, z is the charge on the ion, p is the frictional constant of the ion, K is the reciprocal of the average radius of the ionic atmosphere and r is the radius of the colloidal particle. The field strength, E , was computed from the voltage drop between the microplatinum electrodes in the cell. Kappa, K, was computed from the ionic strength by the equation
(3) where N is Avogadro’s constant, ci is the molar concentration of the i’th ion and zi the corresponding charge. The remaining symbols have the values assigned in the preceding paragraph. The radius, r , of the particle was computed from the absorbance a t 900 mp and the scattering coefficient. Experimental scattering curves9 of K / a , the scattering coefficient, versus a = 2 ~ r / XHzO and the theoretical Rayleigh scattering curve have been plotted in Fig. 1. The expression for a may be rewritten 2.66sr f
f
=
-
A
sir
(4)
(8) J. Th. G. Overbeek, “Advances in Colloid Science 111,” Interscience Publishers, New York, N. Y., 19S0, pp. 97-136. (9) The experimental mattering curve8 were computed by Dr. Nitka of the Imperial Chemical Industries, Manchester, England, from data obtained from the National Bureau of Standards, N.B.S., and the Atomic Energy Commiseion, A.E.C.
m 4/3nrap
( 7)
and m = 0.1878M (8) where m is the mass; M , the molarity and p ( = 6.40), the density of silver bromide; r is the radius of the particle and I is the length of the light path, 1 cm. Substitution of 4, 7 and 8 into 6 gives
K a
12.55
A h M
(9)
When the Rayleigh formula is used, a may be computed directly by a8
Ahair
32.22 --
M
(10)
Curves of u / E , the observed electrophoretic mobility, were plotted against the zeta potential tor twenty-four values of K r . This large graph was then utilized to determine the zeta potential from the observed u / E and K r . The standard deviation of the recorded zeta potentials ’was of the order of 3 to 5 millivolts. The errors were always large near the isoelectric point of bhe sol.
Results and Discussionlo The range of pAg and pH covered in these experiments was experimentally limited. When the concentration of salts became so high that the majority of the current was carried by ions, the determination of the mobility of the colloidal particles became difficult as their electrokinetic velocity became quite small. The particles tended to drop out of the field of the microscope before they had traveled a sufficient distance for the determination of (IO) The photographic terms used in this paper are defined ae: optical sensitizer, a dye which extends the wave length sensitivity of a ailver halide; densitirer, a dye or oompound which lowers the intriaaic light sensitivity of a silver halide; chemicd sensitizer, a compound which increases the intrinsic light sensitivity of silver halide.
.
JAMESE. LUVALLEAND JUANITA M. JACKSON
1218
320 O
P
-
T
1 4
-40 -50
-70 I 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0 10.0 11.0 PA& Fig. 2.-Zeta potential of unprotected silver bromide sols versus pAg for several values of pH: @, pH 4.0; 0 , pH 5.5; n,pH 7.2; 0, pH 9.4; a, pH 9.2 and p = 0.01; 0,pH 10.3; , zeta potentials calculated from Langmuir adaorption isotherm equation of text; - - -, pH 9.4 no control of ionic strength; . . . . , pH 9.2, constant io& strength,
p
= 0.01.
their velocity. Above pH 6.0, the formation of silver oxide placed a lower limit on the values of pAg which could be utilized in these experiments. The effect of pAg on the zeta potential for constant values of the p H is shown in Fig. 2. The curve obtained from the data a t pH 5.5 is very similar to the theoretical curve of Davies and Halliday. 4b These data show excellent agreement with Davies and Halliday’s experimental data4b above pAg 5.0. However, the sharp peak reported a t pAg 3 by these workers was not observed in this investigation. The data a t pH 9.2 appeared to be anomalous when compared to the data at higher and lower values of pH. However, when the ionic strength was held constant a t 0.01, data a t pH 9.4 were found to fit in with the data observed at higher and lower values of pH. This demonstrates then that the zeta potential is dependent on the ionic strength of the solution, as well as the pAg, in agreement with the Stern theory of the double layer.” I n general, no attempt was made to hold the ionic strength constant because of the large excess of ions that had to be introduced into the solutions. It was felt that the introduction of “inert” ions might complicate the observed results. Furthermore, it is indeed difficult t o determine whether the ions are completely “inert,” ;.e., not adsorbed to some extent. Computation of ionic strength for each experiment would be extremely tedious and, as no attempt was made to hold the ionic strength constant, it did not appear worthwhile. However, the data obtained when pAg and pH were held constant while the concentration of adsorbate was varied were effectively at constant ionic strength. At pAg 3, Davies and Halliday reported a decrease in zeta potential with increase in pH between pH 6 and 7 and with decrease in pH below pH 2.5. At pAg 3, the sols were below the isoelectric (11) 0. Stern, Z. Elehbochsm., 80, 508 (1924).
Vol. 61
point where adsorbed Ag+ ions were present. The most plausible explanation of the drop in zeta potential below pH 2.5 is the increase in ionic strength with pH. The rapid fall off in zeta potential between pH 6 and 7 was probably due to the adsorption of hydroxyl ion by silver bromide. I n the present work at pAg 8.1, i.e., above the isoelectric point, the pH had no effect on zeta potential between pH 4 and pH ll, i.e., hydroxyl ion was not adsorbed by the silver bromide above the isoelectric point. As sensitizing dyes are used in the large majority of present day photographic emulsions, the effect of sensitizing dyes on the zeta potential of silver bromide sols was to have been investigated. However, of the dyes supplied to US,'^ only the cationic desensitizers pinakryptol green and pinakryptol yellow had sufficient water solubility for these experiments. Since it was necessary to add two to five milliliters of concentrated dye solution to a liter of sol (after stabilization and the measurement of the optical absorption), the dye had to be reasonably soluble in water. Pinakryptol green and pinakryptol yellow were investigated over the concentration ranges 1.00 X 10-8 t o 1.00 X M and 5.00 X lo-$ to 5.00 X 10-6M, respectively. At pAg 8.3, the zeta potential was independent of pH from pH 5 to pH 11 in the presence of these desensitizers. At pH 9.2, the zeta potential of 1.00 X M pinakryptol green and 5.00 X lo-’ M pinakryptol yellow was about -40 mv from pAg 6 to pAg 11. 1.00 X 10-6 M pinakryptol green showed a linear increase of zeta potential with pAg. When the ionic strength of 1.00 X 10-6 M pinakryptol green solution was held constant a t 0.01, the zeta potential was about - 5 mv. from pAg 6 to pAg 11. The effect of concentration of desensitizing dye on the zeta potential a t pH 9.3 and pAg 8.3 has been summarized in Fig. 3. The data indicate that pinakryptol yellow was more strongly adsorbed than pinakryptol green and that a t dye concentrations below lo-’ M , the effect of adsorbed dye on the zeta potential may be considered negligible for the sols used in these experiments. As quaternary salts have been used to accelerate the development processla it appeared pertinent to obtain some quantitative data on the effect of these salts on the zeta potential. Pinakryptol green and pinakryptol yellow may be considered as quaternary salts and as desensitizers. They were classified as desensitizers because they are commonly used for this purpose. Of the two substances classified here as quaternary salts, phenosafranin and laurylpyridinium bromide, the former compound is also a desensitizer. However, both of these compounds have been used frequently by James as quaternary salts in his investigations of the acceleration of photographic development,12 (12) Dr. J. A. Ieermakers of the Kodak Research Laboratories supplied us with the following dyes: (1) 3,3’-diethylthiacarbocyaninep-toluenesulfonate, (2) 3.ethyl-5-(l-ethyl-2-quinolylidene-ethylidene)rhodanine, (3) erythrosin, (4) pinakryptol green and (6) pinakryptol yellow. (13) (a) T. H. James, P.S.A. J. Phot. Sci. and Tech., 178, No. 9, 73 (1951); (b) T. H. James, J . Franklin Inst., 240, 15 (1945); (0) 240, 83 (1945); (d) 240, 229 (1945); (e) 240, 235 (1945); (f) 248, 327 (1947).
.
Sept., 1957
THEZETA POTENTIAL OF SILVER BROMIDE
hence phenosafranin was classified as a quaternary salt for this work. The data showed that, within the experimental error, the zeta potential of 1.0 X M phenosafranin a t pH 9.5was -36 mv. between pAg 7 and ll, whereas 1.0 X lov5M laurylpyridinium bromide at the same pH had a zeta potential between -20 and -30 mv. over the same pAg range. Figure 4 shows the effect of concentration of quaternary salt on the zeta potential a t pH 9 and pAg 8. Phenosafranin appears to be more strongly adsorbed than laurylpyridinium bromide. Thiourea was chosen as an example of a chemical sensitizer. Data were obtained over a range of concentrations and pAg a t pH 4. It was necessary to make the measurements at pH 4 in order to prevent reaction of the thiourea and the silver bromide. At 1.00 X M thiourea, between pAg 6 and pAg 11, the silver bromide sol showed the same variation of zeta potential with pAg as a sol prepared without thiourea. At 1.00 X M thiourea, the zeta potential was almost constant over the same pAg range. At intermediate concentrations, the zeta potential decreased with increase in pAg over the same pAg range. Figure 5 summarizes the effect of concentration of thiourea on the zeta potential a t pH 4.00 and pAg 7.0,8.2 and 10.2, respectively. The curves resemble adsorption isotherms. The three curves show that at pAg 7,8.2 and 10.2,the isoelectric point is reached at about 3.0 X and 5.0 X loU6M thiourea, 3.0 X respectively. The curves at pAg 8.2and 10.2indicate that, at the lower pAg, most of the available sites for bromide ions are already occupied. Increase in pAg does not cause much additional adsorption of bromide ions. Obviously, information regarding the effect of photographic developing agents on the zeta potential of silver bromide sols is of importance. However, it is necessary to obtain this information under conditions whereby the photographic developing agent does not react with the silver bromide grain to produce silver. Therefore, the zeta potential measurements were made a t considerably lower values of pH than the pH values used in processing film. Hydroquinone was not appreciably adsorbed between pH 4 and pH 7. As this is the range in which the molecule is neutral and also photographically inactive, it appears that the neutral molecule was not appreciably adsorbed by silver bromide. Phenidone (1 -phenyl-3-pyrazolidone) was only slightly adsorbed, while Elon (N-methyl-p-aminophenol) and 2-amino5-diethylaminotoluenewere adsorbed to some extent. Figure 6 shows the effect of the concentration of the latter two agents on the zeta potential. The diamine was more strongly adsorbed than the aminophenol. The presence of an adsorbed layer of phenosafranin did not increase the adsorption of hydroquinone or phenidone; however, the diamine was adsorbed. These data appear to imply that a t the pH of development, the monovalent and divalent anions must be the adsorbed species of hydroquinone, At the present time, most photographic emulsions are prepared in about 10% gelatin; hence, it is of importance to know the effect of gelatin on the
1219
40 1
I
p' I I I I
-80'
I
'
10-7 10-6 10-6 Concn. (moles/l. ). Fig. %-Effect of concentration of desensitizing dye on the pinakryptol zeta potential at PH 9.3 and pAg 8.3 -e, green; -763--, pinakryptol yellow; -e-, pinakryptol yellow, pH 11. 10-8
r-
7 -
1-
-70
-------
I-
d
-80 10-7 10-6 10-5 10-4 Concn. (molea/l.). Fig. 4.-Effect of concentration of quaternary salt on the zeta potential of silver bromide sola at pH 9 and pAg 8: 0 -, laurylpyridinium bromide; --e, phenosafranin. 10-8
-
c:
40 -
; 20-
v
3 Y
0-
;-20 0
a 01
42
-
-40 -
-60
-
10-7
10-6
10-6
10-4
Concn . (moles/l. ). Fig, 5.-Effect of concentration of thiourea and pAg on zeta potential of silver bromide sols held at pH 4.0: -e, - -,pAg 10.2. pAg 7.0; - A -, pAg 8.2; - - -
-
zeta potential of silver bromide sols and the effect of dyes, quaternary salts, etc., on the zeta potential in the presence of gelatin. As a gel containing 10% gelatin was much too viscous to utilize in this investigation, it was necessary to work with solutions between 0.001 and 0.1% in gelatin. The data of Fig. 7 show that there was little observable difference in the zeta potential of a silver bromide sol in 0.1 and 0.01% gelatin solutions. However, a t 0.001% gelatin, the colloidal particles were apparently incompletely covered by gelatin. Figure 8
1220
JAMES E. LUVALLEAND JUANITA M. JACKSON v
Vol. 61
illustrates the effect of pH on silver bromide sols in the presence of 0.01 and 0.1% gelatin solution. The / decrease in zeta potential with increase in pH may I be attributed to the ionization of free carboxyl groups on the gelatin but not to the adsorption of - -I01 hydroxyl ion by the silver bromide. The effect of pAg on the zeta potential of silver bromide sols in 0.01% gelatin is demonstrated in Fig. 9. At pH 4.1, the zeta potential was positive and fell off toward zero as the pAg was increased. At pH 5.1, the experiments were near the isoelectric point of 1 the gelatin and a broad minimum was found in the i zeta potential pAg curve. At pH 9.5, the zeta po- - -9 - tential increased almost linearly with increase i ri pAg. The observed decrease in zeta potential with 10-6 1010-4 io-* io-* increase in pAg at pH 4.1 may be caused by the reConcn. (moles/l.). Fig. 6.-Effect of concentration of developing agent on moval of silver ion complexed with gelatin because the zeta potential of silver bromide sols at pH 4.0 and pAg of the increase in bromide ion concentration 1
i
h
2
-
I
8.3: - -e- - Elon; toluene hydrochioride.
Q
% -50.
2-amino-5-diethylarnino-
-A-,
I
N
L
0.001
0.01
0.1
Gelatin concn. (%). Fig. 7.-Effect of gelatin concentration on the zeta potential of silver bromide sols, pAg 8.3, pH 9.5. 20
x
k 4
I
I
10
v
0 .* + s -10
& -20 a 8 -30 Y
t
7 8 9 10 PH. Fig, %-Effect of pH and gelatin concentration on the zeta otential of AgBr sols at pAg 8.4: 0, 0.01% gelatin, A, 0.18gelatin.
5
4
6
20 I
x
-E
10 -
0-
*$ -10 6
8
I
-
""...$,
......'".......,,,
..............................
.
-20 -
-30 3 -30
-
-40 -
3 4 5 6 7 8 9 1 0 1 1
PAg.
Fig. 9.-Effect of pAg on the zeta potential, 0.01% gelatin-protected sols: -e-, pH 4.1; ---A---, pH 5.2; 4 pH , 9.5.
Gel=NH.Ag+
+ Br- +Gel="
+ AgBr
The linear increase in zeta potential with increase in pAg a t pH 9.5 may not be explained as easily. The pH was held constant; hence additional ionization of carboxyl groups or additional proton addition did not take place. As the pAg increased, both bromide ion and potassium ion concentration increased. Adsorption of bromide ion by the silver bromide or the gelatin would have given an effective decrease in the zeta potential rather than an increase. Further removal of silver ion complexed with the gelatin would have also lowered the zeta potential. However, the concentration did increase with increase in pAg. If the potassium ion were complexed by gelatin or combined with the carboxyl anionic groups, the zeta potential would increase with pAg. As this behavior was not observed with gelatin-free silver bromide sols, the potassium ion was not adsorbed on the silver bromide surface. It therefore appears that a t sufficiently high potassium ion concentrations, the gelatin complexes with potassium ion or potassium ion was adsorbed a t the gelatin-solution interface. The broad minimum observed near the isoelectric point of gelatin at pH 5.1, if not due to experimental error, may be explained by a loss of silver ion, with resultant lowering of the zeta potential, followed by a gain of potassium ion, with resultant increase of the zeta potential. The curves also indicate that the complexing or adsorption of silver ions and potassium ions by the gelatin is a function of the pH. The data a t hand were insufficient to determine the exact relationship, The data on the effect of pH and pAg on the zeta potential of gelatin-protected silver halide sols imply that the experiment now observes the adsorbed gelatin sheath-solvent interface rather than the silver bromidesolvent interface. The data also imply that the adsorbed gelatin sheath must be at least an order of magnitude thicker than the double layer at the silver bromidegelatin interface or that the gelatin was so strongly adsorbed that nothing else could be adsorbed on the silver bromide surface. It will be shown that the latter conclusion is probably invalid. It should be pointed out that the silver bromide-gelatin interface is usually complicated by the presence of soli vent molecules and is in reality a mixed interface in
Sept., 1957
THEZETA POTENTIAL OF SILVER BROMIDE
which gelatin and solvent probably occupy adjacent sites on the silver bromide surface. If the experiment now observes only the interface of adsorbed gelatin sheath and solvent, sensitizing dyes and quaternary salts should not alter the zeta potential of a sol in 0.01% gelatin unless the dye was adsorbed to the gelatin or displaced the gelatin from the grain. The data show that pinakryptol green had virtually no effect on the zeta potential of the sol in the concentration range from to M and that phenosafranin exhib1cf concentrations. ited an effect only above At M phenosafranin, the zeta potential was lowered. This is difficult to understand as a M phenosafranin solution increased the zeta potential of an unprotected sol to zero (Fig. 5 ) . Furthermore, if the quaternary salt adsorbed to the gelatin, it would have been expected to raise the zeta potential. The observed decrease may not be explained by adsorption of the quaternary salt to the silver bromide or to the gelatin. JamesI3 concluded from his investigations with certain photographic emulsions that quaternary salts were indeed adsorbed to the silver bromide grain and that this adsorption increased the zeta potential which thereby decreased the potential hump that anionic developer ions had to surmount to reach the silver bromide grain. MitchellI4 has suggested the quaternary salt acts to lift up the gelatin from the grain, thus making the grain surface more accessible to the developing agents. If James' conclusion is correct, the quaternary salt must be adsorbed on the silver bromide surface and hence the gelatin must not be strongly adsorbed to the silver halide grain. If Mitchell's conclusion is correct, the quaternary salt may perhaps only loosen the gelatin binding by complexing with the gelatin and not actually adsorb to the surface of the grain. It has been demonstrated that sensitizing dyes may displace gelatin from silver bromide surfaces.'S The heat of adsorption of merocyanine B on bromide grains in the absence of gelatin was found to be 12.2 kcal. while the heat of adsorption on grains formed in seven per cent. gelatin was 3.0 kcal./mole which indicates that gelatin was adsorbed to silver bromide with a heat of adsorption of about 9.0 kcal./mole. Micro-Kjeldahl determinations of nitrogen showed that gelatin was displaced by sensitizing dye.16 These data show that the dye may displace the gelatin; however, as the grain retains a gelatin sheath, the dye must tend to lift the gelatin from the surface of the grain a t points where it is most weakly adsorbed. The shapes of the zeta potential-concentration curves resemble closely adsorption isotherms. The charge on the grain is proportional to the number of ions a d ~ o r b e d . ~If the assumption is made that the zeta potential is proportional to the charge14the zeta potential is also proportional to the number of ions adsorbed. It is fully realized that this is only an approximation, but it does permit some approximate correlations with adsorption isotherms. It is (14) J. W. Mitchell-Talk given at Internstionsl Symposium on Photography in Cologne, Germany, September, 1956. 0 5 ) W. West,B. H. Carroll and D. L. Whitcomb, J . Phot. Sci., I, 145 (1963).
1221
then permissible to assume that the zeta potential will be related to the fraction of the surface covered by adsorbed ions by (11)
where {o, t i and Imare the zeta potentials in the absence of adsorbed substance, the zeta potential a t concentration c and the zeta potential when the surface is completely covered. The Langmuir adsorption isotherm may be rewritten as
where k is a constant and c is the equilibrium concentration when 6 is the fraction of surface covered. The constant, k , was computed from the data for thiourea a t three values of pAg and of laurylpyridinium bromide a t one pAg value. These computal e siltions gave a value of k = 3 X 108 ~ m . ~ / m o for ver bromide. This value of k was then used to compute the zeta potential of silver bromide sol in the presence of bromide ion by combining equations 11 and 12
was taken as 30 mv., the zeta potential a t pAg 2 where only Ag+ ions are adsorbed and trnwas taken as -60 mv., the mean value a t pAg 10. The empirical curve computed by equation 13 has been plotted in Fig. 2 as the solid curve. It evidently fits the experimental data quite well. These calculations establish that the zeta potential data obtained in this paper result from an adsorption process. The data of this paper show that, in the absence of gelatin, it is possible to study the effect of many substances on the zeta potential a t a silver bromidesolvent interface. These data also show that, in the presence of gelatin,' the experiment only observes the gelatin-solvent interface and hence no conclusions may be drawn as to the zeta potential at the silver bromide-gelatin interface. It is obvious from the results that any conclusions as to the zeta potential a t the silver bromide-gelatin interface are open to considerable question for gelatin, chemical sensitizers, optical sensitizers and bromide ion may all be adsorbed on part of the silver bromide grain. The presence of all of these agents also indicates that quaternary salts probably do more than merely alter the zeta potential a t the silver bromide surface. It appears that these salts may: (1) effectively loosen or desorb gelatin, dyes, etc., adsorbed to the silver bromide surface by a wetting or detergent action; (2) be adsorbed t o the grain surface and (3) perhaps complex with developer molecules and thereby effectively lower the potential barrier the developer molecule must overcome in approaching the grain. Mitchell14is probably correct in assuming that the availability or non-availability of adsorption sites at or near the latent image may determine the rate of development. James and Vanselow have shown that the adsorption process must play an important role in the photographic
1222
B. RICE, R. J. GALIANO AND W. J. LEHMANN
development process. I n recent papers,ls they have discussed the effect of ionic charge on the adsorption process and have concluded cooperative adsorption processes may be important. This paper clearly shows that adsorption and desorption processes must be very important in the photographic development process. It also indicates that the zeta potential in most photographic emulsions must be considerably less than - 60 millivolts for it is indeed doubtful that the bromide ion completely displaces gelatin and sensitizing dyes (16) T. H. James and W. Vanselow, J. A m . Chem. Soc., 74, 2374 (1952); THIS JOURNAL, 67, 725 (1955); 68, 894 (1954).
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from the emulsion under normal processing conditions. The observation that the zeta potential a t the gelatin sheath-solvent interface increases with pAg because of (we believe) the adsorption of potassium ion also complicates the discussion of the effect of charge on the development process. This paper indicates that an investigation of the adsorption processes and their role in the development mechanism should be undertaken. Obviously the net zeta potential must be a function of the relative strength of adsorption of dye, gelatin, bromide ion, and other species present in the emulsion and the processing solution.
VIBRATIONAL SPECTRA AND FORCE CONSTANTS OF TRIMETHYLAMINE-BORANE AND TRIMETHYLAMINE-BORANE-& BYBERNARD RICE,ROBERTJ. GALIANO AND WALTER J. LEHMANN Department of Chemistry, St. Louis University, St. Louis, Mo. Received May 37, 1067
The Ranian and infrared spectra of trimethylamine-borane and trimethylamine-borane-& were measured and the fundamental frequencies of both molecules assigned. A force constant calculation was made based on a n!e3NBHk (eight part$:) model with the most important interartion terms included in the potential function. The trends in certain characteristic frequencies of some ether and amine-boranes were correlated with the electronic charge distributions and the stabilities of the complexes.
Introduction The interest in the field of addition compounds formed by Group I11 Lewis acidswith Groups IV, V and VI Lewis bases centers at present on the relative stabilities of these complexes. To a first approximation, the relative stabilities can be predicted from a consideration of (1) the position of the electron donor and acceptor atoms in the periodic system and (2) inductive, steric and resonance effects. Other aspects of modern valence theory have recently been invoked to account for the exceptions to the order predicted from the Above factors.' Vibrational spectroscopy can play an important role in this field. From a comparison of the vibrational spectrum of the complex with the spectra of the parent acid and base, detailed information can be obtained on the structural changes that accompany the formation of the addition compound. Furthermore, the force constant of the donoracceptor bond should be related to the stability of the complex. Unfortunately, because of the complicated nature of the molecules or the incompleteness of the spectra, the force constant calculations in most of the reported work are based on molecular models and potential functions too simple for such a relationship to be investigated. I n the present work we have obtained the Raman and infrared spectra and made an assignment of all the observed vibrational frequencies of trimethylamine-borane, (CHs)3NBH3, and its isotopjc analog trimethylamine-borane-&, (CH3)3NBDs. Our normal coordinate treatment, based on an MesNBHs model, enabled a more complete calculation of force constants than is common in (1) W. A. G. Graham and F. G . A. Stone, J . Inorg. NucE. Chsm., 8 , 164 (1958); D. L-I Mclhniel, Science, 126, 545 (1957).
this field. Furthermore, with the frequencies of the isotopic derivative available, the most important interaction terms could be retained in the valence force potential function. Experimental Apparatus.-The infrared spectra were obtained with a Perkin-Elmer Model 21 spectrometer by Dr. William Elliott of the Biochemistry Department, St. Louis University. The Raman apparatus has been previously described in the literature.2 Chemicals.-Diborane (obtained from the Rev. F. J. Koenig of the Chemistry Department) was purified by fractionation. Trimethylamine was dried and purified by reaction with lithium aluminum hydride.8 Diborane-dd was prepared by successive exchange of BzHa with deuterium gas. Other reagents were dried and purified by standard means before use. Two methods were used for the preparation of trimethylamineborane as a check against possible spurious lines in the spectra due to decomposition products. Direct combination of diborane and trimethylamine was first used. Later the compound was prepared by the reaction of trimethylammonium chloride with lithium borohydride in ether slurry.4 Spectra obtained on both preparations agreed in every detail. Trimethylamine-borane-& was prepared by direct addition of the amine and BzDeonly. Measurements .-The Raman and infrared measurements were made on solutions of the addition complexes. The spectrum of the pure complex was determined by subtracting the solvent spectrum. The Raman spectrum of trimethylamine-borane was obtained from solutions in three different solvents-tetrahydrofuran, carbon tetrachloride and diethyl ether. Since this complex is most soluble in tetrahydrofuran, the spectrum obtained from this solution was the most complete. However, four additional lines, obscured by tetrahydrofuran solvent lines, were discovered by the meaRurements on a carbon tetrachloride solution. The measure, ments on the diethyl ether solutions led to no further lines. (2) B. Rice and H. S. Uohida, THIS JOURNAL, 69, 650 (1956). (3) J. Roscoe, Ph.D. Thesis, St. Louis University, 1954. (4) G. W. Schaeffer and E. R. Anderson, J . Am. Chem. Soc. 71, 2145 (1949).
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