591
V O L U M E 25, NO. 4, A P R I L 1 9 5 3 The electrodes last on an average for 6 weeks with a minimum of 3 weeks. The drift circuit dry cell lasts a minimum of 1 week, and an average of 3 weeks; the other 2 dry cells usually last for the entire life of the electrode. Exchange interruption lasts about 15 minutes. ACKNOWLEDGMENT
In the later phases of this development work, the author was ably and devotedly assisted by Frank J. DeLuca of the Mines Safety rlppliance Research Laboratory and in the re-engineering by A. C. McInnes of the Engineering Department, to both of whom a feeling of deep gratitude is herewith expressed. LITERATURE CITED
(1) Berl, W. G., Trans. Electrochem. Soc., 83,253 (1943). (2) Glasstone, S., “Electrochemistry of Solutions,” 2nd ed., p. 335, Kew York, D. Van Nostrand Co., 1937. (3) Glasstone, S , “Textbook of Physical Chemistry,” p. 928, New Tork, D. Van Sostrand Co., 1940.
Jacobson, 11.G. (to Mine Safety Appliances Co.), U. 5. Patent 2,156,093 (.May 2, 1939). Ibid., 2,464 087 (March 8, 1949). Ibid., 2,540,674 (Feb. 6, 1952). Ibid., application pending. Kohlrausch, F., “Pratische Physik,” 17th ed., p. 569, Sew York, Rosenberg, 1944. Kordesch, K., and hfarco, -i., Nikrochemie ver. Microchim. Acta, 36, 420 (1950). Moysseef. A. S., and Brickman. N. M..APPZ. .. Chem. (U.S.S.R.). 12, 620 (1939). Paris, 9.,Chimie & industrie,31, 253 (1933). Weisa, R. S I and daffe, S. S., Trans. Electrochem Soc., 93, 128 (1948) Wiener Isolierrohr, Batterie &- hletallwarenfabrik, Austrian Patent 167,840 (March 10, 1951). Yant, W. P., Jacobson, M. G., and Strange, J. P. (to Mine Safety dppliances Co.), U. S. Patent 2,401,287 (May 28, 1946). RECEIVED for reiiew J u l y 28, 1953. Accepted January 12, 1953. Presented before the Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, March 1952.
Dual Intermediates in Coulometric Titrations Equilibria in Copper( +Bromide
Solutions
PAUL S. FARRINGTON: DALE J. MEIER, AND ERNEST 13. SWIFT California Institute of Technology, Pasadena, Calif. The work has a twofold purpose: to demonstrate the feasibility and usefulness of alternately electrolytically generating oxidants and reductants during the course of coulometric titrations, and to show the application of amperometric measurements made with polarized platinum electrodes to the determination of equilibrium constants. Experiments have shown that having a solution of cupric copper and bromide, one can electrolytically generate a known excess of bromine and then quantitatively reduce this by electrolytically generated cuprous copper, or similarly produce an excess of cuprous copper and reduce i t by bromine. Amperometric measurements can be used to determine very small concentrations of bromine and of cuprous copper. The results demonstrate the possibility of making coulometric back-titrations where an excess of reagent is required because of the slowness of the main titration reaction. A value has been obtained for the formal potential of the half-cell reaction CuBr-- = C u + + 2Bre-.
+
+
A
Ti AMPEROMETRIC method, similar in principle to the
dead-stop method of Foulk and Bawden ( 5 ) ,but modified to make use of the measurement of the current flow between two similar platinum electrodes with impressed potential differences ranging from 50 to 300 mv., has been employed for determining the end point of secondary coulometric titrations involving the electrolytic generation of bromine (10, 12), iodine (11). chlorine (S), and cuprous copper (9). Experience has shown that in many cases this method has certain advantages as compared with conventional end-point procedures. Thus, constant readings are attained more rapidly than is the case with potentiometric procedures, and titrations can be made more rapidly. The latter advantage is emphasized by the linear relation between current and concentration of titrant, since the titration can be made rapidly, the equivalence point overrun, and the excess of titrant accurately 1
Present address, University of California, Loa Angeles, Calif.
determined. rllthough concentrations of the above titrants of the order of 1 0 - 7 formal are readily determined, only limited use has been made of this fact for either analytical or physicochemical studies. In addition, the method is unusual in that continuous measurements of concentrations of the above magnitude can be made over extended periods by means of an ordinary microammeter and, although a significant current is continuously passed through the solution during such periods, the composition of the solution remains unchanged because the same half-cell reaction is proceeding in opposite directions a t the two electrodes. The above considerations have seemed to justify amperometric studies of those half-cell reactions of probable value for either analytical or physicochemical measurements. The results of such a study, made with solutions containing both cupric copper and bromide ion are presented below. Bromide solutions have been used for the electrolytic generation of bromine as an oxidant for secondary coulometric titrations, and solutions containing cupric copper have been similarly used for the generation of cuprous copper as a reductant. I n certain titrations with bromine, such as the determination of bromine numbers, it is necessary to add an excess of bromine, to wait for completion of the titration reaction, and then to back-titrate the excess bromine. Since it seemed that this might be done coulometrically in solutions containing both copper( 11)and bromide, a study was undertaken to investigate the indicator currents resulting from alternate anodic and cathodic generation in solutions containing these substances as dual intermediates. In the course of this study it was observed that the indicator current never decreased t o zero but always had an appreciable minimum value. If it is assumed that this minimum indicator current is caused by the cuprous copper and bromine produced by the reaction, 2Cu++
+ 7Br-
= 2CuBrz-
+ Br3-
it is possible to derive a relationship between the minimum indicator current and the equilibrium constant for the above reaction. I n this derivation both anode and cathode reactions are considered. Kolthoff and Lingane ( 6 ) have made similar calcula-
ANALYTICAL CHEMISTRY
592 tions for the silver iodate solubility equilibrium in which only cathode reactions are involved. The cuprous copper and bromine concentrations in equilibrium with cupric copper and bromide are measured by means of the amperometric indicating circuit. I n order to do this the sensitivity of the electrodes must be determined for each of the reactants. This has been done by means of the coulometric apparatus previously described (9).
up to a volume of 45 ml. with water. With apotential of 100 mv. applied to the indicator electrodes, the indicator current was carefully adjusted to its minimum value by alternate generation of bromine and cuprous copper. The indicator potential was then varied and corresponding value3 of the indicator current were recorded. Aplot of indicator current 1's. indicator potential is shown in Figure 1. Inspection of Figure 1 shows that an indicator potential to about 60 mv. is most suitable for this equilibrium study in f h a t the current is least affected by potential change in that region.
EXPERIMENTAL
Reagents. All chemicals used were reagent grade.
A stock solution of 1.01 F (volume formal) copper sulfate was prepared from copper sulfate pentahydrate. A tenfold dilution of the stock solution was standardized iodometrically. The sodium thiosulfate solution used for the copper titration was standardized against dried potassium iodate. Stock solutions of 1.00 F and 6.0 F sodium bromide were prepared by weight from the salt which had been dried a t least 2 hours a t 120" C. Apparatus. Except for the microammeter the apparatus was the same as that describedpreviously (9). The microammeter was replaced by a Leeds and Korthrup box-type, reflecting galvanometer. A shunt resistance of about 10 ohms made the galvanometer sensitivity 3 divisions per microampere. As a result the potential drop across the galvanometer was always much less than 1 mv. and the potential applied to the indicator electrodes was held constant without adjustment.
L Table I. ki
No.
NO.
7
8
6.75 cut+, M
0.0318 0.0152 0.143 0.0206 0,0059 0.0206 0.0152 0.0138
(Div./lO-7 P ) 5 44 5.79 5.65 5.94 7.15 7.35 4.69 3.65
Br-, M 0 90 0 95 0 69 0 66
0.98
n
fifi
0.95
0.95
c u +, F
i4f
k4
(Div./lO-6 F ) 9.10 9.95 5.40 6.16 8.65 8.41 6.16
8
1 2 3 4 5 6
C o n s t a n t for Cupric Bromide OxidationR e d u c t i o n Equilibrium +
(Div.ja 7.5 9.8 10.6 3.6 5.1 4.3 8.6 10.2
K1
x
1020
0 112 0 0561 0 404 0 0561 0 0224 0 0561 0 0561 0 0496 K2
1.39 2.42 10.7 0.306 0.366 0.226 5.29 2.48
x
104
Br-. F 1 00 1 00 1 00 0 70 1 00 0 70 1 00 1 00 K x lOL~
4.84 1.61 1.5. 3 0 232 0.304 0.232 1.61 1 33
hv.
8
1 scale division
0.29 1,50 0.70 1.32 1 20 0.97 3.28 1,86 1.39
= 0.33 microampere.
Selection of Potential Difference Applied across Indicator Electrodes. When bromide is used independently as an intermediate for coulometric titrations (IO),a potential difference of 200 mv. is applied across the indicator electrodes. This potential does not cause continuous oxidation of bromide ion a t the indicator anode unless bromine is available a t the indicator cathode; under these conditions the indicator current is a direct measure of the bromine concentration. Similarly, in solutions containing relatively large concentrations of copper(I1) and relatively small concentrations of copper(I), the diffusion of the latter to the indicator anode a i l 1 limit the current, and the indicator current is a direct measure of the cuprous copper concentration. However, if cupric copper and bromide are both present, the following pair of indicator electrode reactions may contribute to the indicator current: Anode, 3 Br- = Bra2e-
+
Cathode, C u + +
+ 2Br- + e -
=
CuBrr
The following experiment was performed to establish the proper indicator potential for the equilibrium studies:
A solution containing 5 ml. of 1.01 F copper sulfate, 5 ml. of 3 F sulfuric acid, and 7.5 ml. of 6.0 F sodium bromide was made
I
I
I
I
I
I
I
PO
40
60
80
100
120
140
1
APPLIED P O T E N T I A L DIFFERENCE, MV.
Figure 1. Indicator C u r r e n t vs. Applied Potential Difference Solution 0.112 F i n CuSO4 and 1.00 F i n NaBr, initially adjusted to minimum current at 100 m v .
Procedure for Obtaining Equilibrium Data. I n view of the above experiment, the indicator potential was set at 60 mv. Measurements for the equilibrium calculations were made in solutions 0.33 F in sulfuric acid and with a total volume of 45 ml. By means of pipets, copper sulfate and sodium bromide stock solutions were added to a titration cell (see Table I for experimental concentrations). Bromine was generated in the cell and values of indicator current were recorded after each second of generation. The polarity of the generating electrodes was then reversed, and generation was continued until the indicator current had passed through the minimum point. When excess cuprous copper was being generated, values of the indicator current were recorded after each 2 to 4 seconds of generation. The value of the minimum indicator current was determined by passing very slowly back and forth through the minimum point until a reproducible reading was obtained, Figure 2 shows a typical indicator current curve in the region of the minimum point. The data for this curve were obtained from sample 8. About 2.5 X lo-' equivalent of cuprous copper was produced electrolytically in the solution, then bromine was generated for short intervals of time and the indicator current was observed after each period of generation. Indicator Current Mechanism a t Minimum Point. In order to interpret the current measurements in the vicinity of minimum current it is necessary to determine whether both cuprous copper and bromine contribute to the indicator current or whether only one of these substances controls, even though comparable concentrations of the two intermediates may be present. A solution corresponding to Xo. 8 in Table I was prepared. The potential difference between a platinum electrode in the solution and a saturated calomel electrode was measured for various indicator current values. The cupric bromide solution was replaced by a solution having the same concentration of bromide ion (but no copper). Bromine vias generated in this solution and the potential was again measured against the saturated calomel electrode for various values of the indicator current. When the indicator current in the cupric bromide solution was adjusted to the minimum value of 1.2 microamperes, the solution potential was 0.585 volt (us. calomel). When the sodium bromide solution was brought to the same potential by generation of bromine, the indicator current was 0.4 microampere. I n each solution the indi-
593
V O L U M E 25, NO, 4, A P R I L 1 9 5 3 cator current was raised to 15 microamperes by generation of bromine; the potentials of the cupric bromide and sodium bromide solution were 0.651 and 0.652 volt, respectively.
and,
I n the first case above, the bromine concentrations are shown to be equal by the equality of the solution potentials, but the indicator current given by the bromine is not sufficient to account for the indicator current observed in the cupric bromide solution. That the bromine concentrations are the same at a given solution potential is demonstrated in the second case by the equality of solution potentials and indicator currents (15 ka.) in a region where the bromine alone controls the indicator current.
The oxidation-reduction equilibrium which exists in cupric bromide solutions may be represented by the equation
DISCUSSION
Calculation of Cupric Copper Bromide-Cuprous Copper Bromine Equilibrium Constant. When only one intermediate system is present in a solution, and one of the components of this system is present in a relatively low concentration, the indicator current is linearly proportional to the concentration of that component. In a cupric bromide solution two intermediate systems are present and in the region of the minimum current the evidence presented above indicates that both intermediate systems are contributing to the indicator current. Therefore, the indicator current may be represented by a linear combination of the cuprous copper and the bromine concentrations: Zli[Cu(I)]
z
= kl(CuBra-)
For the equilibrium
BrZ Latimer (?') gives K
=
+ k:,(Br?) + ka(BrS-)
+ Br-
2Cu++
(')
+ 7Br-
+ k4(Br3-)
= 2CuBrl-
(2)
+ BrS-
therefore
(3) Inasmuch as the cupric and bromide ion concentrations are e s sentially constant because of the small amounts of CuBrz- and Bra- which are formed, the product (Cu++)l (Br-)'
=
KZ
is a constant for any given solution. The concentrations of CuBrz- and Bra- may then be related by the constant,
K1 = KKZ
=
(CuBr2-)2(Brg-)
(4)
Combination of Equations 2 and 4 gives the indicator current as a function of the CuBrz- concentration
i = kl(CuBr2-)
+ Zk(Br2)
The solutions under consideration contain sufficient bromide that the cuprous copper exists primarily as CuBr2- and the bromine ie present both as-Brp and as Bra-; therefore,
i
i = kl(CuBrz-)
k Ki + (C'uBr2-)2
(5)
Differentiating this expression with respect to time of generation and combining terms
di & = [kl
-
2k4K1 I] [it(CuBrZ-)] (CuBrZ-) dt
At the point of minimum current di - 0, but d(CuBrl-)# 0 dt dt
= Bra-
17
Therefore, a t the minimum point,
The substitution of (Br2) = (Br3-) into the second term of Equal?(Br -) tion 1 yields Zk(Br2) = (Br3-)
+ ka] [+7 (Br-) k
and (CuBra-).w =
The concentration of bromide can be considered a constant for any given experiment because the anlounts of bromine formed are very small; therefore,
Substitution into Equation 5 and collection of terms give a value for the minimum indicator current, i.~,
The constants li1 and kr can be evaluated by observing the change in indicator current per unit of generation time &hen the concentration of one of the indicating ions is very small. If R = rate of generation (equivalents per second) J' = volume of solution (liters) and (CuBrz-) >>(Bra-) differentiation of Equation 2 with respect to generation time gives
(2) /di\
(2) ~
0
5
10
15
PO
95
30
35
GENERATION TIME, SECONDS
Figure 2.
Behavior of Indicator Current in Minimum Point Region
40
-~ -~( d i / d t ) v
d(CuBr2-) -
R
dt
Similarly, when (BPI-)
>> (CuBrp-)
594
ANALYTICAL CHEMISTRY When the potential value EO = -0.173 for the half-cell given above is combined with the data of Lewis and Randall as given by Latimer ('7, p. 173)
However,
d(Bra-) - .R _ dt - 21;
17(Br-)
~
because both Big and BPI- are produced by generation.
Then
k4
=
2v ~
(2)
[1
+ 17(Br-)] 17 (Br -)
I n Figure 2 it is evident that the indicator current curve becomes a straight line a few seconds from the minimum point. On the left side of the minimum point, for example, the indicator current is linearly proportional to the CuBrn- concentration, and kt can be calculated from the observed value of the slope of that portion of the curve together with the generation rate, R, and the volume V . In similar fashion the slope of the indicator current curve to the right of the minimum point is used to evaluate ka. By the substitution of the calculated values of kl and kc and the observed value of i~ into Equation 6, the equilibrium constant K I can then be calculated. Experimental values of K1 are listed in Table I. Before calculating K2 it was necessary to correct the formal concentrations of cupric and bromide ions for the formation of undissociated complexes. Farrington (2) has determined a constant for the equilibrium
+ Br-
= CuBr+
K = 2.1 As an approximation to the value of the constant for the formation of the dibromide complex, CuBr+
+ Br-
= CuBrz
it was assumed that the constants for the two bromide complexes would have the same ratio as the constants for the corresponding chloride complexes. From the work of McConnell and Davidson ( 8 ) one finds a value of 5.65 for the ratio of the monochloride constant to the dichloride constant. By applying this ratio the approximate value of K = 0.37 was obtained for the dibromide equilibrium given above. Higher bromide complexes of copper were assumed to have a negligible effect on the cupric and bromide ion concentrations, The molal concentrations of cupric and bromide ions were calculated from the formal concentrations by accounting for the two complexes, and values of K Zwere calculated from the molal concentrations. Values of K Zas Re11 as the formal and molal concentrations are listed in Table I. For the equilibrium,
values of K for various cupric and bromide ion concentrations will be found in Table I. From the average value, K = 1.39 X the value of E = -0.468 is obtained for the cell reaction given above. If the value Eo = -1.05 is taken ( 7 ) for the half-cell
3Br- = Bra-
+ 2e-
then one calculates Eo = -0.58 for the half-cell CuBrp-
=
Cu++
+ 2Br- + e -
Combination of the following data,
+ eCuBrz- = Cu+ + 2BrCu = Cu+
(4; 7 , p. 171) Eo = -0.522 K = 1.2 X 10-6
yields a value Eo = -0.173 for the half-cell 2Br-
+ 2e-
Eo = -0.3448
the value EO = -0.52 is obtained for the half-cell CuBrZ- = C u + +
R
Cu++
Cu = C u + +
+ Cu = CuBr2- + e -
Latimer (7) gives a value of Eo = -0.05 for this half-cell, but this value is not consistent with the above data from this reference.
+ 2Br- + e-
.4 possible source of error in these calculations could be an indicator current given by the following electrode reactions: Indicator anode 3Br- = Bra2eIndicator cathode C u + + 2Bre - = CuBrzHowever, the production of significant quantities of bromine or of cuprous copper at the anode and cathode, respectively, would result in the cell potential becoming greater than the applied potential since one would expect the diffusion gradient of these substances from the electrodes to be small. As a result, the total indicator current should remain controlled by the concentrations of bromine and of cuprous copper in the solution. In addition, the curve shown in Figure 1 is similar in shape toa polarographic wave, and it seems reasonable to assume that the indicator current is primarily diffusion controlled and that the decomposition current is not a major portion of the total indicator current. Activity coefficients have not been taken into account in these approximate calculations, and the work has been done a t prevailing laboratory temperatures. Location of Equivalence Point of Coulometric Titrations. Khen copper(I1) and bromide are used as dual intermediates in coulometric titrations the equivalence point does not coincide with the point of minimum current. From Equation 4, if (Bra-) = '/z(CuBrz-),
+
+
+
(CuBr2-)eq,pt. = (2Kd1I3 Substitution of this value into Equation 5 gives the indicator current a t the equivalence point
This value is not identical with that obtained from the expression for the minimum indicator current. The minimum indicator current would correspond to the equivalence point only for a system in which the electron changes of the oxidation and reduction reactions were equal and the electrode sensitivities were identical. I t is evident in Figure 2 that, in the region of minimum indicator current, the change of indicator current with time of generation is comparatively small; therefore, the minimum indicator current is not a satisfactory reference point for coulometric titrations. The best point of reference is the lowest current value which falls on a straight line portion of the indicator current curve. Higher reference current values will tend to introduce errors due to ( a ) magnification of small differences in electrode sensitivity a t high indicator currents (see Figure 4,IO), and ( b ) air oxidation of copper(1) or volatilization of bromine. Other factors being equal, a point on the bromine diffusion controlled portion of the curve is preferable as a reference point because of the steeper slope. The use of an arbitrary value of the reference current involves no complication of procedure because that same current value can be taken as the end point of the blank test made on the reagents; all titrations can then be made to that reference point. The results of a study of an application of the above dual intermediate system to the coulometric titration of aniline have been presented ( I ). ACKNOWLEDGMENT
The authors are indebted to Tucker Carrington for valuable suggestions in connection with the study of the current mechanism in the region of the current minimum and to Keith Rowley for helpful comments and measurements made during this study.
V O L U M E 25, NO. 4, A P R I L 1 9 5 3 One of the authors (P.S.F.) wishes to acknowledge the assistance of the Merck Graduate Fellowship in Analytical Chemistry
(1949-50).
595
Latimer, R. M., “Oxidation States of the Elements,” p. 53, New York, Prentice-Hall, 1938. McConnell, H., and Davidson, K.,J . Am. Chem. S O C ,72, 3164 (1950).
LITERATURE CITED (1) Buck, R. P., and Swift, E. H., ANAL. CHEM.,24, 499 (1952). (2) Farrington, P. S., J. Am C h m . Soc., 74, 966 (1952). (3) Farrington, P. S., and Swift, E. H., ANAL. CHEM.,22, 889
(1950). Fenwick, F., J . A m . Chem. Soc., 48, 860 (1926). (6) Foulk, C. IT., and Bawden, A. T., Ibid., 48, 2045 (1926). (6) Kolthoff, I. hl., and Lingane, J. J., pp. 455-8, “Polarography.” Yew Tork Interscience Publishers, 1946. (4)
hleier, D. J., hlyers, R. J., and Swift. E. H., Ibid., 71, 2340 (1949). Myers, R. J., and Swift, E. H., Ibid.,70, 104T (1948). Ramsey, W. J.,Farrington, P. S., and Swift, E. H., ANAL. CHEM.,22, 332 (1950). SeasP, J. W., Kiemann, C., and Swift, E. H., Ihid., 19, 197 (1947). R E C E I V ~for D review August 27, 1951. Accepted January 24, 1953. Contribution No. 1623 from the Gates and Crellin Laboratories of Chemistry, California Institute of Technology, Pasadena, Calif.
Differentiation of Chelating from Nonchelating Phenols SAUL SOLOWAY AND PERRY ROSEN The City College, College of the City of New York, New York, N. Y . As a result of previous observations in connection with a modified ferric chloride color reaction for phenols, a test has been developed to differentiate chelating from nonchelating phenols. This test is based on the fact that ferric complexes of chelating phenols have a wider range of pH stability than the others. It is shown that the stability of the phenolic-ferric ion complexes varies considerably with the solvent. Results indicate that acetic acid shows promise as a competitive agent against certain types of weakly chelating phenols, and hence may be used to differentiate members of the chelating class from one another. The reason for the differences in chelating strength of some of the common orthosubstituted phenols with ferric ion is discussed briefly.
A
LTHOUGH it is well known that certain ortho-substituted phenols form stable heterocyclic structures with many metallic ions, no simple chemical procedure is available to distinguish such chelating phenols from those not capable of chelation. Practically all phenols give color reactions with ferric chloride under certain conditions (8). However, the chelating types yield complexes which are more stable over a wider range of environment. This fact serves to differentiate the chelaters from the nonchelaters. I t was recorded (8) that acetic, benzoic, and p-hydroxybenzoic acids could inhibit the formation of the typical color reaction between m-cresol and ferric chloride in some solvents, but not with salicylic acid under the same conditions. This observation led to the thought that carboxylic acids might form complexes with ferric ion of a stability intermediate between those of chelating and nonchelating phenols. The procedure developed on the basis of this idea proved to be only partially successful-that is, acetic acid apparently can compete against any nonchelating phenol for ferric ion, but this was also the case for some wellknown chelating derivatives. A second scheme based on the greater p H range stability of chelated phenol-ferric ion complexes proved to be successful for carrying out the differentiation. This test used in conjunction with the first one allows not only the differentiation of chelaters from nonchelaters, but also the distinction of certain chelaters from one another on the basis of their relative stabilities toward acetic acid. EXPERIMENTAL
The tests using acetic acid as a competitive agent against phenols for ferric ion were carried out in five solvents: water. methanol, benzene, diethyl ether, and 2,2’-dichlorodiisopropyl ether. This variation served a twofold purpose: to provide at least one good solvent for all the compounds tested and to furnish some information on the variation of the stability of the complexes with solvent. In the second set of tests a phenol was added to a suspension of
ferric hydroxide in water and methanol. These suspensions were prepared by adding aqueous ferric chloride to dilute solutions of pyridine in water and methanol. The pH, as determined with a set of pHydrion indicator test papers, was found to be in the range 4.6 to 4.8. It was noted (Table 111) that under these acid conditions most of the chelating phenols gave stable colors with the ferric ion present, whereas none of the nonchelaters did. On further acidifying the solutions to the extent of dissolving the excess ferric hydroxide, it was noted that all the chelaters gave colored solutions and/or precipitates easily distinguishable from a blank of ferric ion a t the same p H (1.7 to 1.9). The nonchelaters gave the same color as the blank. These results, which are given in Table 111,demonstrate the greater pH range stability of chelating phenol complexes with ferric ion. MATERIALS AND REAGENTS
The reagents were made up of the best grades of materials commercially available. All the organic solvents were anhydrous. The phenols and other derivatives tested were mostly commercial chemicals. Some few were synthesized. In such cases the purity as judged from physical constants was at least the equal of the materials described in the literature. Solution 1. Ten grams of anhydrous ferric chloride were dissolved in 100 ml. of water. The solution was filtered. Its color was orange. Solution 2. Ten grams of anhydrous ferric chloride were dissolved in 100 ml. of absolute methanol. The solution ~ - y & 9filtered. I t s color was orange. Solution 3. One-half gram of anhydrous ferric chloride was suspended in 100 ml. of anhydrous benzene. After the suspension was shaken for a short time, the solution was allowed to stand and then was separated from the undissolved material by decantation. The color of the solution was dark brown. The solution was not filtered because a trial run gave a filtrate with a muddy brown appearance, presumably due t o the uptake of a small amount of water from the paper with the possible formation of some insoluble hydrate. Solution 4. Five grams of anhydrous ferric chloride were dissolved in 100 ml. of peroxide-free diethyl ether. The solution was filtered. Its color was dark brown.
