analog of triphenylamine, is titratable in the anhydride while the nitrogen compound is not. No base constant was calculated, as there was some evidence of acetylation of the phosphine. Quaternary ammonium halides all gave quantitative titrations. Calculated pK. values are in agreement with literature values. Chlorides and iodides could be resolved as shown in Figure 5. The sample titrated contained 20.4 mg. of tetramethylammonium chloride and 91.9 mg. of tetrabutylammonium iodide. The amounts found were 20.4 and 90.7 mg., respectively. Chlorides and bromides, or bromides and iodides, give only one break in the titration curve. Bisulfates could not be titrated. This may be due either to the high acidity of sulfuric acid or to the formation of a superacid by reaction of solute and solvent. Russell (IO) has noted a reaction between the anhydride and sulfuric acid to yield a mixed anhydride of high acidity: (CH&0)20
+ H2SO4
e
-+
0 0
I1
CHa OS-OH
II
+ CHsCOiH
Titration of monosodium phosphate resulted in a titration curve indicating the presence of a very weak base, pK. -10.8, which is as weak as iodide ion, In aqueous solution the dihydrogen phosphate ion has a PI(, of 1.96. Reaction with the solvent may have produced a superacid: HzPOI-
+ 2(CHsCO)ZO
(CHaC0)zP-0-
II
+ (CHpC0)20
CH~C-ONHI
-+I
0
I/
+
CHsC-NH1 2CHaC02H (4) The titration of aniline hydrochloride is quantitative. That for ammonium acetate is not, implying the reaction may proceed partially to imide formation.
+
LITERATURE CITED
0
//
0
+ 2CHSCOzH
(1) Edwards, J. O., J . Am. Chem. SOC.
(2)
When trisodium phosphate is titrated, two breaks appear in the titration curve. The first end point occurs after the addition of two equivalents of acid, and the second after addition of the third equivalent. The titration curve again illustrates formation of a superacid for the final part of the titration. More condensed molecules should be stronger acids than parent compounds, as illustrated by the strength of pyro acids over the less condensed parent. Other salts such as aniline hydrochloride and ammonium acetate react with the solvent to give amides, which are then the entity titrated.
(1)
0
The corresponding base would be, of course, exceedingly weak.
76,1540 (1954).
12) Fritz, J. S., ANAL. CHEM.25. 407
(1953). ‘ Fritz, J. S., Fulda, M. O., Ibid.,25, 1837 (1953). Gremillion, A. F., Ibid., 27, 133 (1955). Hall, N. F., J. Am. Chem. SOC.52, 5115 (1930). Kolthoff, J. M. Furman, Tu’. H., “Potentiornetr{c Titrations,” p. 329, Wiley, New York, 1926. Lange, N. A., “Handbook of Chem(7) istry,” 4th ed., pp. 1220-1, Handbook Pub., Sandusky, Ohio 1941. Lemarie, H., Lucas, H. J., Am. Chem. SOC.73, 5198 (1951). Pifer, C. W., WoIlish, E. G., Schmall, hl., J. Am. Pharm. ASSOC., Sci. Ed. 42, 509 (1953). Russell, J., J. Am. Chem. SOC.60, 1345 (1938). Seaman, W., Allen, E., ANAL.CHEM. 23, 592 (1951). (12) Streuli, C. A., Ibid., 27, 1827 (1955). (13) Usanovich, M., Yatsimirskii, K., J . Gen. Chem. U.S.S.R. (Eng. Transl.) 11, 954 (1941). RECEIVED for review August 22, 1957. Accepted December 14,1957.
f.
Amperometric Titration of Fluoride with Thorium Using a Rotating Palladium Electrode W. E. HARRIS Department o f Chemisfry, University of Alberta, Edmonton, Alberta, Canada )The amperometric titration method for determining traces of fluoride is performed in a solution containing potassium bromide, potassium sulfate, and aerosol, and buffered with monochloroacetic acid and sodium monochloroacetate. Best results are obtained with about 100 to 200 y of fluoride per 100 ml. of solution, although as little as 20 y in 100 ml. can b e determined. Moderate amounts of chloride, nitrate, sulfate, perchlorate, borate, calcium, or magnesium cause no interference but large amounts slightly decrease the sensitivity. Aluminum and more than 1 mg. per liter of phosphate interfere. A simplified, inexpensive apparatus is described.
A
of determining traces of fluoride was desired which would
METHOD
0
ANALYTICAL CHEMISTRY
be fast, accurate, and subject to a minimum number of interferences. The amperometric titration method d e veloped is performed in a solution containing potassium bromide, potassium sulfate, aerosol, and buffered with monochloroacetic acid and sodium monochloroacetate. Best results are obtained with about 100 to 200 y of fluoride per 100 ml. of solution, although as little as 20 y can be determined. Moderate amounts of chloride, nitrate, sulfate, perchlorate, borate, calcium, or magnesium cause no interference, but large amounts slightly decrease the sensitivity. Aluminum and more than 1 mg. per liter of phosphate interfere. A simplified, inexpensive apparatus suitable for these fluoride analyses is described. The principal methods for the deter-
mination of fluoride involve either the measurement of the bleaching effect of fluoride upon a colored organic salt of a metal such as thorium or zirconium or the titration by standard thorium solutions using an indicator such as Alizarin Red S (IO, IS). Conditions for carrying out the measurements must be carefully controlled, and many substances may interfere with either the development of the color or the detection of the end point. A number of amperometric methods (2, 6, 7 ) use the dropping mercury electrode, but in general these are not conveniently adaptable to the determination of very low concentrations of fluoride. Recently Shoemaker (9) described a polarographic method for the determination of fluoride using the iron fluoride complexes. These complexes are
irreversibly reduced a t a rotating platinum electrode a t more negative potentials than the uncomplexed iron. At an appropriately chosen potential the resultant decrease in diffusion current can be related to the amount of fluoride by a set of calibration curves. Sundaresan and Karkhanavala (11) described an amperonietric titration of thorium with fluoride using a dropping mercury electrode in which ferric perchlorate was used as an amperometric indicator. The present work has combined and extended these two methods, resulting in the use of a rotating electrode. Very low concentrations of fluoride can be titrated with thorium using ferric chloride as an amperometric indicator. The method is sensitive and convenient. The fluoride is titrated with a standard solution of purified thorium perchlorate. During the titration the diffusion current due to iron is measured a t +0.2 volt us. the saturated calomel electrode (9). The thorium reacts with the ferric fluoride complexes to form the more stable thorium fluoride complexes (10) and the ferric iron released causes an increase in the diffusion current. At the end point, when no additional iron is released, the diffusion current becomes constant. A simple amperometric titration apparatus was used in which the potential of +0.2 volt to the rotating electrode was supplied by a specially devised cell. The following conditions have been found desirable for the titration: use of palladium electrodes in preference to those made of other metals to reduce the deleterious effects of some impurities; addition of potassium bromide to maintain a low background current; addition of potassium sulfate to reduce the undesirable effects of traces of phosphate from the solution or reagents; addition of a buffer of monochloroacetic acid and sodium monochloroacetate to buffer the solution to a p H of about 2.8; and addition of Aerosol 1B to keep the electrode in good operating condition (8). EXPERIMENTAL
Reagents. Purified thorium perchlorate was made by putting Fisher certified reagent thorium nitrate through two cycles of thorium oxalate precipitation, thorium oxide ignition, and perchloric acid solution. The final solution of thorium perchlorate was evaporated to perchloric acid fumes after the addition of a small amount of aqua regia to destroy the last traces of oxalate. This solution was diluted to about 0.05M and standardized gravimetrically by precipitating thorium oxalate and weighing the oxide (18). Solutions of 0.001M thorium perchlorate were prepared from the standard by dilution. A 0.1M solution of sodium fluoride
Table 1.
Effect of Variations in Buffer-Detergent Solution on Titrations of Fluoride with Thorium Perchlorate
Chloro-
Sodium PotasPotasChloroslum slum Background BensiAcid, acetate, Sulfate, Bromide, Mole/ Mole/ Mole/ Mole/ Aerosol, Current, tivity End Point Liter Liter Liter Liter % pa. ( ~ i / F-)a y Vol., M1. 0.03 0.015 0.003 0.03 0.1 -0.15 0.007 7.04 0.006 0.003 0.003 0.03 0.1 -0.18 0.012 6.93 0.1 -0.24 0.006 7.06 0.006 0.003 0.017 0.03 0.000 0.03 0.1 -0.20 0.014 6.0-6.4 0.006 0.003 0.006 0.003 0.003 0.17 0.1 -0.22 0.011 7.37 0.006 0.003 0.003 0.00 0.1 1-0.53 0.009 6.70 0.006 0.003 0.003 0.03 0.5 -0.20 0.011 7.03 0.006 0.003 0.003 0.03 0.0 -0.15 0.012 7.12 Sensitivity of an electrode was defined as change in diffusion current in pa. during titration per y of fluoride. acetic
(1
was prepared from C.P. Baker reagent. This solution was also standardized gravimetrically by precipitating lead chlorofluoride (1.2). Solutions of lower concentration were prepared from this one by dilution. Aerosol lB, a dibutyl sodium sulfosuccinate, was obtained from the American Cyanamid Co. Buffer-detergent solution was prepared from 19 grams of monochloroacetic acid, 2.7 grams of sodium hydroxide, 20 grams of Aerosol lB, 10 grams of potassium sulfate, and 80 grams of potassium bromide. Each was dissolved in separate portions of distilled water, combined, and diluted to 1 liter. Indicator solution was prepared by diluting 25 ml. of stock 0.01M ferric chloride in 0.001M hydrochloric acid to 1 liter with distilled water. Apparatus. A modified version of the amperometric titration apparatus described earlier (6) was used (Figure 1). A saturated calomel reference half cell, A , was made from a 15-ml. distillation flask with the side arm bent into shape for a salt bridge. The side arm was connected with a short piece of rubber tubing to a 1-cm. piece of 5-mm. porous glass ( I ) . The palladium indicator electrode, B , was made by sealing a piece of palladium wire into 6-mm. soft glass tubing. Because palladium amalgamates with mercury, electrical connection was made by fusing the palladium to a short piece of platinum in an oxygen flame. The palladium-platinum seal was covered
Table II.
with paraffin wax. Mercury was used to make contact with the end of the platinum wire. The electrode was rotated by a Sargent BOO-r.p.m. Synchronous rotator, C. The required potential was applied by the cell, D, made from two half cells. The more positive was a saturated potassium sulfate-mercurous sulfate half cell. The other was a phosphate half cell with an electrolyte made by dissolving 4 grams of monopotassium phosphate, 4 grams of dipotassium phosphate, and 5 grams of potassium nitrate in 50 ml. of water with an excess of mercurous phosphate. The salt bridge for cell D consisted of a U-shaped tube packed with glass wool and filled with potassium nitrate solution. The electrical circuit was completed through a lamp and scale galvanometer, E , with a sensitivity of about 0.03 pa. per scale division. Recommended Procedure. Ti ansfer 100 ml. of the neutral fluoride solution t o a 150-ml. beaker. Add 5 ml. of the buffer detergent solution, and place the beaker in position as shown in Figure 1. Add 4 ml. of indicator solution, and read the diffusion current. Titrate with standard thorium perchlorate solution in suitable increments, and read the diffusion current after each addition. Apply dilution corrections to the current readings and plot the values obtained. Draw two straight lines through the experimental points to obtain the end point. Relate the end point volume to the amount of fluoride by means of a linear
Effect of Impurities on Titration of Fluoride with Thorium Perchlorate
Impurity
Amt. Added
None Potassium chloride Pot,assium nitrate Sodium perchlorate Calcium bromide Magnesium chloride Borax Aluminum chloride Potassium phosphate
...
1 gram 1 gram 1 gram 1 gram
0.1 gram
10 mg.
0 . 5 mg. 0.1 mg.
0 . 2 mg. 0.3 mg.
1 mg.
Sensitivity (Ail7
F-)
0.012 0.009
End Point Vol., hll.
6.93 6.93 6.97 7.01 7.08 6.95 6.67 3.96 6.86 About 6 ml. Poorly defined, 4 6 ml. No end point
0.012
0.012 0.010 0.013 0.012
0.005 0.010
0.007
0.005
... ~~
VOL. 30,
NO. 5 , MAY 1958
1001
standard curve obtained by the titration of known amounts of fluoride. RESULTS AND DISCUSSION
Rotating Electrode. The usual rotating platinum electrode is not the most satisfactory for the titration of yery small amounts of fluoride. Impurities, probably phosphates, are present in most commercially available reagent grade thorium salts. These depress the diffusion current of ferric iron a t the electrode and foul the electrode surface. The undesirable effects appear to be least with electrodes made of palladium and to become increasingly greater r\-ith 90% platinum-lOyo rhodium, platinum, and gold. It is important therefore that the standard thorium solutions be prepared from the purified chemical as described, because addition of the impurity with the titrant may cause the diffusion current to decrease continuously after the end point. A new palladium electrode or one freshly cleaned with aqua regia may take an hour or more before the background current decreases to acceptable values. Fortunately, when an electrode becomes insensitive it almost never requires drastic cleaning with aqua regia. Full sensitivity can generally be restored by wiping it with a clean cloth. Buff er-Detergent Solution. A buffer containing two parts of monochloroacetic acid to one of sodium monochloroacetate was suitable. Satisfactory end points are difficult to obtain a t high pH values because of hydrolysis of the ferric iron. K i t h a p H of much less than 2, satisfactory titration curves are not obtained because of very lonelectrode sensitivity. Large amounts of the monochloroacetate buffer considerably reduce the sensitivity and hence should be avoided. Table I shows the effect of variations in bufferdetergent solution on titrations of 204 y of fluoride with 10-3M thorium perchlorate. While the inclusion of potassium sulfate in the buffer-detergent solution decreases the sensitivity of the electrode by interfering with the ferric iron reduction (Q), the effect is not serious with moderate amounts of this salt (Table I). A small amount of potassium sulfate is desirable for giving a sharper end point and for minimizing the effect of phosphate on the titration. Some of the data of Table I illustrate the function of potassium bromide in keeping the background current low. I t also appears that the addition of bromide increases the sensitivity. The detergent, Aerosol lB, kept the indicator electrode in a state of nearly constant and high sensitivity for long periods. Without the detergent the 1002
ANALYTICAL CHEMISTRY
electrode required more frequent cleaning and conditioning. Interferences. Because the titration is based upon the transference of fluoride ions from the stable ferric fluoride complexes to the more stable thorium fluoride complexes, there are only a few, if any, of the weakly complexing cations or anions that should interfere. Table I1 shows the effect of various impurities on titration of 204 y of fluoride in 100 ml. of solutionwith 10-3 M thorium perchlorate. On the other hand it is to be expected that some of the multivalent cations with very high ionic potential will form more stable fluoride complexes than does ferric iron and will therefore interfere in the titration-e.g., results obtained in titrations with an aluminum salt present are incorrect (Table 11). The fluoride-aluminum complex is sufficiently stable so that aluminum interferes a t all measurable concentrations. Because the fluoride complexes of a number of other cations (IO), particularly of elements in groups 3 and 4 of the periodic table, are stable, it is expected that some of these will also interfere. Also, anions which form stable complexes with iron and thorium will probably interfere. Phosphate is an example of such an ion, but fortunately traces do not give serious interference if the titration is carried out rapidly. If the titration is carried out slowly a slow reaction occurs and causes the electrode to become increasingly insensitive or fouled. As a result the diffusion current continually decreases, especially near and after the end point. This effect is not serious if the titration is carried out rapidly. With more than trace amounts of phosphate the titration curve reaches a poorly defined maximum before the true end point, and it is not possible to effect a reliable titration. The limiting amount of phosphate which can be tolerated in a fluoride determination is about 1 mg. per liter (Table 11). Miscellaneous. For a given concentration of fluoride, the change in current during a titration becomes greater with increased concentration of the ferric iron indicator. For high sensitivity it is therefore desirable t o use enough indicator solution so that the galvanometer will be almost fully deflected a t the end point. The range of analysis might be extended to higher fluoride concentrations by decreasing the sensitivity of the galvanometer and increasing the concentration of the iron. On the other hand satisfactory conditions for extension of the range t o lower fluoride concentrations by increasing galvanometer sensitivity and decreasing the iron concentration could not be found. With less than about 200 y of fluoride the experimental points for the titration curve lie on two straight lines. With larger amounts the experimental
points deviate from linearity near the beginning and near the end point but without seriously decreasing the accuracy of the end point determination.
Figure 1. Amperometric fitration apparatus
The results of the titrations of varying amounts of fluoride show that the standard curve is linear over the range from 50 to 600 y per 100 ml. of solution. With less than about 50 y of fluoride the results indicate that slightly smaller volumes of standard thorium nitrate solution are required than would be expected from a linear standard curve. The fluoride-thorium ratio a t the end point is about 1.5, which suggests formation of ThF+++ and ThF2++ complexes during the titration. These complexes have been identified and studied (3, 4, 14). Sixteen replicate titrations of fluoride solutions containing 204 y of fluoride gave results with a relative standard deviation of 1.8%. Thirty micrograms of fluoride in 100 ml. can be determined to a precision of about 2001,. Below about 20 y per 100 ml. no end point is obtainable. LITERATURE CITED
(1) Carson, W. N., Michelson, C. E., Kovama. K.. ANAL. CHEM.27, 472 (1955). ’ Castor, C. R., Taylor, J. H., Zbid., 24, 1369 (1952). Day, R. il., Stoughton, R. W.,J . Am. Chem. SOC.72,5662 (1950). Dodgen, H. W.,Rollefson, G. K., Zbid., 71, 2600 (1949). Kolthoff, I. M.,Harris, W. E., IND. ENQ.CHEY.. ANAL. ED. 18, 161 (1946). (6) Langer, A., ANAL. CHEM. 12, 511 (1940). (7) Nordic, K., Rezac, Z., Chem. listy 49, 570 (1955).
(8) Shiner, V. J., Smith, AI. L., ANAL. CHEM.28, 2028 (1956). (9) Shoemaker, C. E., Zbzd., 27, 552 (1955). (10) Simons, J. H., ed., “Fluorine Chemistry,” Vol. 2, pp; 102-57, Academic Press, n’ew 1 ork, 1954. (11) Sundareean, RI., Karkhanavals, M.
D., Current Sci. (India) 23, 258 (1954). (12) Willard, H. H., Diehl, H “Advanced Quantitative Analysis’,’” Van Nostrand, New York, 1954. (13) Willard, H. H., Winter, 0. B., IND. ESG. CHEJI., ANAL. ED. 5, 7 (1933).
(14) Zebroski, E. L., Alter, H. W., Heumann, F. K., J . Am. Chem. SOC. 73, 5646 (1951). RECEIVED for review October 14, 1957. Accepted January 4, 1958. Division of Analytical Chemistry, 132nd Meeting, ACS, New York, N. Y., September 195i.
The Functionality of Phenols by Bromination ARTHUR
K. INGBERMAN
Development laboratories, Bakelite Co., Division o f Union Carbide Corp., Bound Brook, N. 1.
b A solution of bromine in acetic acid, catalyzed by pyridine, provides a specific and quantitative bromination of the unsubstituted ortho and p a r a positions in phenol. The method is applicable to phenols and many monosubstituted phenols, 2,4,6-trisubstituted phenols, the dihydroxydiphenylmethanes, multiring phenols, bis(0-hydroxybenzyl)amine, phenols in the presence of hexamethylenetetramine, ethers, and hydroxymethyl groups.
P
Sovolak resins are generally regarded as linear (non-crosslinked) chains of phenolic nuclei, joined by methylene links with side chains. and very few, if any, methyl01 groups. The cure of such resins with additional formaldehyde or hexamethylenetetramine is probably analogous to a vulcanization process-i.e., the establishment of additional cross links as methylenes, or meth>-lenimine linkages, respectively. between adjacent chains that have unsubstituted reactive positions in the phenolic rings ortho or para to the phenolic hydroxyl groups. A method of determining the number of unsubstituted positions in phenolic nuclei ortho or para t o the phenolic hydroxyl therefore offers not only a means for analyzing a large number of phenols, but also a tool for following the kinetics of the phenol formaldehyde cross-linking reactions. The classical bromate-bromide method of Koppeschaar ( 5 ) has been used for the estimation of phenol and unreacted phenol in phenol formaldehyde condensates. The method can be extended only to meta-alkylated phenols. Variable nonstoichiometric results were obtained with ortho- or para-substituted alkyl phenols as well as with the dihydroxydiphenylmethanes ( I , 9). The bromate-bromide method also requires the use of aqueous systems, which limits its applicability for less soluble phenols. Alkaline iodometric techniques are more selective and useful for cresols as well as HEKOLIC
dihydroxydiphenylmethanes, but are still limited by the use of aqueous systems (8, 7). The variable results obtained with the bromate-bromide method are probably due t o the high reactivity of the brominating agent, H20Br*, which has a tendency to brominate side chains as well as nuclear positions. Experience in this laboratory with a milder reagent (iodine monochloride in acetic acid) was not too promising. Bromination with molecular bromine in acetic acid proved to be a somewhat more facile reaction, and phenol was completely brominated within 45 minutes a t 80” C. to yield 2,4,6-tribromophenol. These conditions were still somewhat too rigorous for the more reactive compounds-i.e., saligenin and p-phenylphenol. The use of pyridine as a catalyst, however, provided specific and stoichiometric bromination of the ortho and para positions of a wide variety of phenols, generally complete within a fen- minutes a t room temperature. It is probable that the active brominating agent is pyridinium perbromide. EXPERIMENTAL DETAILS
An amount of sample that consumed about three milliatoms of bromine was weighed into a 250-ml. iodine flask and treated with exactly 25.00 ml. of a 0.15M solution of bromine in acetic acid. After complete dispersion, a solution of pyridine in acetic acid mas added, the flask was stoppered, and the contents were permitted to stand at room temperature for a minimum of 2 minutes, but preferably not longer than 20 minutes. About 1 ml. of a 27y0 solution of pyridine in acetic acid was found effective. A minimum of 5 ml. of 50% aqueous potassium iodide was added, and the liberated iodine was titrated in the usual manner with standardized 0.15M sodium thiosulfate, A blank was run in the same manner. The pyridine used was Barrett Grade 2A, purified first by fractionation over barium oxide, followed by treatment with 10% of its weight of liquid bromine, fractionated, and finally refrac-
tionated over barium oxide. All other solvents and reagents were standard chemically pure grades. The phenols used for this study -sere all carefully recrystallized or redistilled to constant physical properties with the exception of 2,6-bis(o-hydroxybenzyl)phenoland 3,3‘- bis(o- hydroxybenzyl) -2,2’-dihydroxydiphenylmethane. These compounds were recrystallized only once. Some gelled phenol formaldehyde condensates that were difficultly soluble in acetic acid alone were frequently soluble in a 1 to 1 mixture of acetic acid and dimethylformamide. The use of dimethylformamide did not affect the results. With unknown materials it is advisable to run a series of determinations with varying reaction times to ensure complete bromination. CONSUMPTION
OF BROMINE
The bromine consumed was determined as the difference between the total bromine employed as measured by the blank determination and the residual bromine determined iodometrically. The detection of the end point using starch as an indicator was generally useful for pure compounds. Kith polymers, the use of this indicator was complicated by the formation of colored or insoluble bromination products. In such cases, the titration was carried out potentiometrically using platinum and calomel electrodes. Glass electrodes were not usable because of their high membrane resistance. Generally, it was faster and more convenient to perform a dead-stop end point titration. By this technique, no intermediate readings were taken. The end point of the thiosulfate titration was accompanied by a sudden depolarization of the electrodes system. These electrometric determinations were made with a Leeds & Northrup 7664 pH meter using platinum foil electrodes. The results obtained by the dead-stop end point technique, the usual potentiometric titration, and the visual titration nith starch all coincided. VOL. 30, NO. 5 , MAY 1958
1003
