42
T. H. JAMES SUMMARY
Singie crystals of copper, j in. in diameter and 5 in. long, are grown in a vacuum furnace using a modified form of the Bridgman method. A simple method of preparing a surface on these crystals parallel to a desired crystallographic plane is described. A very bright and smooth finish may be produced by using the electrolytic polishingmethod. This has the decided advantage of removing any strained or amorphous layers on the surface simultaneous with the polishing. The specimens so prepared may be used for studying any of the surface phenomena of single crystals of copper. We are indebted to Professor S. L. Quimby of Columbia University and to Mr. J. Zacharias for furnishing us with details of their furnace used for growing single crystals of nickel. REFERENCES (1) BRIDGMAN, P.: Proc. Am. Acad. Arts Sci. 80, 305 (1925).
(2) JACQUET, P.: Trans. Electrochem. Soc. 69, 629 (1936). (3) QUIMBY, S.: Phys. Rev. 39, 345 (1932). (4) TAMMANN AND SARTORIUS: Z. anorg. allgem. Chem. 176, 97 (1928).
MECHANISM O F PHOTOGRAPHIC DEVELOPMENT. I1 DEVELOPMENT BY HYDROQUINONE T. H. JAMES Kodak Research Laboratories,' Eastman Kodak Company, Rochester, New York Received M a y 10, 19.99
In a preceding paper (3) it was demonstrated that, in the normal course of development (absence of strong silver halide solvent action), the ions or molecules of the reducing agent must penetrate the double-layer electric barrier before reacting with the silver halide grain. The changes which occur in the potential barrier during the course of development and the variation from grain to grain in the magnitude of the original barrier significantly alter the kinetics of development by the ionic reducing agents; further, the barrier effect becomes more pronounced as the negative charge increases. The induction period observed in development by such agent.; as hydroquinone and p-hydroxyphenylglycine may be attributed largely to a decrease in the potential barrier as the reduction proceeds. 1
Communication KO.724 from the Kodak Research Laboratories.
DEVELOPMENT BY HYDROQUINONE
43
A more detailed investigation of the kinetics of development by hydroquinone has been made with the view of obtaining a clearer understanding of the mechanism of the chemical reaction itself. The apparatus and procedure were the same as those previously employed. The majority of the experiments were with sulfite-free solutions of pH 8.0 to 8.9. All solutions were prepared from components carefully freed from oxygen, and development was carried out under an atmosphere of pure nitrogen. The development rates were conveniently slow, and derelopment could be carried to apparent completion without formation of stain or significant fog. The emulsion employed was, as before, an unsensitized motion picture positive silver bromide material of thin and very uniform coating. Borax-boric acid buffers were used throughout. All temperatures, unless otherwise specified, were 20°C. & 0.05". Silver determinations were made by potentiometric titration at 70°C. with 0.0002 M potassium iodide. The determinations were of the amount of silver in a circular area of film in. in diameter, cut from the sensitometer strip by means of a steel punch. EXPERIMENTAL DATA
The pronouncedly S-shaped density-time curve obtained with a pure hydroquinone developer was illustrated in figure 2 of the preceding paper. Obviously, such curves cannot be represented by the simple formulas proposed in the literature (7). The upper third of the curves may be represented approximately b y the simple formula dD/dt = k(D,,,,
- D)
(1)
but this relation fails badly for the lower portion. Three methods of expressing the rate of development have been employed and found to give results consistent among themselves for the simple case of the pure hydroquinone developer. These are as follows: ( I ) determination of k in equation 1 from the data for the upper third of the curve; (6)determination of the slope of the density-time curve a t a fixed density (usually 3 D,,,, , since the curve a t this point is a reasonably straight line); (3) determination of the reciprocal of the time required to attain a given density. The rates so expressed are largely independent of the density chosen when we deal with change of pH or change of developer concentration. I n all of these cases, if a plot of density against log t is made, the curves obtained may be largely superimposed by the proper shift along the log t axis. The maximum density obtainable with a given exposure did not vary with the concentration of developing agent used or with the p H of the solution over the range employed. Likewise, the maximum amount of silver did not change. An example of the results is given in table 1.
44
T. H. JAMES
These data show further that, over a considerable range of exposure, the ratio D/Ag is constant when development is carried to completion. The ratio therefore does not depend upon the number of grains developed over this range of exposure. If, however, data for a fixed exposure and varying degrees of development are examined, it is found that theratio
variation of LOQE
1.15 0.85 0.85 0.85 0.55 0.40
I
TABLE 1 mazimum density m'th exposure
PH
DnUl.
8.89 8.89 8.29 8.00 8.29 8.48
2.50 2.02 2.01 1.98 1.55 1.30
1.68 1.68 1.71 1.68 1.64 1.63
* Silver expressed as milliliters of 0.001 M potassium iodide used in titration.
FIG.1. Silver-density relationship. Log E
=
0.85
D/Ag increases with decreasing degree of development. A plot of log Ag against log D yields a reasonably straight line of slope 1.45;Le., log Ag = 1.45 log D
+ Constant
(2)
(see figure 1). This relationship is interesting in connection with the results reported
DEVELOPMENT BY HYDROQUINONE
45
by Eggert and Kuster (1). They found that, over a rather wide range of conditions, the ratio Ag/D is a linear function of the mean grain diameter. From the relations : Ag/D = kid
+C
Ag = h d a where d is the mean grain diameter, it follows for constant grain number that log Ag
=
1.5 log D
+ 1.5 log (R+ CAg-”8)
(3)
This expression is in good agreement with our experimental results, since the constant C is very small. The agreement between the experimental results and equation 3 is good evidence that we have to deal largely with rates of reduction of the silver halide grains, and not with rates of initiation of the reduction process, since, in the latter case, the ratio D/Ag would be nearly constant for the fairly uniform grain sizes of the experimental emulsion. This conclusion is supported by observations on the bromide effect. The pure hydroquinone developer is very sensitive to small concentrations of bromide, and the effect is markedly to extend the toe of the development curve without appreciably altering the rate of development beyond the toe region (cf. table 4). The effect is more pronounced a t lower exposures than a t higher. These facts point to a strongly preferential action of bromide a t the start of development,-a phenomenon readily understandable on the basis of considerations given in the preceding paper of this series (3). Experiment shows, however, that if development is started in a bromide-free developer and continued to a suitable density (say 0.30 for log E = 0.85), addition of 0.1 g. potassium bromide per liter results in only a slight decrease in the development rate. The same amount of bromide, present a t the start of development, would increase fivefold the time required to obtain a density of 0.10. The dependence of the rate of development upon the solution pH is indicated in figure 2. A plot of the logarithm of the rate against the pH conforms approximately to a straight line of slope 1.0. Hence the rate varies as the first power of the hydroxyl-ion concentration.2 The rate varies approximately as the square root of the hydroquinone concentration. This is shown for two sets of data by the constancy of the ratios given in columns V and VI of table 2. The temperature dependence of the development rate is given in table 3. Temperature coefficient calculations from rates determined at the start of
* More accurately the data, corrected for change in concentration of hydroquinone which is due to ionization, indicate a dependence on the 1.08 power of the hydroxylion concentration and the 0.55 power of the hydroquinone concentration.
46
T. €JAMES I.
development and a t the midpoint are in fair agreement, although a dependence upon exposure appears in the former case, but is not definitely indicated in the latter.
Fro. 2. Val change in concentration of unionized hydroquinone.
applied for
TABLE 2 Variation of rate m'th hydroquinone concentration pH = 8.75 I HYDROQUINONE
gram pa 9w nl.
6.60 3.30 1.65 0.825 0.412
1
1
I
I
I
VI
I1 (10 X Hq)l't
8.12 5.74 4.06 2.87 2.03
RATIO
0.397 0.280
30.0 33.7
0.082
0.192 0.130
0.058
0.090
33.0 35.0 35.0
0.270 0.170 0.123
nm
20.5 20.5 21.2 22.2 22.5
The pure hydroquinone developers employed in this investigation are highly sensitive to the bromide-ion concentration. The major effect of small concentrations of bromide is exerted a t the start of development. Larger quantities of ,bromide, however, affect the entire course of development. This is indicated by the data in table 4. The times employed
47
DEVELOPMENT BY HYDROQUINONE
in determining the rate of initiation of development are for different densities (Le., 0.20 and 0.10) for the two exposures recorded, the densities having been chosen to give equal rates of reaction in the absence of soluble bromide. The depressing effect of small amounts of bromide (less than 0.2 g. per liter) is thus confined almost exclusively to the ear' stages of development, and is most pronounced in the region of low exposures (13). If development is started in the absence of added bromide and then con7
TABLE 3 Dependence of rate upon temperature Hydroquinone concentration = 0.05 M ; pH (at 2OOC.) = 8.89 LOO
E
1.75 1.15 0.85 0.25
1 R
D
AT
8.0%.
1/2
mar.
ATa20.0'c' DR = 112 mar.
~
0.090
1
1
TEYPEBATURE COEPPIC,ENT
1
TEYPEBATURE COEFFICIENTFROM
111; D = 0.10
2.95 3.1 3.3 3.9
0.315 2.90 0.245 2.9 0.056 0.208 3.1 Accurate values could not be determined 0.070
~
,
~
TABLE 4 Effect of added bromide Hydroquinone = 0.05 M; pH = 8.71 MQ
KBr ADDED
1
D = 0.20
-
!
E = 1.46
M O B 0.S
I
1
Rate D .S 1.00
D = 0.10
0.107 0.052 0.0272 0.019 0.0098
0.192 0.190 0.190 0.12 0.030
9.2 24.1 46 65 143
0.123
0.192
8.0
Rata Ill
Rate 1/L
pmma p B lilar
0
9.3 19.3 36.7 53 102
1.00 Washed film. ,. . .
.I
8.1
0.109 0.0415 0.0217 0.0154 0,0070 0.125
tinued in a developer to which 0.10 g. of potassium bromide has been added, no significant change occurs in the rate, provided the initial development is carried to about the end of the induction period. Larger amounts of bromide, however, exert a significant influence over the entire course of development, as indicated by the much smaller rate observed at D = 1.00 when the solution contained 1.00 g. of potassium bromide per liter. It has already been demonstrated that the addition of quinone to the hydroquinone solution increases the development rate. Accordingly, it
48
T. H. JAMES
was desirable to check the kinetic results already obtained by similar measurements made in the presence of added quinone and in the presence of sodium sulfite (which rapidly removes the quinone formed during development). The same dependence of the reaction rate upon the square root of the hydroquinone concentration and the first power of the hydroxyl-ion concentration was obtained in these experiments as in those involving only hydroquinone. This is illustrated by the data in tables 5 and 6. It will be observed that, both in the presence of added quinone and in the presence of sulfite, a fourfold change in the concentration of hydroquinone results in only a twofold change in the development rate. From TABLE 5 The reaction rate i n the presence of added quinone Added quinone = 0.001 M HYDROQUINONE CONCENTRATION
'
I
LOO
HYDROQDINONE
0.05 0.10
I
8.67 8.89 8.89
L O Q E-0.85
j
~
j
0.0196 0.0336
E
RATED =
:
-
0.85
0'1292
0.526
0.40
1
LOO E
--
1.45
RATED 0.W
LOO
0.0296 0.049 0.034 0.068
E = 1.45
0.471 0.69 0.53 0.83
table 5 we see that a change of 0.39 in pH results in a change in log R of 0.35 and 0.42, respectively, while in table 6 a change of 0.22 in pH is accompanied by a change of 0.22 in log R. Hence the rate of development once more is proportional to the hydroxyl-ion concentration and the square root of the hydroquinone concentration, Figure 2 of the preceding paper (3) shows that the induction period of hydroquinone development decreases with increasing concentration of added quinone. Further experiment showed that the maximum density obtained for a given exposure is sharply decreased by the addition of quinone. Figure 3 is based upon data obtained a t pH 8.29. A direct plot of silver against development time has been made here. Curve I is
DEVELOPMENT n Y HYDROQUINONP
49
for pure hydroquinone solution; curve I1 is for hydroquinone solution containing 0.001 M quinone. The maximum amount of silver has decreased from 1.11 to 0.82 (x 10-3 millimoles), and the maximum density from 2.01 to 1.32. This decrease cannot be due to a shift in the equilibrium 2Ag+
+ CaH,Oa--
e 2Ag + CaI402
since thermodynamic considerations show that the reduction would go practically to completion in both cases. The behavior of the photometric equivalent gives the clue\ to the true cause.
FIQ.3. Effect of quinone (0.001 M ) on development by hydroquinone (0.05 M). I, pure hydroquinone; 11, quinone added.
In-table 7 the ratio of silver obtained by development in the hydroquinone-quinone solution to that obtained without the addition of quinone (development carried to the same density) is given for several density values. If, instead of comparing equal densities for the same exposure, we compare equal m a x i m u m densities, we find that there is no significant difference in the photometric equivalent. For example, the maximum density obtained in the presence of quinone is 1.32 for log E = 0.85. This density is matched by the hydroquinone developer for log E = 0.41. The experimental ratio of the amounts of silver obtained was 1.005. Hence, the effect of quinone upon the maximum density is equivalent to a reduclion in the amount of exposure. The data are accounted for if we assume that the presence of quinone effects no change in the diameter of the developed grains, but reduces the number of grains developed. The data
are not accounted for on the assumption that the number of developed grains remains unchanged, but that each attains a smaller maximum size in the presence of quinone. Further evidence in this direction appears from the fact that, if development is started in the absence of added quinone, carried to a density of 0.25 (log E = 0.85), then continued in the quinone-containing solution after thorough washing, a maximum density of 1.95 is obtained in comparison with 2.0 for the pure hydroquinone solution. T h e ratio Ag(added quinone)/Ag(no added quinone) drops to 1.03. I n an attempt to determine to what extent continued development is influenced by the quinone formed in the reaction', a series of tests was made in which the sensitometer strips were partially developed, then thoroughly washed in running water, and,finally subjected to continued development in fresh hydroquinone solution. I n part of the runs, hydroquinone itself
'
DENSITY
UTI0 A g ( a u m o m ) / A g ( N o ADDED QUINONli)
1.32. 1.19 1.00 0.58
1.30 1.23 1.28 1.21
* Maximum development in hydroquinone-quinone solution.
.
was used for the initial development; in the rest, a ferrous oxalate developer was employed. An examination of the data obtained showed that, when the initial development was slight, the intermediate washing process had no effect upon the course of continued development, and the results obtained with initial development by ferrous oxalate checked those obtained with initial development by hydroquinone. As the amount of initial development increased, more and more divergence appeared between the course of the secondary development and that of the uninterrupted development, and the decrease in the rate of secondary development was significantly greater when ferrous oxalate was used for the primary development. If, however, the initial development is carried out by hydroquinone in the presence of a small amount of sodium sulfite, the rate of continued development in pure hydroquinone solution is approximately equal to that found when the initial development is by ferrous oxalate. Under these circumstances, the initial rate of the secondary development does not diverge greatly from the rate of development by a hydroquinone
51
DEVELOPMENT BY HYDROQUINONE
solution containing sulfite or resorcinol when the latter rate is measured a t the same density. Some comparative results are given in table 8. The rates given for secondary development apply to experiments in which the initial development was carried to densities 0.02 to 0.04 unit less than the densities a t which the rate measurements were made. These results clearly show that quinone catalysis is not responsible for the induction period in hydroquinone development. However, an accelerating effect of quinone or some degradation product thereof does persist throughout the entire course of development If development started in a pure hydroquinone solution is interrupted and the film washed in running water for 30 min. or bathed in sodium bisulfite solution for 10 min., continued development in a fresh hydroquinone solution proceeds a t a rate comparable with that of uninterrupted development. This TABLE 8 Effect of removal of quinone during development upon the rate of continued development TYPE OF DEVELOPYEXT
Primary Secondary Secondary Secondary Primary Primary Primary Primary Primary
O Q E = ,oQE 80LQllON
Pure hydroquinone Initial development by pure hydroquinone Initial development by hydroquinone-sulfite Initial development by ferrous oxalate Hydroquinone plus 0.05 M resorcinol Hydroquinone plus 0.10 M resorcinol Hydroquinone plus 0.005 M sulfite Hydroquinone plus 0.05 M sulfite Hydroquinone, 0.005 M sulfite, 0.10 M resorcinol
-
1.45
0.85
ID-0.4
lD-0.30
0.11 0.10 0.038 0.054 0.034 0.031 0.028 0.028 0.028
0.078 0.07 0.023 0.025 0.023 0.022 0.020 0.020 0.023
indicates that the immediate presence of quinone is not essential for the “catalytic” effect. A quinone catalysis of the type which exists in the autoxidation of durohydroquinone (4)is therefore eliminated. An effect of quinone upon the electric barrier may still be involved, since this might persist after the quinone had been removed. However, a comparison of the relative accelerating action of benzoquinone, toluquinone, and mxyloquinone upon development by the respective hydroquinones indicates that a decomposition product of the quinone, rather than the quinone itself, is largely responsible for the effect. I n table 9, T I represents the time required to obtain a density of 0.50 (lob E = 1.15) with the pure hydroquinone developer; Tz represents the time required in a developer containing 0.001 mole of the corresponding quinone. The pH and hydroquinone concentration were adjusted to give approximately the same development rate in the absence of added quinone.
52
T. H. JAMES TABLE 9 Accelerating effect of quinone homologues on development by the corresponding hydroquinone Quinone = 0.001 M
I
I
1
...I I
Hydroquinone. . . . . , . . . . . . . Toluhydroquinone.. . m-Xylohydroquinone.. . . . . .
TIME BEOUIBED POB
1 !::I :85:85 1 _
8":
M/20 M/20 M/80
-
D
ca. ca. 8.0
_
_
_
0.60
_
_ ~ 1.80 1.30 1.08
25.0
TABLE 10 Comparative measurements with substituted hydroquinones Sulfite-free developers ihosphate buffer
--
DBVBLOPMENl BATE
D T 0.60 Mfrn
COYPOUND
ONCBNTBATIOI
mrnnvm RATE
pH
pH-7.9
-___
Hydroquinone. . . . . . . . . . . . . . . . Toluhydroquinone . . . . . . . . . . . . m-Xylohydroquinone . . . . . . . . . Chlorohydroquinone . . , . . . . . . . 3-Chlorotoluhydroquinone . . . . Dichlorohydroquinone . . . . , . . . Bromohydroquinone . . . . . . . . . . Hydroxyhydroquinone . . . . , . . .
0.014 0.021 0.044 0.057 0.100 0.066 0.065 0.053
AUTOXIDATION RAT', K/'J
1.00 1.5 3.1 4.0 7.1 4.7 4.6 3.8
* Values enclosed in parentheses are approximate
-
EIATIVE RATE
7.44
0.00123 0.00479 0.0234 0.0092
1.00 3.9 19.0 7.5
(W* (0.0142) (0.0132) (0.310)
(11.5) (10.7) (2%)
only.
TABLE 11 Comparative measurements w i t h substituted hydroquinone8 Sulfite-containing- developers: borate buffer
I
COMPOUND
BBLA-
D%%
mvm
LUTOXIDATION BATE
BATE
pH
1 pH
Hydroquinone. . . . . . . . . . , . . . . . Toluhydroquinone , . . . . . . . . . , . m-Xylohydroquinone . . . . . . . . . Chlorohydroquinone . . . . . . . . . . 3-Chlorotoluhydroquinone . . . . Dichlorohydroquinone . . . . . . . . Bromohydroquinone . . . . . . . . . . Sodium hydroquinonemonosulfonate. . . . . . . . . . . . . . . . . . .
--
-
DlDVELOPMENT
8.9
0.020 0.039 0.0846 0.089 0.108 0.146 0.109
7.81
1.00 1.95 4.23 4.45 5.40 7.3 5.45
(0.077) (0.052)
0.10
O.OOO53
0.0046
0.0180 0.076
or "",",rPpH-8.8 A@
BEDUCTION
BnLA-
1.00 3.8 16.5 (7.1) (30.0) (16.7) (11.3)
0.038 0.027 0.10 0.47
1.00 0.71 2.6 12.4
0.115 0.0038
TIVE RATE
0.10
DEVELOPMENT BY HYDROQUINONE
53
Clearly, the accelerating effect of the quinone decreases sharply as its stability increases. COMPARATIVE RATES OF DEVELOPMENT, AUTOXIDATION, AND SILVER-ION
REDUCTION
Comparative measurements were made with a number of substituted hydroquinones of the development rate, the autoxidation rate, and the rate of the silver-catalyzed reduction of silver ion. The autoxidation rates were determined according to the procedure given by James and Weissberger (4) and are expressed as “first-order” constants. The rates of the silver-catalyzed reduction of silver ion were determined as described in a previous paper (2), except that sodium sulfite was employed to decrease the silver-ion concentration and thus allow measurements to be made a t a higher pH. Some check experiments showed that the presence of the sulfite effected no significant change in the mechanism of the reduction and the silver sulfite complex wae not attacked to a detectable extent by the hydroquinone. The comparative rates are given in tables 10 and 11. DISCUSSION
The work of Sheppard and Mees (13) and of Luther and Leubner (6) indicated that the divalent hydroquinonate ion is responsible for the reduction of silver bromide by the hydroquinone developer. Our data lend themselves to the same interpretation. In the pH range of 8.0 to 8.9 we may write for constant amount of added hydroquinone:
since the concentration of unionized hydroquinone will be constant within 10 per cent. Further, a t constant pH but varying hydroquinone concentration: [CsH,O2--] = K’[CeHsOe]
(5)
From these equations, it follows that the square root of the-divalent ion varies as the first power of the hydroxyl-ion Concentration and as the square root of the hydroquinone concentration. The rate of development shows the same dependence upon hydroxyl-ion concentration and hydroquinone concentration, and hence varies as the square root of the concentration of divalent ion. This we interpret to mean that the divalent ion is the entity which reduces the silver halide in development. The occurrence of a 3 power in the rate equation may be taken as direct evidence that adsorption plays an important rale in the development process. The kinetics are quite analogous to those frequently observed
54
T. H. JAMES
in heterogeneous reactions in which adsorption follows approximately the Freundlich isotherm and the rate obeys the equation Rate = IcC1ln
(6)
Whether the adsorption is to the silver n u c l d the silver halide itself, or the silver-silver halide interface is not indicated. Rabinowitsch and his coworkers (8, 10) believe that the adsorption is to silver. However, their experiments were so complicated by the occurrence of a large amount of oxygen oxidation that their claims to establishing adsorption must be regarded with suspicion. This is particularly the case since recent measurements made in these Laboratories by Sheppard, Ballard, and Perry failed to show any adsorption within experimental error. Further, it is not apparent why, if adsorption to the silver is the essential thing, the adsorption effects may be detected so readily in the development kinetics but not appear in the silver-catalyzed reduction of free silver ion. The silver nucleus is present in both cases. On the other hand, Wulff and Seidl (16) showed that resorcinol was adsorbed by silver bromide from alkaline solution, and argued by analogy that hydroquinone would be so adsorbed, in agreement with Sheppard’s theory (14). Their experimental evidence for adsorption of hydroquinone, i.e., that resorcinol decreases the rate of development by hydroquinone, is untenable, but the analogy appears reasonable. Further, Rabinowitsch’s failure to detect adsorption of hydroquinone from acid solution furnishes no rebuttal, since Rabinowitsch’s measurements would not have revealed adsorption of the divalent ion or even of the univalent ion. We may conclude with reasonable certainty that, in the process of development by hydroquinone, the reduction occurs heterogeneously at the silversilver halide interface. The Ostwald-Abegg supersaturation mechanism can offer no explanation either of the general kinetics of the development process or of the potential barrier effect which appears so strongly under the conditions we have employed. An alternative mechanism, that of solution of the silver halide followed by silver-catalyzed reduction of.the silver ions, is likewise ruled out by the fact that the development kinetics conform neither to those of solution of silver halide nor of catalyzed reduction of silver ions present initially in the solution. The hydroquinone, as the divalent ion, reacts only after adsorption, presumably a t the silver-silver halide interface. This adsorption undeniably affects the kinetics of the reduction process. However, the def6rmation experienced by the silver ions “adsorbed” a t the active boundary lines of the silver-silver halide interfaces in the developing grains is probably a more fundamental reason for the enhanced reactivity of the nucleated grains which makes possible differential reduction (12). The pirture of the development process which we have presented in
DEVELOPMENT BY HYDROQUINONE
56
this a i d the preceding paper appears to correlate satisfactorily the existing data on sulfite-free hydroquinone developers and solutions containing low concentrations of sulfite. When strong silver halide solvent action is present, other processes may, and probably do, take a significant part in the development process. Some further support is to be found in an examination of the relative rates of development, autoxidation, and silver ion reduction given in tables 10 and 11. There is no correlation between the rates of development and those of silver-ion reduction. Some correlation exists, however, between the development rates and the autoxidation rates. This is explainable on the basis that the active species in each case is the divalent ion, in contrast to the silver-ion reduction where the reaction of the univalent ion predominates. The correlation is simplest in the series: hydroquinone, toluhydroquinone, m-xylohydroquinone. Here, when sulfite is present in the development process, the rate of development varies almost exactly as the square root of the autoxidation rate,-a relation which is interesting in connection with the dependence of the development rate upon the square root of the divalent-ion concentration. The divergence from this simple relationship shown by the rates measured in the absence of sulfite is due to the quinone effect already discussed. It thus appears that the autoxidation rates furnish a clue to the rates to be expected for development by substituted hydroquinones, provided the substituent group or groups are not highly ionized. In the latter event, the electric barrier effect may become dominant. The extent to which the electric barrier effect may influence the relative development rates exhibited by agents belonging chemically to the same class is indicated by a comparison of hydroquinone with its monosulfonate (charge group -3) and its disulfonate (charge group -4). At pH 8.9, the pure sulfite-free solution of the monosulfonate develops a t only about 0.08 times the hydroquinone rate. When M/20,000 phenosafranine is present, however, the monosulfonate rate has been increased to about 0.2 that of hydroquinone. The phenosafranine effects a threefold increase in the hydroquinone rate and a sevenfold increase in the monosulfonate rate. Only the faintest trace of an image was obtained in 4 hr. with potassium hydroquinonedisulfonate a t pH = 9.5. When M/10,000 phenosafranine was added, a good image was obtained in 20 min. This shows that the apparent weakness of the disulfonate as a developer is due to the inability of the highly charged ions to penetrate the electric barrier, rather than to a lack of reactivity. The potential barrier effect enters significantly into the action of elonhydroquinone developers. The density obtained in a given time of development by a hydroquinone-elon mixture is greater than the sum of the densities obtained by development with the two agents separately.
56
T. H. JAMES
This phenomenon has been attributed by various writers (5, 15) to the action of a “metoquinonc” possessed of special developing properties or to an elon-hydroquinone adsorption complex. However, much the same effect can be obtained if the development is started in an elon solution and, after thorough intermediate washing, is continued in a hydroquinone solution. It appears, therefore, that the preliminary development in elon merely accomplishes a partial breakdown of the electric barrier and thus allows a greater rate of penetration by the more sensitive hydroquinone ions. Such action should be kept in mind when considering the r81e which the sulfonates play in development. The experimental results given in the preceding pages do not support the contention of Reinders (11) that “it may certainly be concluded that there is a close relation between the velocity of development and AE.” The dependence of the rate upon the solution pH and the concentration of the developing agent appears quite clear from the kinetic considerations, and requires no association with the thermodynamic quantity for elucidation. The marked dependence of the reaction rate upon the size of the developed silver speck (9) does not follow from Reinders’ argument, while the accelerating effect which quinone has upon the development process is exactly contrary to that argument. It may be noted, likewise, that the decrease in “maximum density” observed when quinone was added to the pure hydroquinone developer is not a true equilibrium phenomenon, and should not be associated with the change in AE. SUMMARY
1. The rate of development by hydroquinone in the absence of sodium sulfite and oxygen varies as the square root of the hydroquinone concentration and the first power of the hydroxyl-ion concentration in the pH range 8.0 to 8.9. The rates measured approximate the average rates of development of the individual silver bromide grain. 2. The addition of quinone to the developer results in an increase in the development rate, which is due to an effect, presumably of a decomposition product of the quinone, upon the double-layer electric barrier. The quinone likewise attacks the development centers, resulting in a decrease in the number of developed grains for a given exposure. 3. The divalent ion is the active developing agent. Reduction occurs heterogeneously a t the silver-silver halide interface. 4. Comparative rates are given for the development reaction, autoxidation, and the silver-catalyzed reaction of silver ion with a series of substituted hydroquinones.
The writer is indebted to Dr. S. E. Sheppard of these Laboratories for many valuable suggestions and criticisms.
SILICIC ACID GEL6
57
REFERENCES (1) EQQERT, J., AND K ~ ~ S T EA.: R , Kinotech. 18, 381 (1936). T . H.: J. Am. Chem. SOC.61, 048 (1939). (2) JAMES, (3) JAMES, T . H.: J. Phys. Chem. 43, 701 (1939). T . H., AND WEISBBEROER, A.: J. Am. Chem. SOC.80, 98 (1938). (4) JAMES, ( 5 ) LUMIBRE,A,, LUMIBRE,L., AND SEYEWETZ, A.: Phot. Korr. 62, 183 (1926). (6) LUTHER,R., A N D LEUBNER, A.: Brit. J. Phot. 69, 632 e t seq. (1912). (7) c f . NIETZ,A. H.: Theory of Development. D. Van Nostrand Company, New York (1922). (8) RABINOWITSCH, A. J.: Trans. Faraday SOC.34,920 (1938). A. J . : Trans. Faraday SOC.34, 921 (1938). (9) RABINOWITSCH, (10) RABINOWITSCH, A. J., AND PEISBACHOWITSCH, 9.: Z. wiss. Phot. 33, 94 (1934). (11) REINDERS, W.: Trans. Faraday SOC. 34, 936 (1938). G. M.: Z. Elektrochem. 36,584 (1929). (12) cf. SCHWAB, (13) c f . SHEPPARD, S. E., AND MEEB,C. E. K.: Investigations on the Theory of the
Photographic Process. Longmans, Green and Company, London (1907). AND MEYER,G.: J. Am. Chem. SOC. 4, 689 (1920). (15) SHIBERSTOV, V.: Photo-Kino Chem. Ind., 2, 141 (1933). (16) WULFF,P., AND SEIDL,K.: Z. wiss. Phot. 28, 239 (1930). (14) SHEPPARD, S. E.,
STUDIES ON SILICIC ACID GELS. I X THEEFFECT OF A CHANGE OF PHUPON THE TIME OF SETOF SOME ACID GELS CHARLES B. HURD AND HARRIS W. PATON' Department of Chemistry, Union College, Schenectady, New York Received January 80, 1059 INTRODUCTION
A recent study of the time of set of silicic acid gels, prepared from solutions of sodium silicate and acetic acid in this laboratory ( 5 ) , has shown that, for mixtures of acid reaction, the time of set is proportional to the hydrogen-ion concentration. A report by Hurd, Frederick, and Haynes (2) showed a similar result when hydrochloric, nitric, or sulfuric acid was used. A discussion of the theory has also been given (1). It was observed, incidentally, that addition of extra acid to mixtures already prepared delayed their setting. We have thought that it would be very important to discover whether there might be some particular time, during the setting of the already prepared mixture, when the introduction of the extra acid would have a particularly marked effect. If such Present address: Corning Glass Works, Corning, New York.
