Oct., 1958
I~ATES OF SOLUTION OF Y I L V ~ R HALIDJG EMULSION CXYSTALS
1189
RATES OF SOLUTION OF SILVER HALIDE EMULSION CRYSTALS BYT. H. JAMXSAND W . VANSELOW Communication No. 1945 from the Kodak Research Laboratories, Eastman Kodak Company, Rochester, N . Y . Received March 3, 1969
A method is described for determining the rates of solution of photogra hic silver halide grains in various solvents. The rate of solution of pure silver chloride in potassium chloride solution &creases to a minimum and then increases as the chloride ion concentration increases. The rate of solution of silver chlorobromide, silver bromide and silver iodobromidu in potassium bromide solution likewise passes through a minimum and then increases as the bromide ion concentration increases. The rate of solution of silver chlorobromide, bromide and iodobromide in potassium chloride solution illcreases almost linearly with increasing chloride ion concentration and shows no miminum. The rate of solution of silver bromide and iodobromide in the presence of sodium sulfite passes through a minimum as the bromide ion concentration increases, in contrast to the behavior of the equilibrium solubility. The dependence of the rate of solution of silver bromide upon the specific area of the solid differs according to the solvent used. The rates are sub-proportional to surface for sulfite as solvent, proportional for bromide and thiocyanate, and super-proportional for ammonia. The mechanism of solution is discussed with respect to the ion complexes involved, the degree of perfection.of the grain surface and surface charge eflects.
The equilibrium solubility of silver chloride and silver bsomide in excess halide, sulfite and other solvents has been determined by a number of investigators (cf. Hill, Zuehlke and Ballard, ' Klein,2n3 Chateau and H e r ~ i e r ,and ~ references cited by them). A knowledge of the solubility of the silver halides and of the composition of the silver ion complexes formed is of importance to the theory of photographic emulsion preparation, development and fixation. However, the rate of solution is often a more significant factor than the equilibrium solubility, and the equilibrium studies do not give direct information on the rate. We have described a method6 which can be used to determine the rate of solution in a developer of silver halide grains coated in a photographic emulsion. By proper choice of the developer, we can use this method to determine the rate of solution of the grains in various' solvents such as sulfite, excess halide and ammonia. The method depends upon the fact that the silver halide grains can be mixed with a Carey-Lea-type silver sol, prepared by dextrose reduction of silver nitrate in the presence of gelatin, and subsequently coated onto film support without thereby making the silver halide grains developable. The unexposed coating is then irnmersed in the solvent solution which contains a developing agent which itself has no solvent action and which is active enough to reduce silver ions as fast as they arrive a t the colloidal silver nuclei. The latter condition is established when the measured rate of accumulation of silver is independent of the concentration of the developing agent, since the specific rate of reduction of silver ions at a silver surface is markedly dependent on the concentration of developing agent. The measured rate could still depend upon the rate of diffusion of silver ions from the silver halide surface to the colloidal silver surface, and it actually does if the concentration of colloidal silver particles is small enough or if the distance between grain and silver nuclei is great enough. However, a t the concenttration used in most of our experiments, where the number of silver particles exceeded the number of (1) M . A . Hill. C. W . Zuehlke and A. E. Ballard. Phot. Sri. Tech.. I111 8, 101 (1956). (2) E. Klein, 2. Elaktrochem., 60, 1003 (1966). (3) E. Klein, Phot. Korr.. 92, 139 (1956). (4) 13. Chatmu and R. Hervier, J . chim. phys., 54, 637 (1957). (5) T. H. James and W . Vanselorv. Phot. Sci. Tech., 1111 2, 135 (1955).
silver halide grains by several orders of magnitude, diffusion is not important. This is indicated by the lack of correspondence between the measured rates and the equilibrium-solubility and by the fact that the temperature-dependence of the measured rates is much higher than that of either the equilibrium solubility or a normal diffusion process. Further evidence will he presented subsequently, when it will be shown that the measured rate is independent of the average distance between the silver particles and the silver halide grains for the concentrations of colloidal silver used in the present work. In this paper, new data are reported on the rates of solution of silver chloride, silver bromide, silver chlorobromide and silver iodobromide grains in excess chloride and bromide ion solutions, the effect of sulfite on the rate of solution of silver bromide in the presence of various excesses of bromide ion, and the dependence of the rate of solution on the size of the silver bromide grains for several solvents. Experimental Materials and Procedure.-The photographic emulsions used consisted of fine- and medium-grain-size silver chloride (Fig. l), fine- and medium-grain-size silver bromide (Fig. 2), a fine-grain 50-50% silver chlorobromide, a finegrain silver iodobromide containing a fraction of a per cent. of iodide, and three silver bromide emulsions of relatively narrow grain-size spread (Fig. 3). The grain size of the latter three emulsions was varied by changing the duration of physical ripening during the emulsion preparatidn, and the specific surfaces, as determined by dye adsorption, were 17,800, 6900 and 2800 sq. cm. per gram of silver bromide. The emulsions were coated at a silver halide spread corresponding to 35-40 mg. halide silver per 100 sq. cm. and contained approximately 0.48 mg. colloidal silver per 100 sq. cm. A normal amount of gelatin was used in these coatings, which were formaldehyde-hardened. A series of special coatings of the silver iodobromide grains was also made with a substantially constant halide silver spread (see Table I ) and a spread of 0.48mg. colloidal silver per 100 sq. cm. The silver-to-gelatin ratios in this series were 2.70, 1.35, 0.68, 0.43 and 0.32. The general procedure for the rate determinations has already heen described6 and the formula for the Metol-ascorbic acid developer published. In the experiments a t low p H , a developer composed of tetrasodium ethylenediaminetetraacetate, 43.2 g.; ferrous sulfate, 6.22 g.; and potasaium bromide, 1 g. per liter, was used and the p H was adjusted b,y addition of acetic acid. All determinations were at 20.0 . The rates of solution are expressed in terms of the change T.H. James and W. Vanaelow, ibid., [I11 8, 69 (1958). (7) T. H. James and W. Vanaelow. Phot. Eno.. 7, 90 (1950).
(6)
1190
T. H. JAMES AND W. VANSELOW
Fig. 1.-Photomicrographs
Vol. 62
of silver chloride grains, 1780 X magnification.
. Fig. 2.-Photomicrographs of silver bromide grains, 1780 X magnification. in optical total density per minute of the colloidal silver solved. This “induction period” does not occur for other measured a t the absorption maximum (approx. 430 mp), solvents. The rates reported in the present paper are those which is essentially equivalent to the amount of silver in determined in the constant rate region. milligrams per 100 sq. cm. per minute. The density at the maximum is proportional to the amount of silver, and the Results proportionality constant under the conditions used is 0.99, Dependence of Rate on Gelatin Concentration.i.e., Ag = 0.99 D,when silver is expressed in milligrams per 100 sq. cm. The density was corrected for any directly re- By keeping the amount of silver halide and colloidal duced silver as previously describeda6 The rates of solution silver coated per unit area constant and varying in some solvents increase a t first with increasing time, but soon reach a substantially constant value which remains the gelatin content, it is possible to change the constant until about 25% or more of the halide has dis- average distance between the silver halide grains
.
Oct., 1958
1191
RATESOF SOLUTION OF SILVER HALIDEEMULSION CRYSTALS
Fig. 3.-Photomicrographs
of silver bromide grains, 1000 X magnification.
3.4
16
3.0 14
2.6 1
12 I
\
.E= E \
E
D
DE
Q
1. 1.8
c
B
m
c
a
g
m
IO
m
808
3
1.4
I0V/
06
04
06
O
2
02
20
2
KCI (g/l),
Fig. 4.-Effect of excess chloride ion on the rate of solution of silver chloride: 0, fine grain; 0 , medium grain.
and the colloidal silver nuclei. Table I gives data for the silver iodobromide grains and various developers. The Metol-ascorbic acid developer at pH 10.2 (M-AA-1) and the iron ethylenediaminetetraacetic acid developer a t pH 4.5 (Fe-EDTA) both contained 1.0 g. of potassium bromide per liter and no sulfite. The Kodak Developer D-19 contains 96 g. of sodium sulfite and 5.0 g. of potassium bromide per liter. The data show little or no dependence of rate upon the silver-to-gelatin ratio. This result, together with the facts previously cited, shows that neither the specific rates of reduction of silver ions nor of diffusion have a significant in-
0
I 4
I
I
8
12
I 16
I
I
20
KBr (g/l),
Fig. 5.-Effect of excess bromide ion on the rate of solution of silver bromide and silver chlorobromide: 0, 50-50y0 chlorobromide; A, fine-grain silver bromide.
fluence on the measured rates under the conditions used. Hence, we conclude that the measured rates are essentially those of solution of the silver halide. The data further indicate that p H has little effect upon the rate of solution in the bromide solution which contains no sulfite over the p H range of 4.5 to 10.2. This has been confirmed by data obtained a t intermediate p H values. Effect of Excess Halides on Rate of Solution.Figure 4 shows the effect of various concentrations of potassium chloride on the rate of solution of silver chloride and Fig. 5 shows the effect of potassium
T. H. JAMESAND W. VANSELOW
1192
Vol. 62
The general shape of the curves is the same, and resembles that previously obtained for the effect of bromide on the rate of solution of the silver iodobromide grains.b In each case, the rate decreased to a minimum as the halide ion concentration increased, then rose again with further increase in halide-ion concentration. The rate of solution thus follows the same general course as the equilibrium solubility, but the two are not proportional. Figure 6 shows the effect of potassium chloride concentration on the rates of solution of pure silver bromide grains and of the silver chlorobromide grains. The rate of solution of each varied almost linearly with the chloride ion concentration, and there is no evidence of a minimum. The rate of solution of silver chloride was decreased markedly by small concentrations of potassium bromide, and the rate of solution in 0.2 g. of potassium bromide per liter was less than 0.1 that in water. The rate then increased as the bromide ion concentration increased beyond 0.2 g. of potassium bromide per liter, and at higher bromide concentrations the rate of solution of the fine-grain silver chloride became about equal to that of the fine-grain silver bromide. Sodium fluoride had only a small effect on the rate 6 of solution of silver chloride. The rates of solution I of the fine-grain silver chloride in solutions of 0, 5, I I I . I I 0 4 8 12 16 20 20 and 40 g. of sodium fluoride per liter were 1-00, KCI WI). 1.00,0.90 and 0.83, respectively. Fig. 6.-Effect of chloride ion on the rate of solution of Figure 7 shows the effect of bromide on the rate silver bromide and silver chlorobromide: 0,50-50% of solution of silver bromide grains in the presence chlorobromide; A, fine-grain silver bromide. of sulfite. The rate curve is similar in form to that obtained in the absence of sulfite, except that the minimum is displaced to somewhat higher bromide concentrations. The rates of solution at pH 9.5 were the same, within the limits of experimental error, as a t pH 10.4. The rate curve in the presence of sulfite does not correspond a t all in form t o the equilibrium-solubility curve obtained by Hill, Zuehlke and Ballard,' which shows no minimum up t o 10 g. of potassium bromide per liter. This curve is also reproduced in the figure. Similar results were obtained in measurements of the rates of solution of the silver iodobromide grains in the presence of sodium sulfite. A distinct minimum was obtained at 1-1.5 g. of potassium bromide per liter, both for solutions of 20 and 60 g. of sodium sulfite per liter, and further increase in bromide concentration caused a rapid increase in rate. KBr [ g l l ) . Fig. 7.-Effect of excess bromide ion on the rate of soluThallous ion increased the rate of solution of tion and equilibrium solubility of fine-grain silver bromide silver bromide grains in water and in dilute broin the presence of sodium sulfite: 0,rate of solution,. no sulfite. A, rate of solution, 20 g. Na&O,/l.; 0, equilibrium mide ion solutions. Thus, the rate of solution of the fine-grain silver bromide in 0.5 g. of thallous solubiiity, 20 g. NaaSOs/l. nitrate per liter was 1.2 times the rate in pure TABLEI water. The same amount of thallous nitrate added E F F ~ COFT SILVER/GELATIN RATIOON RATEOF SOLUTIONt o a solution of 0.2 g. of potassium bromide per liter Halide increased the rate of solution of the silver bromide silver spread mg./100 Bilver Max. rates of s o h in by 2.5-fold. This effect of thallous ion is in the sq. em. Gelatin M-AA-I Fe-EDTA D-19 same direction as that previously found for quat,er48 2.70 0.080 0.062 .. nary ions,b and the mechanism of the acceleration 42 1.35 ,080 .072 0.44 probably is the same for both. 42 0.68 .090 .077 .47 Effect of Grain Size on Rate of Solution.-The 43 .43 .080 .070 .45 effect of grain size on the rate of solution of silver 44 .32 ,080 .074 .47 bromide was determined for three grain preparabromide on the rate of solution of pure silver bro- tions which differed only in the duration of physical mide and of a 50-50 mole % silver chlorobromide. ripening. The grains were coated at the same '
RATESOF SOLUTION OF SILVER HALIDEEMULSION CRYSTALS
Oct., 1958
weight of silver bromide per unit area. The rates of solution in several solvents are listed for each emulsion in Table 11. Rate determinations for the largest grains (smallest specific area) were uncertain in some solvents because of direct attack of the developer on the silver bromide grains. In such instances, the true rate should lie between the limits given. This difficulty was not encountered for the two smaller grain emulsions, however. The variation of rate with respect to surfqce clearly depends on the solvent used, The rate is superproportional to surface area with ammonia as solvent, roughly proportional to area with KBr and KCNS as solvents, and sub-proportional with sodium sulfite as solvent. Data for several commercial developers, Kodak Developers D-19, DK-50 and D-76, and for two of these developers with the addition of KCNS are plotted in Fig. 8. For all of these sulfite-containing developers, the rate is sub-proportional to Furface. EFFECTO F GRAINSIZEO N RATEOF SOLUTION OF SILVER BROMIDE: METOGASCORBIC ACIDDEVELOPER
+
1.6
I.4 A'
r
'E
>1.2 0
9
2 1.0
o="
Q0.8
m 0.6
0.2
D/min. rate for specific areas in sq. am. of 2800 6900 17,800
0.03 M NHs 0.0083 M KBr 0.08 0.185 0.72 0.10 M NHs 0.0083 M KBr . 1.0 3.9 0.167 M KBr 0.12-0.21 0.39 0.96 0.0083 M KBr 0.0103 M KCNS 0.11- -18 0.35 0.96 0.0083 M KBr 0.0206 M KCNS 0 . 3 - . 5 1.01 2.76 0.77 M Na2SOI 0.23 0.34 0.48 0 . 7 4 M N a 2 S 0 4 + 0 . 0 4 M K B r 0.13 0.21 0.39 1.1 2.65 Above, pluR Q salt" 0.54 Q salt = 2 g. of a-picolinium-p-phenylethyl bromide per liter.
.
+
1.8
0.4
TABLE I1
Solvent
1193
+ +
Discussion The rate-of-solution data given in this paper apply to silver halide grains embedded in gelatin. Unfortunately, the method used to determine the rates is not applicable in the absence of protective colloid, but some qualitative experiments indicate that the gelatin significantly decreases the rate of solution. Data by various authorss-'2 on the rate of physical ripening of silver halide grains in the presence of various colloids likewise point to an important effect of the colloid on the rate of solution. The general form of the curve which relates the rates of solution of silver chloride to the concentration of excess chloride ions, and the similar curves for the rates of solution of silver chlorobromide, silver bromide and silver iodobromide in excess bromide ion, can be explained as follows. In the absence of excess halide or solvent other than water, the silver halide dissolves largely by passage of solvated silver and halide ions into the solution. (8) Y. Oyama, Proc. Roy. Phot. SOC.Centenary Intern. Confer. on the Science and Applications of Photography, London, Sept. 19-25, 1953, R. 6. Schultze (editor), Roy. Phot. SOC.1955, p. 37; Y. Oyama and K. Futaki, Sei. et. mds. phot.. [21 27, 41 (1956); Bull. Chem. Soe. Japan, 28, 243 (1955). (9) A. Narath and H. Gernert. 2. W ~ S ~Phof., R. 60, 204 (1965). (10) F. Evva, ibid., 52, 1 (1957). (11) V. V. Jones, ibid., 50, 138 (1955). (12) H. Ammann-Brass, ibid., 60, 173 (1956).
0
4
8
12
16
20 xl000
Totol Grain Area per Gram AgBr. Fig. %-De endence of rate of solution on grain area: 0,D-76; A, 8-19; 0,DK-50; A, D-19 -I-1 g. KCNS/l.; n, DK-50 1 g. KCNS/l. (All Kodak Developers.)
+
The activation energy of this process for' the silver iodobromide grains is approximately 18 kcal. per m0le.5 When a small amount of excess halide ion is present with the corresponding silver halide, some of the excess becomes adsorbed by the surface. The silver ion activity is decreased and the negative surface charge increases. Both factors tend to decrease the rate a t which silver ions can pass into solution. The activation energy of solution of the silver iodobromide increases to 20 kcal. per mole, and some molecular solution may occur.'* When larger amounts of excess halide ion are present, complex formation occurs. I n solution, complexes of the type AgBr5-4, AgCl2-I, AgC14+, etc., have been r e p ~ r t e d . ~ ~It ~ Jseems s probable that the complex involved in promoting solution of the silver halide is AgX2-l. The rate of solution of AgCl in KC1 beyond the minimum is nearly a linear function of the excess chloride concentration. The rates of solution of the silver chlorobromide, bromide, and iodobromide in excess bromide ion increase more rapidly than the bromide-ion concentration, but considerably less rapidly than the square of the bromide concentration. Moreover, the activation energy for solution of the iodobromide in 20 g. of potassium bromide per liter and in 200 g. of potassium bromide per liter is substantially the same,5 16.4 kcal./mole. The higher complexes indicated by the solubility data probably form in the solution rather than on the surface of the silver halide, The rate of solution of silver bromide in potassium chloride varies almost linearly with the (13) Cf.J. H. Jonte and D. 8. Martin, J . A m . Chcm. SOC.,74, 2052 (1952); K. H. Lieser. 2. anorg. u. a2lgem. Chem., 29'2, 97 (1957).
1194
T. H. JAMES AND W. VANSELOW
potassium chloride concentration. The complex involved in promoting solution probably is AgBrC1-’. Chateau and Hervier14 have found evidence for the existence of AgC1Br3+ and AgC13Br3 in solution. The latter complex, which would dominate in the more concentrated chloride solutions, probably forms in the solution rather than on the surface, and contributes t o the equilibrium solubility but not to the rate of solution. Sodium sulfite markedly increases the equilibrium solubility of silver bromide, but its effect on the rate of solution is comparatively small. The rate of solution in the sulfite is smaller by about two orders of magnitude than the rate of solution in an equal concentration of ammonia,b although the equilibrium solubility in the sulfite is greater. The formation of the complex Ag(S03)2-3 on the surface of the silver halide would be impeded by the negative charge of the halide surface, whereas formation of the complex L ~ ~ ( N Hwould ~)~+ not. When excess halide is present in the sulfite solution, therefore, it seems probable that solution is promoted by the formation of the AgBr2-’ complex rather than the sulfite complex. The observed linear dependence of rate upon bromide concentration in the presence of sulfite is in agreement with this suggestion. Equilibrium solubility, on the other hand, depends upon the sulfite complex, which, in turn, depends upon the concentration of silver ions in solution. We have noted previously that quaternary ions5 and thiourea6 increase the rate of solution of silver iodobromide and bromide. The reason for this is not clear, but may be associated with a depression of the negative surface charge of the original grains. Both the quaternary ion and the thiourea probably shift the surface charge t o the positive side115J6 and this may facilitate the direct passage of silver ions into the solution. The thallous ion could act in the same way. An alternative explanation for the effect of the quaternary ions and thiourea, based on a suggestion by Mitchell” that these (14) H. Chateau and B. Hervier, J . chim. phys., 64, 256 (1957). (15) S. Suzuki and Y . Oishi, Phot. Sci. and Eng., 1, 56 (1957). 61, 1216 (16) J. E. LuValle and J. M. Jackson, THISJOURNAL,
(1957). (17) J. W. Mitchell, J . Phot. Sei., 3, 73 (1955).
Vol. 82
agents displace gelatin from the surface of the silver halide, does not appear reasonable as an explanation for the effect of the thallous ion. WelliverL8has noted a similar effect of thallous ion in increasing the rate of fixing of photographic emulsions in thiosulfate solution. The relationship between the rate of solution and the specific surface of the silver bromide grains evidently depends upon the solvent used. We might expect that the more perfect the surface structure, the lower would be the rate of solution per unit area. On this basis, with ammonia as the solvent, the coarse grains would be more perfect than the fine grains. If this is so, however, we must then conclude that some factor is acting t o retard solution of the fine grains in the other solvents, or that a retarding factor is relatively more important for the fine grains than for the coarse. The surface charge, which stems from adsorbed halide ions, could be such a retarding factor. A relatively higher concentration of halide ions is to be expected at crystal edges, corners and areas of imperfection than at relatively perfect flat areas.. Small grains should have a higher average surface charge than large ones, and hence the retarding action of the charge effect toward the approach and adsorption of negatively charged solvent ions should be greater the smaller the average grain size. The relative and absolute effects should be larger the higher the negative charge of the solvent ion. These expectations are in line with the experimental results. The greatest “relative” retardation of solution of the small grains, compared with the ammonia results, is found for the doubly charged sulfite ion as solvent. A quaternary salt tends to eliminate this retardation (see Table 11). The fact that bromide and thiocyanate in the absence of sulfite dissolve the silver bromide grains at a rate which is substantially proportional to the specific surface may be largely fortuitous and represent a balancing of two opposing factors: the surface perfection factor which tends t o decrease the relative rate of solution of the larger grains and the charge effect which tends to decrease the relative rate of solution of the small grains. (18) L. G . Welliver, Phot. Eng., 6,203 (1955).
