T H E OXIDATION-REDUCTION POTENTIALS OF UNSTABLE ORGANIC SYSTEMS1 A, E. CAMEROX
Kodak Research Laboratories, Rochester, New York Received April 9, 1988
The determination of the oxidation-reduction potential of a system requires that the two forms be present and in thermodynamic equilibrium under the conditions of the experiment. The rapidity with which the oxidized forms of most photographic developing agents are converted into other compounds, particularly in alkaline solution, makes direct measurements upon a mixture of reduced and oxidized forms uncertain. Fieser (5, 7 , 8 ) introduced a method of making oxidation-reduction measurements upon such systems, based upon the experimental fact that the oxidized forms of many irreversible systems have a definite but short life in solution. When enough oxidizing solution is added to the reduced form of such a system to produce a half-oxidized mixture, and the readings of the observed potential of a noble metal electrode are plotted as a function of time, it is frequently possible to extrapolate to zero time and thus determine the potential of the mixture before alteration in the concentration of the oxidized form takes place. This paper is the result of the application of Fieser’s method to measurements upon certain organic systems of interest t o photographic theory. EXPERIMEKTAL
The p-methylaminophenol sulfate used was the photographic product Elon, manufactured by the Eastman Kodak Co. This was given the purification treatment outlined by Schering (13) t o eliminate any p-aminophenol present. A hot saturated solution of p-methylaminophenol sulfate in dilute acetic acid was treated with benzaldehyde, allowed to cool, the crystals filtered off, and a second crop of crystals thrown out by addition of acetone. Two recrystallizations were carried out from dilute sulfuric acid solution, with rejection of the first small crop of crystals formed in each case. The monosulfonic acid of p-methylaminophenol was prepared by T. H. James of these laboratories by heating it with sulfuric acid for 8 hr. on a water bath and pouring it into water. No further purifiration Contribution No. titi0 from the Kodak Research Laboratories. 1217
1218
A . E. CAMEROX
was carried out, other than a thorough washing of the precipitate. The hydroquinone used was the photographic product manufactured by the Eastman Kodak Co., and the hydroxyhydroquinone was twice recrystallized before use. All buffer salts used were of reagent quality from reliable manufacturers, as was the potassium ferricyanide used as oxidizing agent. For measurements in acid solutions it was found necessary to prepare potassiuni molybdicyanide. In this preparation the directions of Fieser (6) were followed. The oxidizing solutions werc stored in dark glass bottles wrapped in opaque cloth. The nitrogen used to maintain an inert atmosphere was taken from a commercial cylinder of compressed gas. Gas from the cylinder was passed through two towers of solid ammonium carbonate, and then through a tower of soda lime, leaving a definite partial pressure of ammonia in the gas. The mixture was passed through a quartz tube packed with platinized asbestos free from sulfur and heated to 450-500°C. Ammonia remaining after the reaction of ammonia and oxygen was removed with wash towers of 30 per cent sulfuric acid colored with bromophenol blue to indicate exhaustion, and acid spray and carbon dioxide were removed with a wash of 60 per cent potassium hydroxide. A bottle of a 0.1 per cent solution of reduced indigodisulfonate in the line indicated complete absence of oxygen, even a t high rate&of flow. I t is important to insure saturation of the gas stream with ammonia, however. Exhaustion of any reagent was readily noted, and the line could be operated for long periods of time without attention. Solutions of the compounds invebtigated were made up in 0.02 M concentration and stored in clear glas,i flasks in an atmosphere of hydrogen. The flasks were arranged so that the solutions could be delivered to waterjacketed, IO-mi. burets without destroying the protective atmosphere. The p-methylaminophenol and the monosulfonic acid both showed some sign of decomposition, even in acid colution. They gradually acquired a brownish-red coloration. This decomposition did not procped to the point where a change in titer could be noted. The reference half-cell and bridge employed in the pH detcrnlinations and in the oxidation-reduction potential measurements were 3.5 N in potassium chloride. This solution was prepared volumetrically a t 20°C. This half-cell was assigned the \ d u e of 0.2502 volt against the normal hydrogen electrode at 20"C., and this value was used throughout the calculations. APPARATUS
The essential parts of the apparatus are outlined in figure 1. A doublewalled beaker, the top of which was flattened and ground plane., was
OXIDATION-REDUCTION
POTEII'TIALS
1219
fitted with a flanged glass cap, also ground plane, which carried a water seal for the stirrer, a glass electrode, and a neck in which could be inserted a stopper carrying the platinized electrodes, or blank platinum and gold electrodes. Lead-in tubes for the salt bridge and for the gas supply to a porous alundum bubble head were sealed through the cap. Three small necks on the cap permitted insertion of buret tips with pieces of small rubber tubing for gasketing. A thermometer mounted in this same cap permitted the determination of the temperature of the reaction mixture. Water was pumped through the jackets of the reaction vessel, the half-cell, and the burets from a 35-gallon thermostat operating at 20°C. =k 0.01'.
rr*TER
FIG.1. The apparatus
The addition tube, shown in figure 1, made it possible to add the oxidizing solution rapidly and flush the vessel with buffer solution without admitting atmospheric oxygen. With the pinch clamp on the rubber tubing, A, nitrogen from the reaction vessel swept out compartment B, containing the oxidizing solution, and bubbled through the buffer solution in compartment C, before escaping to the air through the bubbler. When the clamp was removed from the tubing, nitrogen from the reaction vessel could replace the solutions when the lower and upper stopcocks were opened in succession. The stopcocks were both of large bore to permit 10 ml. of solution to be added in about 2 sec. The experimental procedure followed consisted in adding 100 ml. of buffer solution, 0.2 M in buffering ion and with a sodium-ion concentra-
1220
A . E . CALVEROB
tion fixed by addition of sodium sulfate, as necessary, to the reaction vessel, and determining the p H with two hydrogen electrodes. The glass electrode was then read, and the hydrogen displaced with pure nitrogen with the substitution of blank platinum and gold electrodes for the platinized ones. The addition tube was set in place with the measured amount of oxidizing solution and buffer solution for flushing, and half an hour allowed for the sweeping out of hydrogen and air. Ten milliliters of reducing agent was then added from one of the jacketed burets. The lower stopcock of the addition tube was opened, and at the same time an electric timer was started which g a m a single stroke on a bell at 15-sec. intervals. TABLE 1 p-Meth ylaminophenol, 50 per crnt oxidized, at 20°C PH
E
__________
___.____
PH
lolls
0 96 1 37 2 10 2 59 3 38 3 86 4 32 4 66 2 81 1 90 5 28 5 78 5 80 6 -10 7 00 8
In
8 13 8 11 8 25 8 11
0 0 0 0 0 0 0 0 0 0 0
0 0 0 0 0 0 0 0 0
6869 6574 6182 5919 5402 5122 4762 $463 1309 4301 3960 3574 3582 3220 2820 2202 2135 2132 2040 1990
-
E'
--
LdtS
8 9 9 9 9 9 10 10 10 10 10 10 11
57 14 32 52 81 93 33 36 47 75 81 85 02 11 30 11 35 11 69 11 97 12 1s 13 07 13 60
0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 -0
1884 1552 1465 1352 1240 1130 0934 0930 0892 0820 0745 0732 0702 0602 0560 0452 0370 0319 0052 0118
The upper stopcock was then opened and the lower chamber flushed into the reaction vessel. Readings were taken at 15-sec. intervals for 3 min. and at I-min. intervals for 'i min. The glass electrode was then read again, and the pH of the bolution as originally determined with the hydrogen electrode was corrected tor any change in pH which had occurred. Readings were taken with a qtudent-type potentiometer. The potential of the glass electrode was read with the snme potentiometer, using a ballistic galvanometer and a I microfarad condenser as null point indicator. .\ddition of the solutions raised the temperature of the reaction mixture aligiitly, brit thi. effect rould be disregarded, since in most capes the
1221
OXIDATIOS-REDKCTION POTEKTIALS
extrapolation t o zero time was made from a graph which was a straight line up t o 10 min. The p-methylaminophenol and sulfonic acid solutions were standardized against the oxidizing solutions by diwontiniiow titrations, a run being TABLE 2 p-,~~l.lethylaininopiie7~ols~il~on~c acid, 60 p i ' cent oxzdtzed, at 20°C E'
PH
PI1
- ___-___-
Lolls
1 03 1 58 2 05 2 54 2 73 3 90 4 23 4 48 4 87 4 92 5 26
I
I
I
0 0 0 0 0 0 0 0 0 0 0
7323 7000 6729 6414 6221 6389 5102 4902 4585 4542 4229
E' VOll8
I
'I
5 59 6 62 7 o5 8 41 9 21 9 77 10 31 10 70 11 39 11 45 11 81
0 0 0 0 0 0 0 0 0 0 0
l
4062 3422 3194 2359 1942 1622 1317 1175 0902 0890 0774
TABLE 3 Bentohydroqiiinone, 60 per cent oxidized, at 20°C PH
E
E'
gH .__
bolls
6 6 7 7 8
28 0 3342' 0 3371 33 0 2982 03 17 0 2900 0 2370 01 h 97 0 1844 9 64 0 14F9 10 36 0 1209 10 41 0 1189 10 43 0 1172 10 67 0 1104 10 86 0 1025 11 06 0 1022 11 32 0 0916 __* B y addition of quinhgdrone solution t B y addition of quinone solution
,
11 11 11 11 11
32 34 54 60 61 85 00 30 36 43 55
11 12 12 12 12 12 12 78 12 81 13 01
' 1
I
colts
0 0 0 0 0 0 0 0 0 0 0 0 0 0
0957 0945 0874 0888 0872* 0820 0788 0701 0664*
0661t 0560* 0370* 0453 0303*
_____
__
made for each point on the titration curve. T h e hydroquinone solution was standardized by direct titration a t a pH at which decomposition of the quinone did not interfere, and the hydroxyhydroquinone solution was standardized by rapid titration at pH 0 5 , and by a discontinuous titration.
1222
A. E. CAMERON DATA
The data for p-methylaminophenol and p-methylaminophenolsulfonic acid are given in tables 1 and 2. Table 3 contains the data obtained from an attempt t o carry the measurements of the oxidation-reduction potential of the hydroquinone system up into the alkaline range where the quinone is unstable. The data, othcr than those specially marked, were obtained by the addition of sufficient ferricyanide solution t o produce a half-oxidized mixture, or by selecting the midpoint of the titration with a ferricyanide solution in the l o w r pH range. Attempts were made t o investigate the oxidation-reduction potential of hydroxyhydroquinonr over a considerable range of pH. The data for these measurements appear in table 4. TABLE 4 Rydioxyhydroquinone, 50 p e ~cenl oxidized, at 20°C PH l_I - - .
I I
PH
- 1
kOlt8
0 50 1 26 2.30 2.7G 3.40 3.58 3.76 3.87 4.06
I
E'
0 0 0 0 0 0 0 0 0
5707 5310 4690 4394 3896 3710 3690 3487 3797
11
4 4 4 5 5
20 24 e2 32 86 e 04 e 75 7 30
I
I
~
I
-
I
E'
I
volts
1 I
0 0 0 0 0 0 0 0
3310 3490 3217 2542 2102 1840 1255 0802
DISCUSSIOS
The equilibrium between the oxidized and reduced forms of p-methylaminophcnol may be formally represented by the following expression: 2e
Ox.
j-2H+ F' Red.
T h e participation of hydrogen ions in the equilibrium and the possibilities of ionization of the secondary amino groups in both reduced and oxidized forms renders the system rather complicated. An cxprcssion which irns dcrivcd and found to fit the experimental data with pleasing rl\;sctitiide was tlic following'
E = Eo - 0.029 log10 S,/So - 0 029 p H
OXIDATION-REDUCTION POTENTIALS
1223
I n this expression Eois a constant which is characteristic of the system, S R and So are the total concentrations of reduced and oxidized forms, K , is the ionization constant of water a t ZO'C., K , is the acid ionization of the phenolic group in the reduced form, and K O and K R are the apparent basic. ionization constants of the secondary amino groups in oxidized and reduced forms. The basic ionization constants are defined for the oxidized form by the expression:
and by a similar expression for the reduced form. I n [Ox] and [Red] are included the concentrations of both hydrated and unhydrated un-ionized forms.
Figure 2 shows the graphs of the data for both systpms. I n these graphs the lines are drawn with theoretical slopes, and the constants given in table 5 were used in the calculations. It will be noted that the data fit the calculated lines with excellent regularity. Data from other sources are included for purposes of comparison. Fieser found that, worlriiig in the presence of air, he was not able to carry the measurements above pN 8 before the rate of decomposition of the oxidized form became so rapid that extrapolation to zero time of the
1224
A . E. CAMERON
potential-time plot was impossible. In the present work it was found possible to carry the measurements into quite alkaline solutions before the deconipositioii became too rapid. I n all measurements of both compounds, the graph of the potential against time was found to be a straight line for a considerable period. In some cases it remained a straight line up to 15 miri. This indicated, since the drift was toward lower potentials, that the oxidized form was decreasing in concentration by a reaction whose rate was unimolecular. The mechanism of this decomposition has been discussed by Fieser, but no direct data appear t o be available. I n the range of pH from the pKo of tht. two compounds up to pH 8.0, the decomposition leads to the formation of a purple compound. I n lower ranges of pH the iminoquinone was yellc v. The purple tolor appeared to develop as a function of time, indicating that it was due t o a combination of the decomposition product of :he iminoquinone with either the imiiioquinone or the reduced fori,r of the system. The combination with the iminoquinone appeared TABLE 5 Constants used in the calculations
--
p-METEYLAYINOPHENOL
CONSTANTS
EiL Er KO KR KO
KW
-
0 0 1 6 3 4 6
1
7402 volt 6888 volt X 1O-lL 54 x 10-9 x 10-11 16 X 10-l6 (10)
0.6877 volt (5) 1 x 10-10 (5) 1 x 10-8 (5) 6.45 X IO-” (14)
p-METEYLAMINOPEENOLSULFONIC ACID
0.7920 volt 0.7292 volt 5.88 X 8.50 X 2.51 X 10-1’ 6.16 X 10-l6 (10)
to be thc more likely, for titration of the mixture with sodium sulfite, which removed the oxidized form as monosulfonate, gave an end point which coincided with the disappearance of the purple color. Furthermore, the titer decreased as a function of time; this agreed with the disappearance of the oxidized form. The titration curves for these systems did not agree with the curve for a two-electron system. Fieser observed this and ascribed the deviation to the formation of meriquinone. The deviation found by the present writer was less than that observed by Fieser. The meriquinone formation will have no effect upon the measurements of potential at half oxidation unless the meriquinone is formed with other than a one-to-one ratio of oxidized and reduced forms. The effect of the introduction of the sulfonic acid grouping was to shift all ionizations of the compound, so that it was now both a weaker acid and a weaker base. This shift occurred with an increase in the oxidation-reduction potential a t any given pH.
OXIDATIOK-REDUCTION POTESTIALS
1225
Measurements upon the benzohydroquinone system were undertake11 in the hope that measurements in the range in which quinone was unstable might make possible an estimation of the value of the second ionization constant of the reduced form. The potential of the hydroquinone-quinone system should be given by the following expression :
E
=
Eo - 0.029(?) log,, &/So
+ 0.029 loglo (KiKz + Ki[H+]+ [H']')
Inspection of this expression, when Sn = So,shows that the slope for the dependence of the oxidation-reduction potential upon pH should be 0.058 until the pK of the Erst ionization is reached, at which point the slope should become 0.029. When the pK of the second ionization is reached, the slope should become zero. Examination of the graph of the experimental data in figure 3 shows that the change at the pK of the first ionization occurs, but that at about the point where the pK of the second ionization ihould be, there is a change in slope back to thc original 0.058. Furthcrrnore, the slope between pH 9.8 and pH 12.3 iq not the theoretical 0.029 but is 0.0261. This irregular slope indicates that in this region the potential is being determined by a mixture of' two systems. At pH 12.3 the berizohydroquinone system must entirely disappear as far as any effect upon the potentials is concerned, for this change in slope would indicate that a group, common to both forms of the syrtem, had now ionized in the oxidized form. The disappearance of the quinone in the range of pH above 10 was accompanied by the formation of a deep orange-red coloration which faded more or less rapidly, depending upon the pH value, to a yellow. The drift of potentials was rapid, and above pH 11.5 more consistent results were obtained by extrapolating a graph of the antilogarithm of the potential as a function of time. This arbitrary extrapolation agreed closely with the best extrapolation that could be made from the graphs of potential as a function of time. Measurements in this region were made both by adding a ferricyanide solution to a hydroquinone solution and by adding a saturated aqueous solution of quinhydrone to the buffer. The only system which it appeared might exist, even temporarily, ill this region and which might show this behavior was the hydroxyhydroquinone system. The dismutation of quinone is said toorcur in the follorving manner (2, 9, 12):
This reaction may then be followed by (11):
1226
A. E. CAMERON
Hydroxyhydroquinone in equilibrium with hydroxyquiiione might thus exist in the reaction mixture, although not necessarily in an equimolar mixture. Since this system is formed without participation of hydroquinone, mi experiment was run in which a quinone solution was added to the h f f e r instead of a quinhydrone solution. The valuc of potential for this experiment (cf. table 3) checked perfcctly with those obtained with quinhydrom, an& the same color sequence appeared in the solution. X measured quantity of acidified benzoquinone solutior, was reduced with hydrogen and colloidal palladium a t p H 0.5 and titrated with potassiuna rnolybdicyanide solution in an inert atmosphere. The same quantity of quinone solution was then added to 0.1 M sodium hydroxide in an atmosphew of nitrogen, and after 30 min. the solution was acidified with sulfuric acid : C I a pH of 0.5 and titrated with molybdicyanide solution. The titration jndicated that 9.2 per cent of a compound 0.100 volt more negative in on-reduction potential than the benzohydroquinone system, and 67.7 y r cent of hydroquinone had been formed from the quinone upon disniiii rtriori in alkaline solution.
CONSTANTS
BENZOHYDROQUINONE
HYDROXYEYDROQUINONE
_ _ _ _ I _ _ _ _
E o ,. . . . , .
..
,
.
K I .. . .
KO.. . . . . . . . . . .
0.7049 volt 0.7029 volt (1) 1.46 X 10-lo 1.33 X (4) 1.75 X l0-lo (14)
0.6014 volt
0.5960 volt (3)
10 X 10-4
II__...._.
Measurements upon the hydroxyhydroquinone system were carried out, but in this case, also, the limiting factor was the stability of the quinone. The measurements shown in figure 3 indicate that a group in the oxidized form has a pK of approximately 4.0, for above this p H value the slope has changed from 0.058 to 0.087. Above this pH value the solution is orangered in color, and below it, it is yellow. The quinone was not found to be stable in any solution used. As the solutions %ere made more and more alkaline, the dismutation of the quinone became more rapid, until above pH 7 the potential-time curves could not be extrapolated to zero time. The potentials showed an initial rise which was very rapid and which iiidicated formation of the hydroquinone from the quinone. The constants determined for the two systems are given in table 6. KO great significance is to be attached to the discrepancy in the value of Eo for the benzohydroquinone system, for e:xtreme precision was not being sought in this case. The chief interest in the investigation lay in the alkaline solutions. Thus it appears from these investigations that the method introduced
OXIDATIOX-REDCCTION POTEKTIALS
1227
by Fieser is capable of extension and application to a great many organic systems ordinarily regarded as irreversible. The greatest success will be obtained with systems which do not give rise t o the reduced form of an oxidation-reduction system upon decomposition of the oxidized form. SUMMARY
1. The oxidation-reduction potentials of the systems arising from p-methylaminophenol and from its monosulionic acid have been measured over a wide range of p H by the discontinuous titration method of Fieser. 2. Measurements have been attempted upon the benzoquinone system in alkaline solution and upon the hydroxyhydroquinone system, but such measurements were hampered by the production of the hgdroquinones upon disinutation of the unstable quinones. REFERENCES (1) (2) (3) (4) (5) (6) (7) (8) (9) (10)
(11) (12) (13) (14)
BIILMANN, E . : Ann. chim. 16, 109 (1921). BOGISCH, A.: P h o t . Korr. 37, 273 (1900). J. B., AND FIESER, L. A , : J. Am. Chem Soc. 46, 1858 (1924). CON.~NT, YON EULER,H., A N D BRGSIGS,E.: Z.physik. Chem. 139, 615 (1928). FIESER,L. F.: J. Am. Chem. SOC.62, 4915 (1930). FIESER,L. F.: J. Am. Chem. SOC.62, 5204 (1930). FIESER, L. F., A K D FIESER, RI.: J. Am. Chem. SOC.66, 1565 (1934). FIESER, L. F., AKD FIESER, XI.: J. .4m. Chem. SOC.67, 491 (1935). HESSE,0.: Ann. 220, 367 (1883). International Critical Tables, Vol. VI, p. 152. LIcGraw-Hill Book Co., Inc., Kew York (1929). LEUBNER, A.: Dissertation, Dresden, 1911. SCWEID, B.: Ann. 218, 227 (1883). SCHERING: German patent 208,434; Chem. Zentr. 1909, I, 1367. S. E . : Trans. Am. Electrochem. SOC. 39, 440 (1921) SHEPPARD,
