April 15, 1931
INDUSTRIAL A N D ENGINEERING CHEMISTRY
It is believed that this method should prove useful in the analysis of limestones and dolomites, for it is shown that the results obtained without any double precipitations or evaporation are as exact as could be obtained by making a double precipitation of the calcium by the oxalate method and then evaporating down the filtrates and washings and making a double precipitation of the magnesium by the usual method. The method herein described thus eliminates two precipitations and an evaporation without any sacrifice of the accuracy of the determinations and a t a considerable saving of time. Like other molybdates this precipitate is quite soluble in concentrated mineral acids and in strong bases. Therefore, care should be taken that the solution from which it is precipitated be nearly neutral.
129
The writer prefers to precipitate calcium from a slightly alkaline solution. No trace of molybdenum was found in the magnesium precipitates. Reducing gases should not be allowed te come in contact with the calcium molybdate during ignition. A large excess of ammonium molybdate shouId be avoided. Before filtering, a drop of the supernatant liquid should be tested with a saturated solution of pyrogallol in chloroform, A brown coloration denotes an excess of molybdate. This test is not decisive unless the solution tested has been boiled at least 10 minutes, for the precipitate of calcium molybdate forms rather slowly. Literature Cited (1) Smith and Bradbury, Bev., 24, 2930 (1891).
Conductometric Titration of Sulfate and Barium' I. M. Kolthoff and Tohru Kameda SCHOOLOF CHEMISTRY, UNIVERSITY OF MINNESOTA, MINNEAPOLIS, MI".
In the conductometric titration of alkali sulfate with it was thought necessary to T PRESENT, we have barium chloride, the knick point is found before the make a more systematic inneither an inside nor equivalence point. The error depends upon the divestigation on the reliability an outside indicator lution, alcohol content, acidity, and the kind of cations of the conductometric titraavailable for the detection of present. Especially in the titration of 0.1 Nsolutions are tion of sulfate with barium the end point in the direct or the results more or less irregular, the deviations being and the reverse procedure. reverse t i t r a t i o n of sulfate smaller in the titration of lithium sulfate than of poThe first part of the work has with a barium solution. A tassium sulfate. been carried out in the usual potentiometric method canIn more dilute solutions, the results are reproducible way-i. e., at room temperanot be used either, for the within 1 per cent. Therefore, in series of determinareason that there is no suitture, without taking precautions of sulfate, the application of the conductometric able indicator electrode for tions to keep the temperature method can be very useful if the standardization is this reaction. Therefore, the completely constant. In ormade under the same conditions as the experiments application of the conductoder to be certain that the unwith the unknown. metric method to the titrafavorable results could not be The deviation from the theoretical results cannot be tion of sulfate with barium or a t t r i b u t e d to i r r e g u l a r attributed to a coprecipitation of sulfate. By special the reverse procedure has spechanges in temperature, all experiments it has been shown that the conductivity in cial advantages. It was first the work was repeated by perthe presence of the precipitate is distinctly higher than recommended by Dutoit ( 1 ) forming the titrations in a in the filtrates. This may be caused by the charge of in 1910. Though there is an thermostat at a temperature the barium sulfate in suspension and by its influence extensive literature on this of 25" * 0.05" C. Only this upon the cell constant. titration method, there is no part of the work will be reIn the reverse titration of barium with sulfate, the general agreement as to its ported. knick point is found nearer the equivalence point a c c u r a c y . According to Apparatus and Solutions though a small deviation of the order of 1 per cent is obseveral authors, the knick Used served. The conductometric method is very useful in point in the sulfate titration the titration of very dilute solutions in the presence of C o n d u c t i v i t y cells with occurs before the equivalence vertical electrodes were used. 30 per cent alcohol. point and the results are corThey were sealed t o a heavy The mobility of the barium ions is decreased much respondingly low; a few claim more by alcohol than that of the alkali or sulfate or chloplatinum wire, the lattermakt o find accurate results. ing contact between the elecride ions. From the r e c e n t monotrode and the mercurv in the graph of Jander and Pfundt - . (f?),one gets the impression that the conductometric sulfate vertical glass tubes fused to the glass of the cell. Thi"ck coptitration under different conditions gives theoretical results. per wires dipping in the mercury made the electrical contact in However, the experimental evidence they present is too poor the circuit of the Wheatstone bridge. The biggest part of the conductivity cell dipped into the thermostat during the t o make the statement convincing. I n one of the writers' investigations, it was necessary t o titration. The microburet with the reagent was mounted make a series of analyses of the sulfate content in dilute zinc above the opening in the upper middle part of the cell. After sulfate solutions, and for this purpose the conductometric each addition of reagent, the content of the vessel was mixed method was applied. However, in testing it with solutions by careful shaking, and the final readings were made when the of known content, very disappointing results were found, the conductivity did not change further after longer standing. latter sometimes being more than 5 per cent low. Therefore In order to prevent a change in concentration by evaporation during the titration, the opening of the cell was closed by a 'Received October 20, 1930. Chapter from a thesis submitted by rubber stopper after each addition of reagent. Tohru Kameda to the Graduate School of the University of Minnesota, in All the salts used were pure substances. The potassium, partial fulfilment of the requirements for the degree of doctor of philosophy.
A
130
ANALYTICAL EDITION
sodium, and lithium sulfate were chemically pure commercial products, twice recrystallized from water, and dried in the proper way. Normal and tenth-normal solutions were prepared ip conductivity water and the strength was controlled by a gravimetric standardization according to reliable procedures. Solutions of purified barium chloride were standardized in a similar way. In part of our titrations, barium acetate was used as a precipitating reagent. Though the acetate ion has a much smaller conductivity than the chloride ion, we cannot recommend this salt for general use. Difficulties are encountered when the solutions contain some k h d of an acid. The acetate ions of the reagent form a n acetate-acetic acid buffer with the acid in bhe solution, and the conductivity changes in an irregular way on account of the variation of the hydrogen-ion concentration during the titration. The water used in all the experiments was conductivity water; the acetic acid was purified by distillation over chromic acid and the alcohol over lime. Titration of Sulfate with Barium Chloride
VOl. 3, No. 2
Table I-Titration of Sulfates w i t h Barium Chloride SULFATEAMOUNTSULFATE CONCN. TAKEN FOUND DEVIATION SOLVENT N Millimols Millimols % POTASSIUM SULFATE
0.1 0.1 0.1 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01
2.498 2.498 2.498 0.2463 0.2463 0.2463 0.2463 0.2463 0.2463 0.2463 0.2463 0.2463
2.385 2.378 2.378 0.2377 0.2399 0.2387 0.2351 0.2387 0.2327 0.2337 0.2322 0.2237
0.1 0.1 0.1 0.1 0.1 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01
2.420 2.420 2.420 2.420 2.420 0.2420 0.2420 0.2420 0.2420 0.2420 0.2420 0.2420 0.2420
2.373 2.383 2.273 2.348 2.298 0.2374 0.2369 0.2349 0.2319 0.2352 0.2312 0.2307 0.2168
-4.5 -4.8 -4.8 -3.5 -2.6 -3.1 -4.5 -3.1 -5.5 -5.1 -5.7 -9.2
SODIU[M SULFATE
- 2- . 0 .
acetic acid
acetic acid acetic acid acetic acid acetic acid
__
W.tW . . --
-1.6 -6.1 -3.0 5.1 -1.9 -2.1
alcohol iy alcohol 1 acetic acid
-4.5 -4.7 -10.4
10% alcohol 30%alcohol 50% alcohol
-
10% alcohol Water Water
+ 1% acetic acid
alcohol
+ 1'7 1% acetic acid + acetic acid + 1.9 o acetic acid
2.180 0.1 I n the following table, the results of the titrations are 0.1 2.180 reported in a condensed form. In order to save space the 0.1 2.180 0.1 2.180 changes of the conductivity during the titration and the 0.1 2.180 0.1 2.180 plotted values for finding the knick point are omitted. In 0.1 2.180 the column expressing the deviation between location of 0.1 2.180 2.180 0 . 1 equivalence and knick poiit, the minus sign means that the 0 . 0 1 0.2180 0 . 0 1 022180 knick point occurs before the equivalence point; therefore, 0.01 0.2180 too little sulfate is found. In cases where alcohol was added, 0 . 0 1 0.2180 the titrations were not begun before the temperatures of the 0 . 0 1 0.2180 solutions were the same as that of the thermostat (25" (3.). 0 . 0 1 0.2180 In the titrations of lithium sulfate, three sets of readings 0.1 0.2180 0 . 0 1 0.2180 have been made; one after three to five minutes after addi0 . 0 1 0.2180 0 . 0 1 0.2180 tion of the reagent, one after fifteen to thirty minutes' stand0.01 0.2180 ing, and the last repeated after shaking the cell, the precipi0.01 0.2180 0.01 0.2180 tate thus being suspended in the solution. 0.001 0.0215 0.001 0.0215 I n most cases the knick point is found before the equiva0.001 0.0215 lence point, the error being the largest in the titration of 0.001 0.0215 potassium sulfate (deviation about 5 per cent). The addition 0 . 0 5 1.250 of a little alcohol (10 per cent) with or without 1 per cent 0 . 0 5 1.250 0 . 0 5 1.250 acetic acid has practically no influence upon the results. 1.250 0.05 It should be remarked, however, that the shape of the precipi1.250 0.05 0 . 0 5 1.250 tation curve improves markedly by the addition of alcohol. 0 . 0 5 1.250 0 . 0 5 1.250 In titrating 0.1 N potassium sulfate solution in water the 0.05 1.250 conductivity changes during the precipitation more or less in 0.05 1.250 0.005 0.1250 an irregular way, and, therefore, the slope of the line which 0.005 0.1250 0.005 0.1250 combines the different points is more or less uncertain. For 0.005 0.1250 this reason the titration figures found are not well repro0.005 0.1250 0.005 0.1250 ducible. The same holds, though to a lesser extent, in the 0.005 0.1280 0.005 0.1250 titration of 0.1 N sodium and lithium sulfate. The titration 0.. 005 0.1250 of 0.1 N sodium sulfate yields better results than that of 0.005 0.1250 a Reading after 3 minutes. d Reading after 5 minutes. potassium sulfate, the error being about 2 per cent. Addition b Reading after 30 minutes. e Shaken. of 10 per cent alcohol had no influence; in the presence of c Reading after 15 minutes. 50 per cent alcohol, however, the results are about 6 per cent low. Acetic acid in a concentration of about 1 per cent in- The angle between the precipitation line and the reagent line creases the error from about 2 to 3 per cent, and the combi- is much more obtuse in the titrations of zinc sulfate than in nation of 10 per cent alcohol and 1 per cent acetic acid even to that of alkaline sulfates which makes the titration less precise 5 per cent. The best 'results are obtained in the titration of (compare Figure 1). It could be expected that in the titra0.1 N lithium sulfate. Even in the presence of 10 per cent tion of more dilute sulfate solutions, the difference due to the alcohol and 1 per cent acetic acid, the deviation is not larger different cations would disappear or decrease. Still more than about 1 per cent from the theoretical value. In the marked differences were found in the titration of 0.01 N titration of 0.05 N zinc sulfate the knick point is found about alkali and zinc sulfate solutions. The titration of 0.01 N 6 per cent before the equivalence point. It is peculiar that potassium sulfate yields results which are about 3.5 per cent the addition of alcohol improves the results, the error being low. The procedure is not very pleasant with water as a 5 per cent in the presence of 10 per cent alcohol, 3.5 per cent solvent because a fairly long time is required before the of 20 per cent alcohol, 2 per cent of 30 per cent alcohol, 0.0 per conductivity ceases to change on standing. I n the presence cent of 40 per cent alcohol, and +1.9 per cent of 50 per cent of alcohol constant readings are obtained much sooner, alcohol. The combination of alcohol and 1 per cent acetic whereas alcohol itself has little effect upon the results. It is acid has a tendency to make all the results more uniform. recommended that not more than 30 per cent alcohol be
INDUXTRIAL A N D ENGINEERING CHEMISTRY
April 15, 1931
added. I n the presence of alcohol and acetic acid the results are much too low, especially when the alcohol concentration i s larger than 40 per cent. The titration of 0.01 N sodium sulfate gives results which are about 2 per cent low. Addition of alcohol, or alcohol with 1 per cent acetic acid, has qualitatively the same effect as described for potassium sulfate. With lithium sulfate, approximately the same results
131
Interpretation of Results
As we have seen above, the knick point is always found before the equivalence point, the error being especially large in the titration of 0.1 N potassium sulfate and zinc sulfate. Of course, one would feel very strongly inclined to attribute the low results to a coprecipitation of alkali or zinc sulfate with the barium sulfate. The occlusion of potassium and zinc, especially, should be very large in this case. However, the following experiments show that coprecipitation does not account for the low results. To different portions of 50 cc. 0.1 N potassium sulfate, normal barium chloride was added corresponding to 95, 96, 97, 98, 99, 100, and 101 per cent of the equivalent amount. After standing, the mixtures were filtered and the filtrates tested with more barium chloride. A decreasing cloudiness was observed in the first five filtrates, but even the filtrate of the mixture to which 99 per cent barium chloride had been added still contained some sulfate, whereas after addition of 100 per cent, the filtrates were sulfate-free. With zinc sulfate, another type of experiment has been made. About 10 per cent excess of barium chloride was added to the 0.1 N solution, the precipitate collected, and washed with water until free of barium. The zinc content in the filtrate was determined by the simple and accurate gravimetric o-oxyquinoline (oxin) method and a recovery of exactly 100 per Whna LIIW Cblonds I P M cent, *0.5 per cent, was found. Hence, there was no coFigure 1-Titration of Zinc Sulfate with Barium Chloride precipitation of zinc sulfate or zinc oxide (hydrolytic adsorpA aqueous solution tion) with the barium sulfate. In other experiments, separate E ’ 10% alcohol solution portions of 50 cc. 0.1 N zinc sulfate were added to 95, 97, 99, C; 20% alcohol solution D.30% alcohol solution 100, 101, and 110 per cent of the equivalent amount of E 40v alcohol solution P: 50% alcohol solution barium chloride; the precipitates were collected, washed out, and weighed. The weights found corresponded to 95.3, are found as with 0.01 N sodium sulfate. Comparable 97.7, 99.0, 100.1, 100.0, and 99.5 per cent precipitation. results are obtained in the titration of 0.005 N zinc sulfate These experiments show quite conclusively that coprecipitasolutions. When the concentration of the alcohol present is tion is not responsible for the low results of the conductonot larger than 30 per cent, the error is about -3 per cent; metric titrations. with more alcohol, the deviation increases rapidly (in 50 per cent alcohol about 7 per cent). It is peculiar that even in the titration of 0.005 N zinc sidfate, the angle between the precipitation line and reagent line is so much more obtuse than in the titration of 0.01 N alkali sulfate. Addition of alcohol makes the angle still more obtuse. As an example, Figure 1 represents the titration lines of zinc sulfate with barium chloride in water and different concentrations of alcohol. The straight line drawn through E . P. indicates the location of the equivalence point (theoretical end point). The slope is about the same for 0.05 N as for 0.005 N solutions. (Compare the angle with that in Figures 3 and 5, where titration lines of alkali sulfates are given.) Finally, the titration of very dilute sulfate solutions of the order of 0.001 N (46 mg. SO4 in a liter) was examined. There, of course, the gravimetric method is a very painstaking procedure and the conductometric method gains special significance. I n the or Buim ChlDndF absence of alcohol the titration is very impracticable; it Figure 2-Titration of 0.001 N L i t h i u m Sulfate with 0.1 N Barium Chloride takes too long a time before the conductivity is constant after A , 6 minutes after addition of solution addition of reagent. I n the presence of 20 to 30 per cent E . shaken after A alcohol, the results obtained were about 1.5 to 3 per cent high I n the titration of potassium sulfate with barium chloride, and quite well reproducible. I n Figure 2 is shown the titration curve of 0.001 N lithium sulfate in 30 per cent alcohol. the sulfato is replaced by the chloride ion until the equivalence For the estimation of the sulfate content of very dilute solu- point has been reached. I n order to be sure that the line tions which do not contain relatively large amounts of other which represents the conductivity of mixtures of 0.1 N electrolytes, the conductometric method is the best and yields potassium sulfate and chloride, respectively, had no irregular results which are more reliable and better reproducible than shape, it has been determined experimentally. Though the those obtained by the nephelometric method. Quite gener- line is not quite straight, the deviation is so small that it ally, it is recommended to add not more than 30 per cent cannot account for the displacement of the knick point in the alcohol. With more alcohol the precipitate often has a conductivity titrations. Therefore, the remaining possitendency to enter into colloidal solution, especially with an bility of explaining the deviations found in the conductometric titrations is that the presence of the precipitate itself excess of barium. ~~
3.P.
01
Yafolme
lddd
132
ANALYTICAL EDITION
affects the readings of the conductivity. Moreover, we have to consider the adsorbent properties of barium sulfate. When preiipitated from a sulfate solution, the precipitate will have a negative charge as long as there is an excess of sulfate due to some adsorbed alkali sulfate on the surface. This adsorption will decrease the more we approach the equivalence point and will be about zero a t that point. The precipitate will then have a tendency to adsorb barium ions (and acquire a positive charge) with an excess of reagent. That freshly precipitated barium sulfate has a tendency to adsorb barium ions could be proved experimentally. It even seems that the barium ion is much more strongly adsorbed
Vol. 3, No. 2
Four sets of determinations have been made, in which the points of the precipitation and reagent line were determined in the filtrate. Sometimes it was very difficult to get a clear filtrate. Moreover, great care has to be taken that no evaporation occurs during filtration. If the precipitate has no influence upon the conductivity, we might expect that the slope of the precipitation and reagent line would be the same in the filtrate as in the direct titration. Actually it was found that in the filtrates the slope of the precipitation line and, t o a lesser extent, of the barium chloride line, is much less than in the direct titrations without filtrations. This shows conclusively that the conductivity is higher in the presence of the precipitate than in the filtrates. The more barium sulfate present, the larger the difference between the reading in the unfiltered and filtered solutions. It may be that at the equivalence point the charge is due to adsorbed alkali chloride. In Figure 3, ABC represents the change in conductivity in the direct titration (measurements after 3 and 10 minutes, respectively), whereas ADE shows the precipitations and reagent line as measured in the filtrate. In the latter case, the knick point is found much nearer the equivalence point, the deviation being now of the order of 1 to 2 per cent. I n order to confirm the results, a set of readings has been made with mixtures of potassium sulfate and potassium chloride of the same composition as could be expected in the filtrate. The slope of this line coincides with that of the precipitation line in the filtrate. The higher conductivity in the presence of the barium sulfate cannot be attributed to
n Figure 3-Titration of Potassium Sulfate w i t h Barium Chloride A B C , measurement taken in presence of barium sulfate A D E , measurement taken with filtrate
than the sulfate ion; with an excess of the latter ion, there is no difficulty in getting a clear filtrate; with an excess of barium, however, the filtrate is usually cloudy and it requires a very dense filter to obtain a clear filtrate, Moreover, if there is 40 to 50 per cent alcohol present, the precipitate shows a strong tendency to go into a colloidal suspension. The barium sulfate precipitated at room temperature consists of very fine particles; on standing, the size probably increases and the adsorption decreases. For this reason, somewhat better results may be expected if the conductivity is measured after some time of standing. If the barium sulfate contains some sulfate or barium adsorbed a t its surface, the charged particles will show some conductance of the electric current when suspended in the solution. Therefore, it could be expected that the knick point would be found much nearer the equivalence point when the conductivity was measured in a clear solution from which the precipitate had settled or had been removed by filtration. Special experiments were conducted to study the influence of the precipitate. As a matter of fact, the following experiments have no direct analytical significance as the technic is very impracticable; the results, however, may contribute to an understanding of the error in the direct titration of sulfate with barium chloride. Potassium sulfate 0.1 N was titrated with 1 N barium chloride. The conductivity was measured 3 and 10 minutes, respectively, after addition of the reagent. Moreover, readings have been made after 10 minutes’ standing and after stirring up the precipitate. Four sets of determinations have been made. in^ some cases the points representing the conductivity data deviated less from a straight line than other cases. Under the different conditions mentioned, the knick point was found 4.5 to 6 per cent before the equivalence point.
C
Figure 4-Titration
/
I
of Barium Chloride w i t h L i t h i u m Sulfate
its solubility as the latter is very small and moreover has the same value in the filtrates. The only facts to account for the irregularities and deviation in the presence of the precipitate are that the barium sulfate increases the conductivity of the solution because of its charge, and that the cell constant may be affected more or less by the presence of the suspended particles. Titration of Barium Chloride with Sulfates
The results of the titration of barium chloride with sulfates are reported in Table 11. I n the titration of 0.1 N barium chloride with an N sulfate solution, the knick point is found * 1 per cent before or after the equivalence point. With potassium sulfate as a reagent, the deviation was about +1 per cent, with sodium sulfate - 1.5 per cent, and with lithium sulfate -0.5 per cent. There-
INDUSTRIAL AND ENGINEERING CHEMISTRY
April 15, 1931
fore, the mutual agreement is much better and the knick point much closer to the equivalence point than in the reverse titration of the different alkali sulfates with barium chloride. From the last two sections of Table I1 we see that addition of alcohol shifts the knick point more to the right, resulting in a larger error. With 50 per cent alcohol as a solvent, the deviation was +5 per cent using sodium sulfate as a reagent. One per cent acetic acid and especially the combination of alcohol and acetic acid have a tendency to increase the deviation still more.
133
if the concentration of the latter is lower than about 30 to 40 per cent. Of special practical significance again is the conductometric titration of very dilute barium solutions of the order of 0.001 N . I n the presence of 30 to 40 per cent alcohol, the results are quite reproducible though the knick point appears about 3 per cent before the equivalence point. It is interesting to note that the angle between the precipitation and reagent line becomes more obtuse with increasing alcohol content of the medium, especially the slope of the
Table 11-Titration of Barium Chloride with Sulfates BARIUMAMOUNT BARIUM CONCN. T A K k N FOUNDDEVIATION SOLVENT
N
Millimols Millimols
%
POTASSIUM SULFATE
0.1 0.1 0.1
2.514 2.514 2.514
2.541 2.566 2.640
+1
0.01 0.01 0.01 0.01 0 01 0.01 0.01 0.01 0 01 0.01
0.2503 0.2503 0.2503 0.2503 0.2503 0.2503 0.2503 0.2503 0.2503 0.2503
0.2539 0.2569 0.2559 0.2554 0.2564 0.2554 0.2554 0.2564 0.2600 0.2637
$1.4 4-2.6 +2.2 +2.0 +2.4 f2.0 +2.0 +2.4 f3.8 +5.4
0.1 0.1 0.1 0.1 0.1 0.1 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01
2.514 2.514 2.514 2.514 2.514 2.514 0.2503 0.2503 0.2503 0 2503 0.2503 0.2503 0.2503 0.2503
2.476 2.572 2.642 2.572 2.567 2.688 0.2509 0.2514 0.2552 0.2586 0.2528 0.2562 0.2567 0.2586
$2 +5
Water alcohol alcohol 4- 1 % acetic acid Water 10% alcohol 20% alcohol
:$
igQ :;::E 5 0 8 alcohol
1% acetic acid 10% alcohol 20% alcohol 40% alcohol
++ +
SODIUM SULRATE
-1.5 +2. 1 +5.1 +2.3 +2.1 +6.9
+0.2 +0.5 f2.0 $3.3 11.0 +2.3 +2.6 $3.3
Water lO$-alcohol alcohol 17 acetic acid 10 alcohol 17 acetic acid 1% acetic acid 5 0 k alcohol Water 10% alcohol 20% alcohol 50% alcohol 1% acetic acid acid 10% alcohol acid 20% alcohol acid 50% alcohol
++
++ +
LI7“IU M SULF‘ATE
0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01 0.01
0.2503 0.2503 0.2503 0.2503 0.2503 0 2503 0.2503 0.2503 0.2503 0.2503 0 2503
0.2486 0,2490 0.2486 0.2486 0.2505 0.2505 0,2503 0.2503 0.2473 0.2477 0.2542
-0.7 -0.5 -0.7 -0.7 +0.1 +o. 1 0.0 0.0 -1.2 -1.0 +1.5
0.01
0.2503
0.2551
C1.9
Watera Waterb alcohol %% alcoholb ,507’ alcohol
9
50 alcoholb 1% acetic acid 1 % acetic acidb 50% alcohol 4- 1 % acetic acid 50% alcohol 1% acetic
+
0.0239 -5.3 0.001 0.0252 0.0239 -5.3 0,001 0.0252 0.0243 -3.5 0,001 0.0252 0.0244 -3.1 0.001 0.0252 0.0244 -3.3 0.001 0.0252 0.0244 -3.3 0,001 0.0252 0.0245 -2.8 0.001 0.0252 -2.8 0.0245 0.001 0.0252 a , Reading taken after 5 minutes. b Shaken.
Figure 5-Titration of Lithium Sulfate with Barium Chloride A 17’ acetic, 50% alcohol solution B: 1% acetic acid solution
precipitation line where the barium is replaced by the alkali ion, which is much less in the presence of alcohol. The effect is very much pronounced, as may be seen in Figure 4. Here the change of conductivity in the titration of 0.01 N barium chloride in water, 10,30, and 50 per cent alcohol, respectively, with lithium sulfate is represented. On the other hand, in the titration of sulfate solutions, alcohol has very little influence upon the slope of the precipitation line, whereas the slope of the reagent (barium chloride) line is much flatter in the presence of alcohol. This is shown in Figure 5, where lines are given for the titration of 0.01 N lithium sulfate in water and 50 per cent alcohol, respectively, with barium chloride. From these phenomena, it may be inferred that the mobility of the barium ions is much more decreased by alcohol than that of the alkali chloride and sulfate ions. Whereas the mobility of the barium ions in water is much larger than that of the sodium ion, the difference between the two is very small in 50 per cent alcohol.
I n the titration of 0.01 N barium chloride with any of the three alkali sulfates, theoretical results are found within 1 per cent. It is practicable to add some alcohol if we are dealing with 0.01 N or more dilute solutions, as the conductivity then soon becomes constant after addition of the reagent. In concentrations lower than 30 per cent, the alcohol has almost no influence upon the location of the knick point. The same holds for the combination of 1per cent acetic acid and alcohol
(1) Dutoit, P., J . chim. phys., 6, 27 (1910); for literature review compare I. M. Kolthoff, “Konduktometrische Titrationen,” p. 69, Steinkopff,
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for their operation 45 horsepower per hour. Each unit consists of 12 arc lamps, each requiring 60 amperes a t 50 volts across the arc. The top carbons are 0.875 inch in diameter; those a t the bottom, 0.50 inch; and a grade of cored carbons, designated as “C carbons,” which give the most ultra-violet light, are those employed. Similar units are in operation a t their Reidsville and Durham plants. Five and one-half units are required for this operation in all of their plants.
Literature Cited
Dresden, 1923. (2) Jander, G., and Pfundt, 0.. “Die visuelle Leitfahigkeitstitration,” p. 32, Verlag Enke, Stuttgart, 1929. (3) Kolthoff, I. M., IND.END.CHEM.,Anal. Ed., 2, 225 (1930).
