ANALYTICAL CHEMISTRY
142 that the mass spectrometric method described is suitable for this analysis, however. I t Lvas of some interest to obtain the monoisotopic mass spectra and compare the dissociation patterns of these st~ructurallysimilar compounds. Although some of these are already available, it was thought important in such a comparative study that all data lie obtained under very similar instrumental conditions. In order to obtain the monoisotopic spectra, correction was first made for carbon-13, using the observed isotope ratio C13/C‘2 = 0.0112. Next the interference from singly and doubly charged mercury ions was deducted, using the isotope abundances found experimentally, as follows, which are in excellent agreement with the data of Inghram ( 3 ) : Isotope Percent
198 10.03
199 16.77
200 23.20
201 13.20
202 29.86
204 6.85
The final operation consisted in applying statistical considerations to the chlorine isotope distribution in the various ion fragments, using the experiment,ally determined ratio CI3’/Cl3j = 0.319. The procedure is identical with that used by Dibeler and Bernstein ( 1 ) in obtaining the monoisotopic spectrum of chloroform and chloroform-d. Table T’ summarizes the anticipated pattern of the chlorine-35 and chlorine-37 isotopes, assuming a completely random distribution of the isotopes in any given species of ion fragment. Experimentally this distribution has been verified for all fragments, for every molecule except tetrachloroethane and methylene chloride, where a direct test is difficult because of superposition of C135H2 peaks upon CI37 peaks. In these cases the distribution given in Table V was assumed and utilized in computing the monoisotopic patterns. For the doubly charged ions where appreciable overlap occurred, it was necessary to solve appropriate simultaneous linear equat,ions involving the appropriate distribution ratios. Because these ions are usually present in low abundance, and often masked by singly charged ions, the relative errors are somewhat larger in these cases. Table V I lists the monoisotopic patterns for the seven compounds studied; also a quantity termed “maximum-ion sensitivit!.,” based upon the ion fragment whose relative intensity is given as 100.00. This quantity is defined as ( Z I j / P )where Zl j ~
1’
i
represent,s the sum of all the isotopic (C1) modifications of the ion prior to normalizing, and P is the pressure in microns. Figure 2 illustrates graphically the appearance of the monoisotopic spectra for each of the compounds studied. For purposes of simplicity, only singly charged ions of relative intensity >2% are indicated in the figure. Certain qualitative features are immediately apparent on in-
spection of the monoisotopic patterns. First’, as anticipated, no detectable parent peak was found for carbon tetrachloride and hexachloroethane. In a study of the mass spectra of carbon tetrafluoride, silicon tetrafluoride, and sulfur hexafluoride, Dibeler and Mohler (2) found that, dissociation by electron impact rather than formation of t,he molecular ion was by far the most probable occurrence. It is of interest, to note from Table VI that the molecular ion is the most abundant. species in the spectrum of t,etrachloroethylene, however. For carbon tetrachloride, chloroform, and methylene chloride, it is seen t’hat the most probable process is the rupture of C-CI bond by electron impact. However, for hexachloroethane and sym-tetrachloroethane, the cleavlink is most probable. I n the case of pentaage of the C-C chloroethane, both processes occur iTith equal probability. Certain regularities are also noted in the ionization of the pairs of moleculespentachloroethane-chloroformand tetrachloroethanemethylene chloride. For pentachloroethane and chloroform the relative probabilities of removal of a hydrogen atom, a chlorine atom, and a hydrogen chloride fragment, compared to molecule ion formation, are seen to be 0.28, 78, 1.1, and 0.56, 43, and 1.9, respectively, while for the pair tetrachloroethane and methylene chloride these values are 0.08, 0.97, 0.23 and 0.035, 1.3, 0.13. The relative stability of t’he molecular ion compared to the dissociated ion of largest abundance is seen to decrease in the series tetrachloroethane, pentachloroet’hane, and hexachloroethane and likewise in the series methylene chloride, chloroform, and carbon tetrachloride, as the number of electronegative chlorine at,oms is increased. This is reasonable in the light of the high electron affinity of the chlorine atom and the lessened tendency toward removal of an electron to form the molecular ion as the number of chlorine atoms is increased. ACKYOWLEDGMEYT
The authors appreciate the valuable cooperation of D. V. Kniebes with regard to the mass spectrometer. Partial support of this work by the A4tomicEnergy Commission is gratefullv acknowledged. LlTERATURE CITED (1) Dibeler, V. H., and Bernstein. R. B , J . Chem. P h y s , 19, 404 (1951).
(2) Dibeler, 1‘.H.. and Mohler, F. L., J . Research,~at2.BzLr.Standards, 40, 25 (1948). (3) Inghrarn, M., Phys. R e v , 71, 560 (1946). (4) ?ne, P. F., and Druschel, M . L., A N ~ LCHEM., . 24, 626 (1962). (5) T ogel, A . I., “Textbook of Practical Organic Chemistry,” p. 174, London, Longmans, Green and Co., 1948.
RECEIVED for reviea July 7,
1952.
Accepted October 13, 1952
Thermal Stability of Potassium Acid Phthalate EARLE R. CALEY
AND
ROBERT H. BRUNDJNl
Department of Chemistry, T h e Ohio S t a t e Liniversity, Columbus, Ohio
P
OTASSIUM acid phthalate has been widely accepted as the most satisfactory substance for standardizing base solutions, but a lack of agreement exists as to its thermal stability. The experiments of Hendrixson (1) indicate that it is stable a t temperatures up to 150’ C. but Kolthoff and Stenger (3) caution against drying above 126 O C. because loss of phthalic anhydride may occur. Hillebrand and Lundell ( 2 ) recommend 120” C., and others-Willard and Furman ( 4 ) , for example-specify 110” C. as the most suitable drying temperature. For USP in acidimetry, the Kational Bureau of Standards, in it6 certificate of analysis for potassium acid phthalate, recommends that the salt, after being 1
Present address, U. S. Military Academy, West Point, N. Y .
crushed to about 100-mesh size, be dried a t 120’ C. for 1 to 2 hours. The bureau found no significant change in weight after 2 hours a t this temperature. The purpose of the present investigation was to study the thermal stability of potassium acid phthalate in the temperature range 110” to 200” C. Because of the way it is usually dried for use as a primary standard, only its stability in an air oven a t normal atmospheric pressure was investigated. DECOI\IPOSITION PRODUCTS
I n the temperature range here considered the products of the thermal decomposition of potassium acid phthalate are water,
V O L U M E 2 5 , N O . 1, J A N U A R Y 1 9 5 3
143
The thermal decomposition of dry potassium acid phthalate first becomes noticeable around 145" C., and the rate becomes fast enough for convenient measurement around 150' C. Data are given on the initial rates of decomposition at various temperatures up to 200' C. A safe upper limit for drying potassium acid phthalate for use as a primary standard is 135' C. The products of the thermal decomposition of potassium acid phthalate up to 200" C. are water, phthalic anhydride, and normal potassium phthalate. Mixtures of potassium chloride and potassium acid phthalate used as unknown weak acid samples have about the same thermal stability as potassium acid phthalate itself.
phthalic anhydride, and normal potassium phthalate. Thc existence of these products is not easy to estahlish except at temperatures near the top of this range, for at temperatures near the middle of the range the amounts formed in a rrxasonable time are so small that they are not easily isolated and detected. Even a t these higher temperatures runs of at least 2 or 3 days by the methods here described are necessary in order to isolate these decomposition products in amounts sufficient for proper identification, The formation of water as a product was demonstrated by passing dried air over a heated sample and testing the exhaust air for moisture. The sample, previously dried to constant weight a t 110' C., was placed in a horizontal electrically heated glass tube. A slow stream of air, dried by passing it through a column of ordinary Drierite and then through a test column of indicating Drierite, was passed over the sample and exhausted through a second test column of indicating Drierite. The amount of water so indicated was greater than any possible amount of hygroscopic moisture that could have remained in the sample dried a t 110" T o demonstrate the formation of phthalic anhydride, a sample was heated in a small flask provided with a heating jacket and a cold-finger condenser. The melting point and other physical properties of the product that collected on the condenser, and the properties of its amide, showed that this sublimate was phthalic anhydride.
c.
S o volatile products other than water or phthalic anhydride could he detected, and these also appeared to be the sole volatile decomposition products a t temperat'ures up t o 250" C. When a 10-gram sample of potassium acid phthalate was heated for a week a t 250' C., and then for 3 weeks a t 200" C. a neutral residue TTas finally obtained that no longer gave off any volatile products. This residue was analyzed for potassium in the following way. Half-gram samples were first, carefully ignited in platinum over a gas burner, and then ignited in an electric muffle a t 650" C. The product of this ignition was potassium carbonate. The residues of potassium carbonate were dissolved in 10-ml. portions of water, and the solutions were neutralized by the very cautious addition of hydrochloric acid. The resulting solutions of potassium chloride were evaporated to drj-ness, and the residues were ignited to constant weight a t 550" C. and weighed as potassium chloride. Ten individual determinations gave a mean of 32.19% potassium, with a low of 32.11y0 and a high of 32.27%. These results agree satisfactorily n-ith the theoretical figure of 32.27%, though they arr generally a little loiv. However, slight mechanical loss by spraying on the neutralization of the potassium carbonate solution with the acid may account for a t least some of the lower individual determinations. On the nrhole, therefore, these results show definitely that the neutral stable residue from the complete thermal decomposition of potassium acid phthalate consists of normal potassium phthalate. I t is probable that the initial step in the decomposition of potassium acid phthalate is the formation of the normal salt and phthalic acid, but the latter apparently decomposes immediately into phthalic anhydride and water. i i o evidence could be obtained for the presence of phthalic acid in residues from the partial decomposition of potassiuni acid phthalate, a t least a t temperatures in the upper part of the range here considered. Therefore:. any residue from the partial thermal decomposition of
potassium acid phthalate should ho a mixture of this salt and the normal salt. Because a stoichiometric relationship exists between the proportion of the two salts in such a residue and t,he proportion of both water and phthalic anhydride that is lost on heating, the weight of potassium acid phthalate remaining in any such residue of known weight may be calculated from the loss of weight on heating. Furthermore, a direct estimation of the weight of potassium acid phthalate in such a residue is possible by titration, and a close agreement of the results obtained by these t'wo independent methods n-ould shorn that such a residue consisted of only a mixture of the acid salt and the normal salt. For obtaining the results shown in Table I samples of pot.assium acid phthalate dried to constant weight a t 110" C. and weighing about 1.5 grams were heated at 200" C., until the loss of weight was around 20%, the exact loss of weight being carefully measured.
Table I. Weight of Potassium .4cid Phthalate Present in Residues after Partial Thermal Decomposition T e i g h t Calcd. from Loss of Weight on Heating Grams
Weight Found by Titration Grams
0,6775 0.8354 1.1547 0.6444 0.8567 1.1662 1.1101 1.0945 0.6266 0.8109
0.6756 0.8340 1.1678 0.6445 0.8560 1.1703 1.1090 1.0938 0.6271 0.8123
Difference Mg. -1.9 -1.4 +3.1 +o. 1 -0.7 -4.1 -1.1 -0.7 +0.5 +1.4
N e a n -0
5
72 - 0 28 -0.10 +0.27 +0.02 -0 08 -0 35 - 0 10 - 0 06 +o 08 + O 17
-0
04
The residues were dissolved in hoiled distilled water and titrated with standard carbonate-free sodium hydroxide solution, using phenolphthalein as the indirator. The agreement of the individual results is as good as could he expected from the accuracy of the two analytical methods, and the mean result is very satisfactory. DRYING AYD IIEhTIhG EXPERIMENTS
These experiments were all performed on material taken fiom a single bottle of a well-known brand of reagent-grade potassium acid phthalate, which had been found from previous experiments to be satisfactory for use as a primary standard. In order to measure small percentage losses of weight accurately, 10-gram samples were used. The samples were heated in glass weighing bottles, and similar empty bottles used as tares in weighing were placed alongside the filled bottles in all the operations of heating and cooling. A 40-minute period of cooling in a large desiccator containing Drierite was uniformly used. The electric air oven was of the de Khotinsk:. type provided with controls that maintained the temperature within + l o C. of a given setting over many days of use. The mercury thermometer used in this oven was calibrated against a thermometer certified by the National Bureau of Standards. The analytical balance was sufficiently sensitive to weigh accurately to 0.1 mg., and the weights were calibrated just before the experiments were started. When samples of the salt were dried a t 110" C. for ten consecutive hourly periods, constant weight was reached a t the end ot the
144
ANALYTICAL CHEMISTRY
second hour. Closely agreeing results on four samples gave a mean loss of aeight of 0.020%. N o further loss of weight was observed a t temperatures up to 130" C. Hillebrand and Lundell (2) report that when the Kational Bureau of Standards sample of potassium acid phthalate So. 84 a a s tested, it lost 0.02% in weight when dried for 2 hours a t 120' C., and that the dried sample regained this lost weight on exposure for 8 days to an atmosphere of 90% relative humidity. It seems safe to assume, therefore, that the observed loss of weight a t 110" C. was due to the escape of hygroscopic moisture. When a sample dried to constant weight a t 110" C. was heated to temperatures in the range 130' to 140' C. a further loss of w i g h t of 0.010% was observed. Possibly this loss a t the higher temperature range was simply due to the escape of moisture tenaciously held in the crystals of the salt, as these were not ground to a powder for these drying experiments, though it seems somewhat more likely that the observed loss was due to the escape of some other volatile substance or substances, such as n-ater and phthalic anhydride arising from the decomposition of phthalic acid, which is known to be a common impurity in potassium acid phthalate. S o furthrr loss in weight was observed when the sample a a s held at 139" C. for 6 days.
Table 11. Initial Rate of Loss of Weight and Formation of Nor mal Potassium Phthalate from Potassium Acid Phthalate Heated at Various Temperatures
e 10-Gram Sample
c.
Mg./g./hr.
Mg./g./hr.
149 159 164 174 184 199
0.0023 0.0048 0.0094 0.0302 0,0927 0.335
0.0033 0.0070 0.0137 0.0441 0.135 0.488
O
Phthalate %/hr.
0.0003 0,0007 0.0014 0.0044 0.014
0,049
Thermal decomposition of potassium acid phthalate begins in the neighborhood of 145' C., though a t this temperature the progressive loss of weight on heating cannot be detected or measured quantitatively unless heating experiments are performed over very long periods of time. Hence the exact temperature a t which thermal decomposition begins is not easy to establish by direct experiment. Hon ever, extrapolation of data on the rate of loss in weight a t higher temperatures indicates that it begins slightly below 145' C. Around 150' C. the rate of decomposition becomes fast enough for convenient measurement. The rate of loss of weight on heating increases rapidly as the temperature rises, as is shown in Table 11. -4typical graph for one of the temperatures is shown in Figure 1. This graph and the data in Table I1 indicate, of course, only initial rates of decomposition. EFFECT OF POTASSIUM CHLORIDE
Mixtures of potassium acid phthalate and potassium chloride are rather commonly used as unknown ueak acid samples for students of elementary quantitative analysis. As no data on the thermal stability of such mixtures could be found, it seemed worth while to investigate this also. If the potassium chloride reacts with the potassium acid phthalate lvhen such mixtures are heated, this should be revealed by the evolution of hydrogen chloride, a rate of decomposition different from that of potassium acid phthalate itself, and a discrepancy between the amount of potassium acid phthalate calculated from loss of weight and that found by titration.
To test for the evolution of hydrogen chloride, a slow stream of urified air was passed over a heated mixture, the evhaust air was fkbbled through water, and the water was teated for chloride with silver nitrate. The sample was contained in a horizontal electrirally heated tube, the ail rvns purified by passing it through a
Figure 1.
I
I
I
2
I
I
I
3 4 5 TIME OF HEATING, DAYS
I
I
6
7
Loss in Weight on Heating Potassium Acid Phthalate at 159' C.
column of soda lime, and water free from chloride was placed in the bubbler tube. In a critical experiment in which a mixture of 2 arts of potassium acid phthalate and 1 part of potassium chlorixe was heated a t 200' C., no chloride was detected on adding silver nitrate solution acidified with nitric acid to the water in the bubbler tube after a run of 7 days. For preparing samples for heating experiments, potassium acid phthalate and potassium chloride were separately dried to constant weight a t 110' C., 10- ram samples of the phthalate were accurately weighed, about hayf as much of the chloride was added, the two salts were intimately mixed in a mortar, and the final weight of the mixed sample was determined. The initial rate of loss of weight on heating such samples to various temperatures was substantially the same as for potassium acid phthalate alone. Results similar to those shown in Table I were obtained when carefully prepared and accurately weighed mixtures containing about 1.5 grams of potassium acid phthalate and about 0.5 gram of potassium chloride were heated a t 200' C. until a loss in weight of about 20% in terms of the potassium acid phthalate had occurred, and the weight of the residual phthalate calculated from this loss in weight was compared to that found by titration. The mean difference for ten trials was +0.12%. CONCLUSIONS
Thermal decomposition of potassium acid phthalate is first detectable from direct loss in weight experiments a t about 145' C., though therate isextremelyslowatthis temperature. The extrapolation of data from heating experiments a t higher temperatures indicates that thermal decomposition actually starts slightly below 145' C. ilround 150" C. the rate of decomposition is fast enough for convenient measurement, and the rate increases exponentially with the rise in temperature. A safe upper limit for drying potassium acid phthalate for use ae a primary standard is 135" C. The products of the thermal decomposition of potassium acid phthalate up to 200" C. are water, phthalic anhydride, and normal potassium phthalate. The residue from the partial thermal d e
V O L U M E 25, NO. 1, J A N U A R Y 1 9 5 3 composition of potassium acid phthalate up to the same temperature consists of a mixture of the acid salt and the normal salt. The presence of admixed potassium chloride has no significant effect on the initial rate of decomposition of potassium acid phthalate, and mixtures of the two saks used as unknowns may be dried in the same way as potassium acid phthalate itself. LITERATURE CITED
( 1 ) Hendrixson, W. S., J . Am. Chem. Soc., 42, 726 (1920).
145 (2) Hillebrand, W. F., and Lundell, G. E. F., “Applied Inorganic Analysis.” p. 140, New York, John Wiley & Sons, 1929.
(3) Kolthoff, I. M., and Stenger, V. A., “Volumetric Analysis,” 5’01. 11,p. 94, New York, Interscience Publishers, 1947. (4) Willard, H. H., and Furman, N. H., “Elementary Quantitative Analysis,” p. 136, New York, D. Van Nostrand Co., 1940. RECEIVED for review November 2 , 1951. Accepted October 30, 1952. Constructed from a thesis presented to T h e Graduate School of T h e Ohio S t a t e University in partial fulfillment of the requirements for t h e h1.S. degree, June 1961.
Photometric Determination of Silica in Alkalies 0. A. KENYON AND H. A. BEWICK Solvay Process Division, Allied Chemical & Dye Corp., Syracuse, N. Y . Because the usual gravimetric determination of silica in alkali products is lengthy and inadequate for very small amounts, an investigation was undertaken. A photometric method employing the molybdenum blue complex was developed, which is rapid and has precision and accuracy suitable for quality control. A standard deviation of f 0 . 9 microgram in 100 micrograms of silica has been realized. No treatment of the sample other than solution, pH adjustment, and proper salt concentration is required. No serious interference is encountered. and the method has a wide range.
N
O PREVIOUS critical study of the photometric mol?-bdate method for silica in thp high salt concentrations encountered in the analysis of sodium hydroxide and other alkalies is known. A recent review (3) ha. indicated the need for a study of the effect of salinity on color development of molybdenum complexes in the determination of silica in sea n-ater. The usual gravimetric determination of silica in sodium hydroxide or other alkali products is lengthy and inadequate for very small amounts. The photometric determination of silica using the molybdate reaction is the recognized procedure for small amounts of silica, and this paper reports the development of the proper conditions for its application t o alkalies. Early investigations of this method Tvere made by Jolles and Xeurath ( 7 ) and Dieneit and Wandenbulcke ( 6 ) . The method with various modifications has been used for the determination of silica in fresh water ( 9 ) , in sea tvatcr (10, 15, 1 6 ) , and in boiler feed water ( 4 , 8,11, I S ) . The optimum conditions for the color reactions were determined with special emphasis upon the effects of pH and varying concentrations of sodium or potassium chlorides. Other important factors such as concentration of reagents, time required for color development, color stability, method of neutralizing samples, and reproducibility of the method were studied. I n this work the authors xere unable to obtain reproducible results with Bunting’s molybdate reagent ( 4 ) . Stoloff’s reagent (12)was found to be satisfactory. The volume of the sample a t the time the molybdate reagent was introduced was found to be important. Great variations in the intensity of the color formed were noted for all volumes below i 5 ml. Published data vary concerning pH recommended for maximum color formation. I n this study color formation was found to be jointly dependent upon pH and chloride concentration. The findings of this work may also be applied to the photometric determination of the much larger amounts of silica present in alkali silicates according to Hiskey’s technique (e),in which the spectrophotometer is adjusted to zero absorbancy using a high silica standard and a lyider slit width. APPARATUS
The absorbancy measurements were made with a Beckman RIodel DU spectrophotometer and 1.00-cm. cells. The red-sensi-
tive phototube was used a t 660 and 825 mp to measure the molybdenum blue complex and the blue-sensitive phototube was used a t 410 mp to measure the yellow molybdisilicic acid. The reference cell contained distilled water. Slit widths were 0.09, 0.10, and 0.15 mm., respectively. Other apparatus required included a Beckman pH meter, Model H-2, 100-ml. Kohlrausch flasks, and 100-ml. mixing cylinders. REAGEhTS
Hydrochloric acid, 10.0 Y. Citric acid solution, 10% w./v. Tartaric acid solution, 10% w./v. Sodium chloride solution, 2 M, 116.9 grams per liter. Adjusted to pH 1.3. Potassium chloride solution, 2 X , 149.1 grams per liter. Adjusted to pH 1.3. Ammonium molybdate solution, 5%. Five grams of ammonium molybdate Tvere dissolved in 80 ml. of warm distilled water. The solution mas cooled, 2.8 ml. of concentrated sulfuric acid were added, and the solution was made up to 100 ml. Ammonium hydroxide solution, silica-free. Air was passed for a t least 1 hour through 250 ml. of concentrated ammonium hydroxide in a 500-ml. gas washing bottle and conducted into 250 ml. of distilled water in a polyethylene bottle. As this ammonia solution is used only for p H adjustments, the strength is unimportant. Reductant Solution. Sinety grams of sodium bisulfite, NaHSOS, were dissolved in 800 ml. of distilled water. Seven grams of anhydrous sodium sulfite and 1.5 grams of l-amino-2naphthol-4-sulfonic acid (Eastman No. 360) were dissolved in 100 ml. of distilled water. The two solutions were mixed, diluted to 1 liter, and stored in an amber glass bottle under refrigeration. Silica Solution. A stock solution was prepared by dissolving 5.1 grams of sodium metasilicate nonahydrate, Ka2Si03.9Hz0 in 1 liter of distilled water and standardized gravimetrically ( 2 ) . Appropriate dilutions were used for the working standard silica solution in the preparation of calibration curves. EXPERIMENTAL WORK
In this investigation the concentration of reagents, stability of color complexes, method of sample neutralization, effect of pH, effect of varying concentration of sodium and potassium chlorides, and precisionof the method were carefully studied. The procedure of Bunting ( 4 ) was used as a starting basis in this method of investigation. Stability of Silica Solutions. Investigators have stated ( 4 , 9) t h a t silica solutions stored in a hard-rubber container will not deteriorate. During the present work it was found that a standard
