amounts of inorganic chlorides. Chlorides were also determined on each of the samples after filtration, and after treatment with aluminum hydroxide with subsequent filtration, to eliminate interference due t o turbidity. Results shown in Tables IV and V are a t the 95% confidence level. They show that single chloride determinations by the Mohr method yielded results within =t4 to 1 2 6 mg. per liter of the average chloride content a t concentration levels from 760 to 4759 mg. per liter. This would result in an error of hO.9 to h 5 . 9 mg. of oxygen per liter correction of the C.O.D. The precision of chloride determinations a t the 9400 mg. per liter level mas found to be less. A single determination would be viithin 1 1 4 8 mg. per liter from average, yielding an error of 1 3 3 mg. of oxygen per liter correction of the C.O.D. REFLUX TIME REQUIREMENTS
The reflux time required to obtain maximum oxidation was determined by studying limitations of the modified
method for C.O.D. It has been established (Table 111) that a t least a 1-hour reflux period before catalyst addition is necessary to oxidize inorganic chlorides completely in concentrations up to 12,000 mg. per liter. Several standard samples, to which varying amounts of inorganic chlorides had been added, were analyzed in duplicate by the modified procedure. The total reflux time was varied from 1.5 to 4.0 hours with the reflux time prior to catalyst addition kept a t 1.0 hour for each determination. The results as shown in Table VI indicate that the oxidation obtained within a total reflux time of 2 hours is near the maximum level and that there is little gain in extending the time beyond 1.5 hours for most of the compounds tested.
ferences due to inorganic chlorides. This method should be of value in EStimating the strength of wastes, especially those resulting from petrochemical industries, whose process and cycle waters may contain appreciable amounts of chlorides. LITERATURE CITED
(1) American Public Health Association, New York, “Standard Methods for the Examination of Water, Sen-age and Industrial Wastes,” 10th ed., p. 60,
1955.
12) Ibid.. D. 260. (3j Ibid.; ‘p. 333.
CONCLUSION
(4) Moore, IT,-4., Kroner, R. C., Ruchhoft, C. C., ANAL. C m M . 21, 953 (1949). (5) Moore, W. A., Ludzach, F. J., Ruchhoft, C. C., Ibid., 23, 1297 (1951). (6) Pierce, W. C., Haenisch, E. L., “Quantitative Analysis,” 2nd ed., p. 258, Wiley, New York, 1937.
This modification of the standard dichromate reflux method for determining C.O.D. of industrial wastes offers all the advantages of the standard method plus a n elimination of inter-
RECEIVED for review September 9, 1957. Accepted April 28, 1958. Divisions of Industrial and Ennineerine Chemistrv and Water, Sewage, aGd Sanit&on Chemistry, 132nd Meeting, ACS, New York, N. Y., September 1957.
Cellulose Supported Thorium-Alizarin Red S Reagent for Fluoride Ion Determination STANLEY K. YASUDAl and JACK L. LAMBERT Department of Chemistry, Kansas State College, Manhattan, Kan.
b The reagent and procedure described are the result of a study of colorimetric methods to produce a color in solution directly proportional to fluoride ion concentration. Fluoride ion in the concentration range to 15 p.p.m. reacts selectively with the reagent to release the thorium-Alizarin Red S chelate (absorption maximum, 520 mp) into solution. Concentration limits for a number of possible common interfering ions were determined. A polymeric structure for the reagent and a mechanism for the ion exchange reaction are proposed. The combining ratio of thorium to dye was found to be 1 to 2 by gravimetric and spectrophotometric methods. Chelation of thorium through a carbonyl oxygen and the 1-hydroxy group in the Alizarin Red S was substantiated by infrared absorption data. The method should be convenient for rapid visual or spectrophotometric determination of fluoride ion within the conditions specified. Present address, Los Alamos Scientific Laboratories, Los Alamos, N. M.
T
colored compounds obtained in solution or suspension from the reaction of thorium or zirconium salts with hydroxyanthraquinone dyes have been described as reagents for fluoride ion determination (2, 4, 8-15). These methods generally involve a change in hue of the sample solution as a measure of fluoride ion concentration. I n an investigation of metal-dye compounds as analytical reagents ( l e ) , thorium-Alizarin Red S compound supported on filter paper was found to undergo rapid and selective exchange with fluoride ion to release anions into solution. Preliminary work (6) had indicated the potential usefulness of thorium-Amaranth compound supported on filter paper as a reagent for fluoride ion, and this Tvork extended the study to include a number of dyes of various types. HE
SPECIAL REAGENTS AND EQUIPMENT
Thorium nitrate tetrahydrate, reagent grade, 1% solution buffered to p H 1.8 to 1.9 with 0.2M hydrochloric acid0.2M potassium chloride buffer solution.
Alizarin Red S, indicator grade, filtered 1% solution. Standard fluoride ion solution, 100 p.p:m., 0.221 gram of reagent grade sodium fluoride per liter of solution. Amberlite IR-120 (H) nuclear sulfonic cation exchange resin, analytical grade. PROCEDURE
Prepare the thorium-Alizarin Rcd S reagent paper by treating 1.5-inch squares of Whatman No. 42 filter paper individually with 1% Alizarin Red S solution. Blot between filter paper to remove excess dye solution, and dry a t room temperature. Store in a brown glass container. Prepare each reagent paper immediately before use by immersing in 1% thorium nitrate solution buffered to p H 1.8 to 1.9. Remove excess thorium ion solution by washing with several changes of distilled water and blot free of water before use. Convert a 6-inch column of Amberlite I R 1 2 0 in a 250-ml. buret (inside diameter approximately 3.1 cm.) to the acid form by treatment with hydrochloric acid solution, and rinse with distilled water until the eluate no longer tests acid. Pass 75 ml. of sample VOL. 30, NO. 9, SEPTEMBER 1958
0
1485
Table I.
Effects of Interferences
Concn., P.P.M.a Substance
Ab 500
500 500
Bc 500
500
500
500 200d
500 500 100
150d 400d 2d
500
1
150d
Id
500
500 500 500 3
500d 500 1v
15
500d
was an interference not easily removed, the reaction time was chosen so that the interference from 500 p.p.m. of sulfate ion was negligible, Results of reaction time studies are shown in Figure 1. If sulfate ion is known to be absent, the method can be made more sensitive by allowing a longer reaction time between the reagent and the sample solution. aso
O /'
a25
Vaximum concentration investigated, 500 p.p.m, Maximum concentration which did not produre a color greater than that obtained with distilled water. c Maximum concentration not affecting the color produced by 2 and 10 p.p.m. fluoride ion. d Permisvible levels after treatment of the sample to minimize interferences. a
*
solution through the column a t a flow rate of 60 drops per minute. With the water level just above the top of the column, close the stopcock and add 150 ml. of sample solution. Allow the solution to flow out a t 60 drops per minute until the meniscus is a t the 30nil. mark. Collect the next 100 ml. of sample eluate in a 150-ml. beaker and adjust the p H to 3.5 with 5N nitric acid. Boil the solution gently for 5 to 6 minutes with glass beads to prevent bumping. Add 2.5 ml. of 1% silver nitrate solution and boil another 2 to 3 minutes. Cool to room temperature and filter through an ultrafine fritted borosilicate glass filter if turbid. Adjust the p H to 7.0 with approximately 0.5N sodium hydroxide solution. Transfer to a 100-ml. volumetric flask and dilute to volume. Pipet a 25.0-ml. sample into a 90-mm. porcelain evaporating dish. The method may be simplified for sample solutions known t o be free from interfering cations or anions. Place a freshly prepared thoriumAlizarin Red S reagent paper in the sample solution and swirl gently for 5 minutes a t a constant rate. Measure the absorbance of the solution a t 520 mp with the spectrophotometer, and determine the fluoride ion concentration from the calibration curve. A Beckman DU spectrophotometer with 10-mm. Corex cells was used to obtain the calibration curve. DISCUSSION
The thorium-dye reagent supported on Whatman No. 42 filter papers is insoluble, stable, and capable of rapid, selective ion exchange with fluoride ion. The absorbance of the substance released into solution was found to be directly proportional to the fluoride ion concentration in the range from 0 to 15 p.p.m. Four determinations were made a t each concentration. 1486 *
ANALYTICAL CHEMISTRY
Investigation of the supporting media indicated TThatman No. 42 filter paper to be superior to other grades from the standpoint of uniform chelate formation, and the rate with n-hich thoriumdye reagent undergoes exchange with fluoride ion. Apparently Khatman Nos. 40 and 41, because of higher porosity, caused uneven formation of the compound, resulting in mottled light and dark areas on its surface. TThatnian Xo. 50 filter paper produced uniform compound distribution, but the reagent paper reacted very s l o ~ l y nith fluoride ion. The compound supported on starch granules exhibited rapid exchange due presumably to the exposure of a large surface area to the sample solution, but this reagent was not so satisfactory as that supported on filter paper. Although the method of applying the dye on the papers was not quantitative, the gravimetric data shomed the amount of dye per gram of cellulose to be identical for different paper samples. The specified order of application of dye and thorium to the paper was necessary to obtain uniform compound formation. The dye must be added first and the paper dried, then the buffered thorium ion solution applied. The dyed papers were found to be stable for a t least 2 months when stored in a b r o w glass bottle. Photodecomposition of the dye presumably was kept a t a minimum. Reproducible results were obtained when the p H of the buffered thorium nitrate solution was between 1.8 and 2.2. Under these conditions, Kraus and Holmberg (6) found the degree of hydrolysis of thorium ion to be negligible, so that the cation can be represented as Th+4 (as). As sulfate ion in high concentration
Figure 1, Reaction time studies 1.
Blank solution
2.
400 p.p.m. sulfate ioti
solution
3.
500 p . p m sulfate ion
solution
4.
2.0 p.p.m. fluoride ion
solution
5.
5.0 p.p.rn. fluoride ion
solution
Table I summarizes the results. Most of the divalent cations which formed insoluble fluoride salts showed negative interference by removing fluoride ion. Phosphate and bicarbonate ions formed insoluble thorium salts or stable complex ions, and caused the release of the dye molecules. Potential interfering ions commonly found in natural and domestic water supplies were investigated in the presence and absence of fluoride ion. Serious negative interferences were found with aluminum and ferric ions because of their stable complex formation n-ith fluoride ion. Most cations were removed by a column of Amberlite IR-120 in the acid form. Wthout the use of this column, many divalent cations in the presence of fluoride ion interfered above a concentration of about 25 p.p.m. The interfering anions, bicarbonate and phosphate, were removed by boiling the acid solution, or by precipitation with silver ion. Comparison of the interferences of the various ions mith and without the presence of fluoride ion, showed marked difference as to the maximum ion concentration that could be present. The standardization curve of the exchange reaction of the reagent paper with fluoride ion was linear in the con-
centration range 0 to 15 p.p.m. The curves determined a t three different temperatures indicated slight changes in the slope, but the differences were within experimental error. Using this standard curve, multiple analyses were made of Manhattan, Kan., city water and synthetic samples to which known amounts of fluoride ion were added. The results obtained by this method agreed with the known values of fluoride ion concentration as summarized in Table 11. An equation relating absorbance to fluoride ion concentration was derived from the standardization curve and found to be C = 126 ( A
Table II.
Sample
S o . of Detns.
Ia
4
STRUCTURE OF REAGENT
The combining ratio of thorium to Alizarin Red S TVSS determined for the pH range 1.8 t o 1.9, which was used in this procedure. Based on the method of continuous variations, 1.0 to 10.0 nil. of 1 x 10-3,1~thorium nitrate solution mere pipetted into test tubes to which sufficient illizarin Red S solution of the same concentration and p H \vas added to bring the total volume t o 10.0 ml. A confined spot test filtration apparatus (7) with a Whatman No. 50 filter paper clamped between the two 18/9 socket joints was used with freshly precipitated barium sulfate filtered directly on the filter paper. The latter was prepared by mixing 5.0 ml. each of 0.02M barium chloride and 0.02M sodium sulfate solutions. The barium sulfate mat was rinsed with distilled water and excess water adhering to the inside walls of the filtration apparatus was removed with folded filter paper. Within the filtering flask, a 2.5 x 11.5 cm. test tube was placed directly under the outlet of t h e filtration apparatus.
F - Added, 0
F - Contained, P.P.11. 0.8-0.9
2
2,8-2.9
P.P.31.
F - Found, P.P.11. 1 0 1 2 1.1 1 2 3 4 3 1 2 7
IIb
10
10.8-10.9
3
2
2.0
3
5
5.0
29 10.6 9.9 10.5 10.6 2.0 1.8 2.1
- 0.005)
where C equals concentration of fluoride ion in parts per million and A equals absorbance through 10 mm. of solution. Because visual comparison methods would hasten the analysis, standard fluoride ion solutions from 0 t o 14 p.p.m. were prepared, and three 25.0-ml. samples of each concentration were treated according to the procedure described. The resulting solutions were combined and thoroughly mixed, and transferred to 50.0-ml. Nessler tubes. These mere compared visually with previously prepared standards. Batchwise analyses were also carried out by agitating 100-ml. samples rvith four thorium-dye reagent papers and comparing the solutions visually. These results compared favorably Ivith those obtained with the combined solutions from the three 25.0-ml. samples as shoivn in Table 111. Both the spectrophotometric and visual coniparisoii methods are applicable t o this method of analysis.
Results of Spectrophotometric Analyses
4.9
4.5
5.3
*
hianhattan, Iian., city water supply, obtained from wells in the Blue River valley. Synthetic samples.
Table 111.
Visual Comparison Method
F- in Synthetic
Sample I” I1 I11
IT’
r-6
1-1 1-11
- Samples, P.P.M. Contained 1
2 5
10 2 5
Found 1 2 5
102 3
10 10 Sample- I to IV, three 25.0-ml. Samples were comhined for comparison. b Samples 1‘ to 1-11 were batchwise analyses using 100.0-ml. samples. 0
The prepared mixtures of the dye and thorium solutions were then filtered with suction through the spot test filtration apparatus. The filtrate was collected and the absorbance was measured spectrophotometrically a t 520 mp. TKO buffers, pH 1.9 and 2.2, were used in preparing the thorium nitrate solutions to determine the p H effects on the combining ratio of dye to thorium (curves 1 and 2 in Figure 2 ) . The filtered precipitate was dissolved with 10.0 ml. of dilute ammonium hydroxide and further washed with 20.0 ml. of distilled water. The absorbance of the total volume of filtrate was measured a t 520 mp (Figure 3). Gravimetric analyses of thorium and dye were performed t o supplement the data obtained by the variations method. Ten Whatman No. 42 filter papers, 1.5 inches square, were placed in a desiccator and weighed periodically until a constant weight was obtained. The papers were then treated with 1% iilizarin Red S solution, blotted, dried, and reiveighed under desiccator conditions. The gain in weight showed the amount of dye present on the filter paper. These dyed papers were reacted with 1% t8horium solution, washed, dried, and again weighed until a con-
2
0
ML
4
6
8
0
1 L ZARIN R E D S
Figure 2. Determination of thoriumdye ratio b y spectrophotometric analysis of filtrate 1 . pH 1.9 2. pH 2.2
1
0
.
2 ML. ALIZARIN RED 9
e
I
10
Figure 3. Determination of thoriumdye ratio b y spectrophotometric analysis of dissolved residue VOL. 30, NO. 9, SEPTEMBER 1958
1487
Table IV.
Results Obtained b y Gravimetric Analysis
Samples I I1 I11 1.3673 1.3698 1.3559 I.3816 1.3674 1.3790 0.0118 0.0118 0.0115 3 . 6 x 10-5 3.6 x 10-5 3 . 5 x 10-6 mole mole mole 1.3918 1.3769 Thorium-dye paper 1.3890 0.0102 0.0095 Thorium6 0.0100 1.8 X 1 . 7 x 10-5 1.8 x 10-5 mole mole mole 28.6832 28.6474 Crucible 28.8303 28.6923 28.6584 Crucible plus Tho2 28.8407 0.0101 ThOz 0.0104 0.0110 Thorium. 0.0097 0.0103 0.0095 1 . 8 x 10-6 1 . 9 x 10-5 1.7 x 10-5 mole mole mole Ratio of Th-dyed 1:2.0 1:1.96 1:2.0 G. dye/g. cellulose 8.6 X 8 . 6 x 10-2 8.5 x Average weight of 10 papers, 1.5-inch squares. * Weight of thorium determined by the gain in weight. c Gravimetric determination of thorium. d The ratio calculated is based on thorium determined gravimetrically. Weight, G. Papera Dyed paper Dye
stant weight was obtained. The amount of thorium present on the filter paper was determined from the gain in weight, and also from the weight of the thorium dioxide after ignition. The latter was determined by two methods. The reagent papers from the thorium and dye treatments were initially charred in a n-eighed porcelain crucible over a Bunsen burner and then ignited at 1000" C. in a muffle furnace for 1 or 2 hours until a constant weight v a s obtained. The other method involved the initial digestion of the thorium-dye reagent papers with a minimum amount of a 2 to 3 mixture of concentrated nitric acid and 70% perchloric acid at 240" C. This acid treatment dissolved the cellulose completely as well as the dye of the chelate, and converted the thorium to salt. The solution was evaporated to dryness and the residue was redissolved in 10 ml. of dilute sulfuric acid with heating. The solution was transferred into weighed crucibles, evaporated to dryness, and ignited a t 1000" C. From these data, correlation of the amount of dye and thorium present in the chelate, their combining ratio, the average amount of dye per gram of cellulose, and the reproducibility of absorbing the dye on the papers \$ere established (Table IT). Samples I and I1 were analyzed by the initial charring of the thorium-dye reagent, and samples I1 and IT' by mixed acid digestion. Infrared spectra of the Alizarin Red S dye anion and its thorium chelate were studied to determine the functional groups of the dye anions involved in the compound formation. Approxi-
1488
ANALYTICAL CHEMISTRY
IV 1.3990 1.4104 0.0114 3.5 x 10-6 mole 1.4212 0.0108 2.0 x 10-5 mole 28.8931 28.9051 0.0120 0.0112 2.0 x 10-6 mole 1:1.79 8 . 2 x 10-3
mately equal amounts of the dye and the chelate samples were ground and thoroughly mixed with 0.5 gram of potassium bromide. The mixture was placed in a pellet mold, evacuated with a vacuum pump, and the pressure applied to form the pellet. Both the sodium chloride and the lithium fluoride prisms were used for the spectrum range 2 to 12 microns and 2 to 4 microns, respectively. Three samples-a blank, the dye, and the chelate-were examined (Figures 4 and 5 ) .
The method of continuous variations and the gravimetric analyses showed the combining ratio between thorium ion and Alizarin Red S dye anion to be 1 to 2. As the dye molecule has three types of functional groups, this n-ould suggest three possible structures for the compound, with bonding between carbonyl oxygen and the adjacent hydroxyl group, or between the two adjacent hydroxyl groups, or possibly through the sulfonate group. Because the 2hydroxyanthraquinone-3-sulfonic acid, the 2-hydro~yanthraquinone~and the 2-anthraquinone sulfonic acid do not form insoluble compounds with thorium, this would indicate that any such compound formation v ith thorium and Alizarin Red S dye lvould most likely take place through a carbonyl oxygen and the adjacent hydroxyl group. The infrared absorption spectrum of illizarin Red S dye shon.ed characteristic bands for the hydroxyl, carbonyl, and sulfonate groups present in the dye molecule. Hydroxy1 bands a t 3342 and 3080 cm.-l wave numbers were assigned to the free hydroxyl group not adjacent to the carbonyl, and the intramolecular hydrogen-bonded hydroxyl group, respectively. The broad shape of the 3080-cm.-' band indicated hydrogen bonding to be the case. The carbonyl bands a t 1661 and 1630 cm.-', which were not of equal intensity, gave further evidence that the first band mas hydrogen bonded, and the second band a t 1630 em.-' mas that of free carbonyl. The assignment of these hydroxyl and carbonyl bands agreed with the hydroxyanthraquinone
b Figure 4. Infrared spectra using sodium chloride prism 1. 2.
Alizarin Red S Thorium-Alizorin S compound
c
Red 2
4
6
8
10
12
WAVE L E N 6 T H . p
c
2.5
3.0
3.5
WbVE LENGTH, ,u
Figure 5. Infrared spectra using lithium fluoride prism 1.
2. 3.
Alizarin Red S Thorium-Alizarin Red S compound Same os 2 with more compound in pellet
spectra data given by Flett (3). The sulfonate bands according t o Bellamy (1) occurred at 1250, 1198, 1163, 1064, 1043, and 676 cm-l. These bands correspond well with the dye spectrum shown in Figure 4. A comparison of the thorium-iilizarin Red S chelate spectrum with that of the dye anion indicated that bands due to the sulfonate group were symmetrical and a t identical positions with relatively small changes in the intensity of the over-all spectrum. These changes are probably attributable to the electrostatic nature of the bonding betiveen the thorium and the sulfonate groups.
The carbonyl band at 1630 em-’. showed a slight shift but no indication of bonding. The most pronounced changes occurred a t 3080 and 1661 em.-’, which were assigned to the hydrogen-bonded hydroxyl and the adjacent carbonyl bands, respectively. The 3080-cm.-l band completely disappeared while the 3342-cm. band became broad and unsymmetrical in shape. Very vceak intensity with a slight shift of 1 6 6 1 - ~ m . - ~carbonyl band was evident. These changes indicated that these groups were involved in chelation with thorium ion. The spectra of related dye, 1,2-dihydroxyanthraquinone (alizarin) and its thorium chelate n-ere taken, and the changes observed were consistent with those of the Alizarin Red S dye and its thorium chelate. Similar studies of 2-hydroxyanthraquinone-3-sulfonic acid indicated no chelate formation. The absence of an hydroxyl group adjacent to the carbonyl would explain this. The observed spectra data of these two dyes further substantiated that chelation in the thorium-Alizarin Red S compound takes place between carbonyl oxygen and the adjacent hydroxyl. The presence of a new band in the vicinity of 1497 cm.-’ was found with the thorium compounds of Alizarin and Alizarin Red S but was absent in the free dyes. Such a band was not observed with the thorium-2-hydroxyanthraquinone-3-sulfonate compound. An attempt was made to identify this new band by preparing the analogous beryllium-Alizarin Red S compound. The lower mass of the beryllium ion conipared with the mass of the thorium ion should result in a shift to higher frequencies of any band dependent on the vibration of the metal in a chelate ring. The spectrum of the beryllium compound showed a comparable band at 1518 cm.-l with a shift of approximately 21 em.-’ from the corresponding band in the thorium-Alizarin Red S chelate. Although this change was small in magnitude, the shift to higher frequency indicated that this band probably resulted from the vibration of the nietal in the chelate ring which was formed. Structure 1 is probably the predominant species formed between thorium and Alizarin Red S a t pH 1.9. As the dye molecule has three types of functional groups, two other monomeric structures for the compound formation are possible with bonding betreen the two adjacent hydroxyl groups, and
1
0
0
o,+t
0
T‘ h
o? ‘0
I/
0
Structure 1
possibly through the sulfonate groups. However, the latter two structures are considered less favorable than structure 1 because they would be ions of finite size and hence should be more soluble than the thorium-Alizarin Red S chelate is observed to be. They are also ruled out from the infrared data and chemical properties of related dyes with thorium ion. Structure 1 may form an infinite, three-dimensional polymer by ionic coordination of two sulfonate groups from two other dye anions.
on curve 1, Figure 2, and the apparent molar absorptivity of the compound in solution. The latter was calculated from the absorbance of a solution consisting of 1.0 ml. of 1 X lO-3-M Alizarin Red S dye and 99.0 ml. of 1 X lO-3M thorium nitrate solution at pH 1.9. KO precipitate was formed in this solution. From the absorbance of this solution a t 520 mp, which was found to be 0.036, the molar absorptivity was calculated from the equation E
AC-IL-1
(1)
where E equals molar absorptivity, A equals absorbance, C equals molar concentration, and L equals length of the light path in cm. A value of 7200 \vas calculated for the molar absorptivity. The absorbance a t point S \vas 0.045 and from Equation 1, the solubility was found to be 6.2 x 1OP6M at normal room temperature (28’ C.). ACKNOWLEDGMENT
The authors acknowledge the help of Basil Curnutte, Department of Physics, Kansas State College, in the interpretation of the infrared absorption spectra. LITERATURE CITED
(1) Bellamy, L. J., “The Infra-red Spectra of Complex Molecules,” p. 301, Wiley,
Structure 2
The resulting polymer, Structure 2, mould be a mixed chelate and salt compound of very high molecular weight, which could be considered trapped in the interstices of the filter paper fibers, The electrically neutral polymer would be subject to attack by fluoride ions only a t or near the outer surface of the polymer particle. The point of ion exchange is favored a t the thorium-sulfonate group ionic bond, resulting in the release of a chelate fragment. This is substantiated by the color of the released soluble substance which is the reddish violet of the chelated compound rather than the yellow color of the free dye. The diffusion rate of the released fragment out into the solution is the limiting reaction step, rather than the essentially ionic displacement reaction of the sulfonate group by the fluoride ion. The solubility of the unsupported, monomeric thorium-Alizarin Red S compound was determined from the absorbance a t the stoichiometric point, S,
Ken- York, 1954. ( 2 ) de Boer, J. H., Basart, J., Z. anorg. u. allgem. Chem. 152, 213-20 (1926). (3) Flett, M. St. C., J . Chem. SOC.1948, Part 11, 1443. (4) Icken, J. M., Black, B. bl., ANAL. CHERI.25, 1741-2 (1953). ( 5 ) Kraus, K. il., Holmberg, R. W., J . Phys. Chem. 58, 325-30 (1954). (6) Lambert, J. L., AXAL. CHEM. 26, 5.58-9 119.54’1 \ - -
- I .
(7) -Lambert, J. L., Xoore, T. E., Arthur, P., Ibid., 23, 1193 (1951). (8) Lothe, J. J., Ibid., 28, 949-53 (1956). (9) Sanchis. J. RI.. ISD. ENG. CHEhl.. ANAL. ED. 6, 134 (1934). (10) Scott, R. D., J . Am. Water W o r k s ASOC. 33, 2018-20 (1941). (11) Shvedov, V. P., Lab. Prakt. (U.S. S.R.), 2-3, 22-5 (1939). (12) Talvitie, N. A., ISD. ENG. CHERI., A N A L . ED. 15, 620 (1943). (13) Willard, H. H., Winter, 0. B.. Zbid., 5, 7-10 (1933). (14) Yasuda, S. K., Ph.D. dissertation, Kansas State College, 1957. ~
RECEIVEDfor review August 12, 1957. Accepted May 12, 1958. Supported by grant D-207 of the National Institutes of Health, U. S. Public Health Service. Portion of a dissertation presented as partial fulfillment of the requirement for the degree of doctor of philosophy in chemistry at Kansas State College.
VOL. 30, NO. 9, SEPTEMBER 1958
1489
