13.5 13.5 13.0 13.4 13 9 13 8 (17 29) (14 87) (18 26)
mz
36 35 37 36
37 37 85 43 26
b 0.836 0.841 0.850 0.835 0 841 0 837 0 778 0 838 0 929
218 219 227
Presumably the third moments of the distributions could be computed also. This should give a measure of skewness. The authors have not investigated any moments of higher order than the second. One could expect less and less precision the higher the moments.
216
232 228 384 (384 2 ) 264 (266 2) 359 (360 6)
contribution to < p * > and m2 is too small to warrant consideration in detail. Much time is saved by discontinuing the measurements after 98% of the material has settled. Plot these eight values on a sheet of clear cellulose acetate laid over the standard curves and by inspection estimate a value of b by comparisons with the standard curves. Compute < p > and m? from the relations above. Table I summarizes the results of six measurements of a powder of unknown distribution and on the three standard curves of Figure 2. The parenthetical values of < p > and were derived from the detailed distributions evaluated by the tangential method a t 2-micron intervals. The unit of length is the micron.
LITERATURE
CITED
(1) Carbide & Carbon Chemicals Corp., "Synthetic Organic Chemicals," 12th ed. ( 2 ) Dotts, W.&I,,ISD. EKG.CHEM.,A x . ~ LED., . 18, 326 (1946). (3) Duncombe, C. G., a n d Withrow, J. R., J . Phvs. Cliem., 36, 31 (1932). (4) Goodhue, L. D., a n d Smith, C . Ll,,IND. EXG.CHEM.,A N A LED., 8, 469 (1936). (5) Kelly, IT. J., I n d . Eng. Chem., 16, 928 (1924). (6) Kottler, F., J . F r a n k l i n Inst.. 250, 339, 419 (1950); 251, 499, 617 (1951); J . Phys. Chem.. 56, 442 (1952). (7) L a m b e r t , R. H., ASAL. CHEX, 19, 283 (1947). (8) Oden, Sven, Kolloid-Z., 18, 5 (1916); 26, 5 (1920). (9) Sane, S.C., et al., A N A LCHEH.,22, 617 (1950). (10) Schweyer, H. E., a n d Work, L. T.. Am. SOC.Testing Materials, Symposium on S e w Methods for Particle Size Determination in the Subsieve Range, March 4 , 1941. (11) Sumner, C. G., .Isa~. CHEM.,19, 939 (1947). (12) Sumner, C. G., T r a n s . Faraday Soc., 28, 20 (1932). (13) Wiegner, 0. A , Landzcirtsch. Vers. Sta., 91, 41 (1918). RECEIVED for review May 2 , 19,52. Accepted January 9, 1953.
Titration of Aprotic Acids in Thionyl Chloride ESTHER B. GARBER, LEON E. D. PEASE, J R . , AKD W. F. LLDER .\'ortheustern University, Boston, Muss. Although some qualitative titrations to show how aprotic acids affect indicators have been performed, \cry few quantitative titrations with aprotic acids have been reported. As part of a program for the investigation of the typical properties of such acids, a search for examples of their titration against familiar bases was undertaken in this laboratory. Stannic chloride, aluminum chloride, and aluminum bromide were successfully titrated in thionj 1 chloride against pjridine with crystal violet and malachite green as indicators. Aluminum bromide was also titrated against quinoline with crystal violet as the indicator. Stannic chloride was titrated conductometrically against benzophenone. If such titrations can be performed in a variety of solvents, the procedures should w-iden the scope of the techniques available to the analytical chemist.
I
K A S earlier paper ( d ) , the typical behavior ( 3 ) of indicators
toward aprotic acids-such as ferric chloride-in benzene, chlorobenzene, and 1,2-dichloroethane was reported. Another paper ( 2 ) , described the typical catalytic action ( 3 ) of the acids aluminurn chloride, stannic chloride, and ferric chloride on the reaction of metals with thionyl chloride. This paper presents another typical property of acids and bases in thionyl chloride: neutralization, as observed by a conductance titration and by several indicator titrations. The acids aluminum chloride, stannic chloride, and aluminum hromide have been titrated against pyridine xvith both crystal violet and malachite green as indicators. Aluminum bromide was also titrated against quinoline with crystal violet as the indicator. Because so many of the usual bases are sulfonated, chlorinated, or polymerized by thionyl chloride, it was hoped that 1)enzophenone might be a strong enough base in thionyl chloride
to be used as a standard. However, benzophenone proved to be too weak to be titrated by means of an indicator. Therefore, a conductance titration n s tried, successfully. INDICATOR T I T R 4 T I O N S
Thionyl chloride and stannic chloride were distilled in all-glass stills with technique described previously ( 1 , 2 ) . Aluminum bromide was distilled three times, the last time just before using. Aluminum chloride was resublimed. Pyridine (Eastman white label) was dried over calcium hydride and distilled through a Vigreux column just before using: only a small middle fraction was taken. Quinoline was treated in the same way, except that an ordinary still was employed because of the high boiling point of quinoline. Aluminum chloride, aluminum bromide, and stannic chloride are extremely hygroscopic. However, in an attempt to make the technique as simple as possible, sealed ampoules were not used for weighing the solutes. Instead, a dry box was employed for transferring them from the receivers to weighing bottles. The solutions were made u p in 250-ml. volumetric flasks to known molarities. They were titrated from 50-ml. Gpissler burets lubricated with silicone stopcock grease. The basic solution was added to the acidic solution in a 250-ml. Erlenmeyer flask. Although many more runs were made, the results presented in Table I are typical of those obtained after the preliminary difficulties were ovwcome. Crystal violet and malachite green were chosen as indicators because the solvent is so acidic. In fact, thionyl chloride is such a strong acid that the color of crystal violet alone in it is close to its intermediate shade of yellowish green. However, both indicators are unstable in thionyl chloride. Because of this instability, a better color change was obtained when the indicator was added in solid form just before the end point was reached. First, a preliminary run was made to determine an approximate end point. Then in the subsequent titrations as the end point was approached the indicator was added by dipping a glass stirridg rod into the bottle and tapping off the few adhering crystals. Care was taken not t o add too much indicator: better results were
ANALYTICAL CHEMISTRY
582 obtained when only enough crystals were added to give the acidic solution a pale yellow color. For some titrations one indication seems to be better than the other; this is indicated in the table by naming the indicator used. When pyridine solution was added to stannic chloride solution a white precipitate, which did not interfere with the end point, was formed, However, when an attempt to perform a similar titration using quinoline as the base instead of pyridine was made, the amount of precipitate was too large to permit a good end point.
Table I.
Acid-Base Titrations Using Crystal Violet and Malachite Green in Thionyl Chloride Volume A , M1.
Volume B , M1.
Moles A
c5Hax
25.94 26.01 26.05
29.17 29.29 29.29
2.002 2.004 2.001
0.09821 M CsHaS
33.09 33.06 33.01
37.32 37.35 37.26
2.007
Base, B
Acid, .4 Crystal violet 0.05517 M SnClr
0.09821
Malachite green 0.05517 M SnC1,
Also, independent measurements of the change in resistance with time were made using a k e d concentration of either the acid or the base alone in the solvent. In the acid solutions, the specific conductance increased about 100% over a period of several hours before coming to a constant value. Apparently, a slow reaction occurs between the acid and the solvent. However, in all the runs the same type of sharp break a t 2 to 1 ratio of base to acid shown in Figure 1 was observed. Also, as the end point was approached, the resistances became constant. These facts indicate that the reaction between the acid and the base takes precedence over the reaction of the acid with the solvent. Consequently, when readings are taken a t similar time intervals, rapid runs can be made, from which smooth curves, all showing the same sharp break a t the expected 2 to 1 ratio, are obtained.
2.011 2.009
Crystal violet 0.08994 M AlBn
0.09299 M CsHaN
30.00 30.00 30.00
28.82 28.80 28.84
1.004 1.005 1.004
Malachite green 0.1144 M AlCli
0.08599 M CsHaN
26.35 24.99 25.10
34.13 33.72 33.97
1.008 1.010 1.013
Crystal violet 0.06404 .M AlBn
0.07207 JI C9H:S
30.00 30.00 30.00
26.32 26.34 26.20
1.010 1.009 1.015
After a number of titrations were carried out with purified thionyl chloride, several runs were tried using thionyl chloride directly from a glass-stoppered reagent bottle in which it had been stored for several weeks without purification. First, a blank was run on the same volume of solvent that would be used in a titration, usually 55 to 60 ml. About 0.06 to 0.08 ml. of the acid solution were required to give the same end point color as was used in the titrations. When allowance was made for this in the calculations, the same results were obtained as with purified solvent. Apparently, purification of the thionyl chloride is unnecessary.
CONDUCTANCE TITRATION
When an attempt to titrate benzophenone according to the technique just described was unsuccessful, a conductance titration was tried, Thion 1 chloride and stannic chloride were distilled as described previousyy (1,2) (against a slow stream of dry nitrogen) into the conductance cell of the Erlenmeyer type, and weighed by difference in the cell. The specific conductance of the solvent was usually about 1 X 10-7 mho. Benzophenone, dried overnight in a vacuum oven, was added in small increments from a long-tailed weighing tube. The conductance cell was in a thermostat kept at 25 O f 0.05 O C. : nitrogen was flowing slowly out the top whenever the cell was o en. Figure 1 gives tge results of two typical rune. Because the resistances of the acid side slowly changed with time, more than 20 runs were made over a range of initial stannic chloride concentration from 0.01 to 1.1molal. Some runs were made as rapidly as possible; others were made over a period of 10 to 12 hours.
MOLES BASE
3
2
I
All the results are high in the amount of base required to react with I mole of acid. The explanation may lie in the choice of an end point, which was taken as red-violet. Both indicators have rather wide transformation intervals. Considerable practice was required to recognize and duplicate the same shade. Whatever the explanation, the fact that the results are all high means that if they had been calculated in terms of one of the substances chosen as a standard the results would have looked better. This wm not done because none of the reagents is of the type ordinarily considered suitable for a standard.
/ MOLES
ACID
Figure 1. ConductanceTitrations of Stannic Chloride and Benzophenone at 25' C. 1. 2.
1.885 millimoles of SnCh a t initial concentration of 0.0207 M 6.77 millimoles of SnClr a t initial concentration of 0.0455 M
Examination of the curves in Figure 1 discloses that benzophenone is a very weak base in thionyl chloride. Also relative to the solvent, stannic chloride is not a very strong acid. The reason is that thionyl chloride is nearly as strong an acid, as indicated by its action upon metals (2) and by the color of crystal violet in it as explained in the first part of this paper. Although the rapid increase in slope before the end point is hard to interpret, the rise in the curves to the left of the 2 to 1 ratio can be explained as follows. Before any base is added, the conductance of the acid solution is only slightly higher than that of the solvent. As base is added, the rise in the conductance must be caused by the replacement of the slightly ionized stannic chloride with the more highly ionized SnClr.2(C6&)&0. (The dielectric constant of thionyl chloride is high enough to permit considerable ionization of a salt that has large ions.) This change could be represented by the following equation: SnCh
+ ~(C&)ZCOe [SnC12.2(Ce&)&O] + 2c1'+
The reason for showing benzophenone rather than its ionized form in this equation is evident when the shape of the curves beyond the 2 to 1 ratio is considered. The slow rise upon further addition of benzophenone shows that even in the strongly acidic
583
V O L U M E 25, NO. 4, A P R I L 1 9 5 3 solvent benzophenone, btting such a weak h s e , ia slightlv ionized. This slight ionization might be represented hy the following equation :
+ SOClp
(CsHo)zCO
Zs
[\CaHj)iCO:SOCI]
+
+ ‘21-
In such a ztiongly acidic bolvent aq thioriyl chloride the curves are. as one would predirt on the basis of other information. tvpical of the titration of a 11 e,ik acid against a weak baie
LITERATURE CITED
IT.H., Jr., and Luder, W. F., J . Am. Chem. S O C , 66, 107 (1944). (2) Hubbard, R. A,, 2nd, and Luder, W. F., Ibid.,73, 1327 (1951). (3) Luder, W.F., and Zuffanti, S., “Electronic Theory of Acids and Bases,” New York, John Wiley &- Sons, 1946. (4) Rice, R. V.,Zuffanti, S., and Luder, IV. F., ANAL. CHEM.,24, 1022 (1952). (1) Btoniley,
RECEI!E D f o r r e r k w AIarch 28, 195’7. Accepted January 21, 1953.
Determination of Oxalate Ion and Calcium Ion by Indirect Colorimetry FERNANDO BURRIEL-JIARTi, JUAN H.4SIiREZ-MU%OZ, AND ENRIQUE FERNLNDEZ-CALDAS Analytical C h e m i s t r y Laboratories, Madrid University, a n d High Council of Scientijlc Research, Madrid, Spain
An attempt has been made to apply the decrease in optical density when the ferric-salicylate complex is treated with oxalate to determine oxalate ion by indirect colorimetry. Beer’s law is followed. The influence of the acidity and presence of tartrate and citrate ions has been studied. This method can be applied to the determination of small amounts of calcium even in the presence of magnesium, and results compare satisfactorily with those obtained by volumetric methods. Calcium is precipitated in the presence of an excess pf oxalate, and oxalate is determined colorimetrically. The method is reliable in the concentration range 2.5 X to 2.5 x 10-6 mole of oxalate. An acceptable precision is achieved in the range 1.2 to 2.2 mg. of calcium.
EVER.41, colorimetric methods have been advanced recently, in which the intensity of color of a particular system is reduced by addition of the ion being determined. Because little work has been done on the colorimetric determination of oxalate ion (11, 12), an attempt has been made t o utilize the fading in color when the ferric-salicylate complex is treated with oxalate, and a t the same time to apply this method t o the determination of small amounts of calcium. The ferric-salicylate complex has been utilized in the colori-
metric determination of iron (S, IO). llehlig ( 6 )has studied the influence of several factors, especially cations and anions, in this type of colorimetric work. Ions such as oxalate, citrate, and tartrate must be avoided in the determination of iron with salicylate (8). I n acetic acid medium, ferric iron in presence of salicylate yields complexes whose constitution depends on the p H of the medium. According to Monnier, Rusconi, and JTenger ( 6 ) and Babko ( l ) ,three complexes may be formed, violet, red, and yellow, in accordance with the theoretical and experimental findings of Bertin ( 2 ) . The p H recommended by Llehlig ( b ) is the most suitable for the present method. At a p H of about 5 an equilibrium is reached betaeen salicylate and acetate complexes. At higher p H these complexes may be destroyed by formation of colloidal ferric hydroxide. I n acetic acid medium, salicylate, oxalate, and ferric ions being present, the equilibrium depends on the amount of oxalate added to the system, since when oxalate ion increases, ferric ion decreases, as the complex OxFe+ is formed from the complex RFe“. Itoreover, from the complex OxFe+, the follon ing complexes may be formed: O.\Fr+ Ox-- c-f O.\iFcOx2FeOx-- e Ox3Fe---
+ +
SOLUTIONS AND APPARATIJS
v 50
0
I
2
3 X I O ” HOLE O X A L A T E
Figure 1.
Effect of Amount of Reagent
5
1. Iron standard solution, prepared from ferrous ammonium sulfate. 1 ml. = 0.1 mg. of ferric oxide. 2. Colored reagent. Mix 500 ml. of solution 1 with 250 ml. of 1% sodium sahylate, add ammonium hydroxide (1 to 1) dropwise until a yellow color appears, then add 10 drops of ammonium hydroxide (1 to 1) in excess. Make up to 1000 ml. Kith acetic acid (1 t o 1). 20 ml. = 1 mg. of ferric oxide. 3. Oxalic acid solution. 1 ml. = mole of oxalic acid. 4. Oxalate solution. 1 ml. = mole of sodium oxalate. 5 . Tartrate solution. 1 ml. = 2 X mole of sodium
tartrate.
