January 15, 1931
TNDUXTRIAL AND ENGINEERING CHEMISTRY
115
Direct and Reverse Titration of Sulfuric Acid with Barium Hydroxide' I. M. Kolthoff and E. B. Sandell SCHOOL OF
CHEMISTRY,
UNIVERSITY OF MINNESOTA, MINNEAPOLIS, MI".
In the precipitation of sulfuric acid with barium chloPotassium biphthalate (Bureau of Standards) is to be recommended for the standardization of barium hy- ride, a small part of the acid is occluded. In precipitation droxide. Oxalic acid is less suitable for this purpose as a t room temperature the occlusion is larger than a t the precipitate of barium oxalate carries down some higher temperature. The titration of barium hydroxide with sulfuric acid bioxalate or oxalic acid. By titrating in hot solution leads to erroneous results. A part of the base is occluded this error is decreased to 0.1 per cent. The titration of sulfuric acid with barium hydroxide and a part adsorbed by the precipitate. The adsorbed cannot be used for precise work. The best results are base can be removed by boiling the mixture with a small obtained if the titration is performed a t room tempera- excess of sulfuric acid; the error by occlusion is still about ture and if, after the first color change, the mixture is 1 per cent. The best results are obtained if the barium boiled and more base added until the color of the indi- hydroxide is added to a n excess of potassium sulfate at cator persists for 15 to 30 seconds. After continued boil- room temperature. The mixture is titrated a t room teming for 5 minutes, the solution is cooled and the titra- perature and, after the color change, boiled for 5 minutes tion finished at room temperature. The accuracy obtained with a small excess of sulfuric acid and titrated back a t is 0.1 to 0.2 per cent. In the titration of a hot solution room temperature. The accuracy obtained is 0.0 to 0.2 per of sulfuric acid there is a greater occlusion of barium cent. If the potassium sulfate solution is added to the barium hydroxide, about 1 per cent of the base is occluded. hydroxide. . . . . . .. . . . . . . .
I
T IS a well-known fact that barium sulfate precipitated from a sulfate or barium solution has a tendency to carry down other ions present in the solution. In the titration of sulfuric acid with barium hydroxide or in the reverse procedure, the possibility exists that either sulfuric acid or barium hydroxide may be co-precipitated. No accurate data on the accuracy of this titration are to be found in the literature. E. von Drather (2) concluded from rough experiments that in the titration of 0.2 N barium hydroxide with 0.2 N sulfuric acid, about 2 to 3 per cent too little base was found. A special study of the deviations found in the titration of sulfuric acid with barium hydroxide and in the reverse procedure has been made. All titrations were made with weight burets, and precautions were taken against contamination by atmospheric carbon dioxide. About 0.2 N sulfuric acid and 0.2 N barium hydroxide solutions were prepared in carbon dioxide-free water and the normality of both determined by means of various standard substances. From the figures obtained, the normality ratio sulfuric acid to barium hydroxide was calculated and compared with the experimental ratio found by direct titration. As the calculated value is a little uncertain owing to errors made in the standardization of the acid and base, the same procedure was followed with hydrochloric acid and sodium hydroxide. In the latter case the experimental ratio can be determined without encountering the errors arising from the formation of a precipitate during the titration. Pure standard substances were prepared and tested for purity according to the directions given by one of the authors (3). The results have been corrected for the titration error, which was determined in an experimental way (for details compare Kolthoff, 3). The results are given in Table I. The experimental and the calculated ratios agree within 0.02 to 0.03 per cent. Calculated ratio of normality : sodium hydroxide to hydrochloric 0 18914 acid = 0.18084 = 1'0460 Ratio found: 1.0462, 1.0462, 1.0463 (bromocresol green) 1.0464, 1.0464 (phenolphthalein) Ratio found: Average: 1.0463 I
Received September 26, 1930.
T a b l e I-Standardization of R e a g e n t s STANDARD SUBSTANCE NORXALITY FOUND INDICATOR Equivalents 9ev 1000 grams HYDROCHLORIC ACID
0.18091, 0.18079, 0.18082, 0.18087, 0.18087; av. 0.18085 NazCOs (product 11) 0.18083, 0.18085; av. 0.18084 KHCOJ (recryst. from HzO) 0.18092, 0.18086, 0 18084; av. 0.18087 KHCOs (pptd. with alcohol) 0.18084, 0.18083, 0.18084, av. 0,18084 Borax I 0.18081, 0.18080, 0.18086; av. 0.18082 Borax I 0.18082, 0,18082; av. 0.18082 Borax I1 0.18083, 0.18085; av. 0.18084 Average normality: 0.1 8084
NarCOs (product I)
Bromocresol green Methyl orange Bromocresol green Bromocresol green Bromocresol gieen Methyl red Bromocresol green
SODIUM HYDROXIDE
Benzoic acid (B. S.)
0.18912,0.18913,0.18910;
av. 0.18912 Benzoic acid (Kahlbaum) 0.18918, 0.18913, 0.18911; av. 0.18914 Oxalic acid 0.18917, 0.18915, 0.18918; av. 0.18917 Oxalic acid (fMgC12) (0.18935, 0.1894) Potassium biphthalate 0,18909, 0.18912, 0.18915; av. 0.18912 Average normality: 0.18914
Phenolphthalein Phenolphthalein Phenolphthalein Bromocresol greeq Phenoiphthalein
BARIUM HYDROXIDE
Benzoic acid (B. S.) Potassium biphthalate ( Q ) Potassium biphthalate (b) Oxalic acid ( c ) Oxalic acid ( d )
0,19229, 0.19231, 0.19231, 0.19233; av. 0.19231 0,19221, 0.19222 0.19232, 0.19233, 0.19234;
- . . n-.-1~2.13 - - --
0.19300, 0.19305 0.19248, 0.19250, 0.19256 Average normality. 0.19232
The sulfuric acid was standardized against the various standard substances used for the hydrochloric acid solution, an average normality of 0.20493 (equivalents per 1000 grams) being found. A few words may be said about the standardization of barium hydroxide with potassium biphthalate and oxalic acid, respectively. In a titration of an oxalic acid solution with barium hydroxide, the barium oxalate precipitating during the procedure carries down some oxalic acid or bioxalate. If the mixture is heated to boiling after the first color change of phenolphthalein, then cooled to room temperature and the titration continued until the color becomes pink again, the deviation is still 0.35 per cent from the theoretical value. By direct titration at room temperature Bruhns ( 1 ) found an error of 0.24 per cent, wherea8
ANALYTICAL EDITION
116
Schmitt (4) found a deviation of 0.3 per cent. According to the latter the titration of a hot solution of oxalic acid gives theoretical results. This statement was not confirmed in the present work, a deviation of 0.1 per cent from the theoretical value being found, if a hot solution of oxalic acid was titrated with barium hydroxide. Therefore, for work of high precision oxalic acid cannot be recommended for the standardization of barium hydroxide. Potassium biphthalate, on the other hand, is very useful, and gives theoretical results if a 0.05 N solution is titrated at a temperature of 40" to 50" C. At room temperature some barium phthalate may crystallize out at the end of the titration and carry down a trace of barium hydroxide. The error, however, is very small, amounting to about 0.05 per cent. In Table I the results of the standardization of barium hydroxide with benzoic acid, potassium biphthalate, and oxalic acid, respectively, are also reported. The titration of potassium biphthalate (a) was carried out a t room temperature, the results given under (a) were obtained at 40" to 50" C., and the results reported for oxalic acid (c) were found by titrating a t room temperature until a temporary end point was reached. The mixture was heated to boiling, cooled with protection of a soda-lime tube, and the titration finished with barium hydroxide (error 0.35 per cent). In experiment (d) the boiling solution of oxalic acid was titrated with barium hydroxide. After the end point had been reached, the color of the indicator did not change if the mixture was heated to boiling and cooled to room temperature.
Experiment 3b. After the precipitation of barium sulfate with barium chloride, the liquid was heated to boiling and kept a t 100" C. for an hour, and titrated after standing overnight; the titration was made as in Experiment 3a. Experiment 3c. Hot barium chloride was added drop by drop t o hot sulfuric acid. Titration with barium hydroxide was made in hot solution; near the end point the liquid was boiled for a few minutes. The final end point was obtained a t room temperature. If the precipitate was allowed t o stand an hour a t 100" C. and then titrated in the cold after standing overnight, practically the same results were obtained. Experiment 4. Titration was made in hot solution; the final end point was obtained after cooling. Experiment 5 . The mixture was titrated after cooling to room temperature, the final end point obtained after boiling the liquid and cooling.
Titration of Barium Hydroxide with Sulfuric Acid
The titration of barium hydroxide with sulfuric acid is given in Table 111. Table 111-Titration of Barium Hydroxide w i t h Sulfuric Acid Calculated normality ratio HzS04 :Ba(OH)z, 1.066 Av.
EXPERIMENT
la lb
E3
Table 11-Titration of Sulfuric Acid with Barium Hydroxide Calculated normality ratio HzSOo : Ba(OH)z, 1.066 Av. NORMALITY Av. DEVIANo. RATIO DEVIA- TION OF HzSO4: TION FROM EXPERITITRA-Ba(0H)z FROM TRUE MENT ADDITION TEMP. TIONS FOUNDMEAN RATIO O
1 20 2b 3a
3b
3c 4 5
...... ...... I
.
.
.
.
.
0.5 N BaClz added in excess to HzSOo at room temp. 0.5 N BaCh added in excess to HzSOn at room temp. 0.5 N BaCh added in excess t o HzSOo at 100' C. HzSOa a d d e d t o e x c e s s BaClz at 1000 c. Ba(N0dz added to HzSO4at1OO"C.
c.
%
%
Room 100 100
4 3 3
1.067 1.072 1.070
0.1 0.2 0.1
+o. 1
Room
2
1.057
0.0
-0.85
Room
2
1.057
0.05
-0.85
100
2
1.062
0.0
-0.4
fO.6
$0.4
Room
2
1.062
0.1
-0.4
Room
2
1.062
0.0
-0.4
Explanatory R e m a r k s Experiment 1. The end point obtained in the cold was not permanent. Therefore, the liquid was heated t o boiling and barium hydroxide added until the mixture was slightly alkaline; the boiling was continued for 5 minutes and the titration then finished a t room temperature. Experiment 2a. Barium hydroxide was added rapidly to the hot sulfuric acid solution. When the end point had been reached (first tinge of pink) the mixture was boiled for 5 minutes; the pink color of phenolphthalein faded only slightly during the boiling. The titration was finished a t room temperature. Experiment 2b. In this case barium hydroxide was added very slowly to the hot sulfuric acid. End point obtained as in Experiment 2a. Experiment 3a. Twenty per cent excess neutral 0.5 N barium chloride solution was added drop by drop to the sulfuric acid, and the titration made immediately after the precipitation; after the first color change the mixture was boiled for 5 minutes and the titration finished a t room temperature.
TEMP.
ADDITION
36
4b 5a 5b
Ba(0H)z added to excess K&Oa at room temperature Ba(0H)z added to excess K&Oo at room temperature Ba(0H)z added to excess KzSOo at 1000 c. Ba(OH)z added t o excess KzSOo at 1000 c. Excess KzSOa added to Ba(0H)z at room temperature Excess KzSOr added to Ba(0H)z at room temperature
OF NOR- TION TITRA-MALITY FROM TIONS RATIO MEAN
c.
FROM
TRUE
RATIO
%
%
Room Room 100 100
3 2 2 3
1084 1.076 1082 1.075
0 3 0.15 0.1
0.1
$1.7 +0.95 $1.5 $0.85
Room
2
1 070
0 0
$0 4
Room
3
1.064
0.2
-0.2
Room
3
1.083
0.3
+1.6
Room
3
1.066
0 1
0.0
Room
2
1.083
0.1
+l.6
Room
3
1.078
0.3
$1.1
O
. ....
2a
2b 3a
40
Calculated ratio of normality : sulfuric acid to barium hydroxide = = 1.0660 0.19223
Av. DEVIADEVIA- TION
No.
Titration of Sulfuric Acid w i t h B a r i u m Hydroxide
The titration of sulfuric acid with barium hydroxide is given in Table 11,
Vol. 3, No.1
Explanatory R e m a r k s Experiment l a . The average normality ratio was obtained by taking the first color change in the cold as the end point. When the mixture was boiled the liquid became alkaline owing t o slow liberation of co-precipitated barium hydroxide. Thus, in a particular titration the ratio found by taking the end point as the first disappearance of the pink color of phenolphthalein a t room temperature was 1.087; after boiling 5 minutes, cooling, and adding acid to restore neutrality, the ratio became 1.085, and after another 5-minute period of boiling, 1.083, etc. Experiment l b . When the end point had been reached in the cold, about 1 gram of the standard sulfuric acid was added in excess and the mixture boiled for 10 minutes, then cooled, and the acid back-titrated with barium hydroxide. Experiment 2a. After the color change in hot solution, the liquid was kept a t the boiling point for 10 minutes, then cooled, and the titration finished a t room temperature. The amount of barium hydroxide liberated by boiling for 10 minutes corresponded to about one drop of 0.2 N sulfuric acid. Experiment 2b. After the color change a t 100" C., about 1 gram of sulfuric acid was added in excess and the mixture boiled for 10 minutes; after cooling to room temperature the titration was finished with barium hydroxide. Experiment 3a. The barium hydroxide was added drop by drop to 50 cc. of 0.5 N potassium sulfate solution and the mixture then titrated with sulfuric acid. Since the end point obtained in the cold was not permanent, the mixture was kept a t 100' C. for 5 minutes, then cooled, and the titration finished a t room temperature with sulfuric acid. The ratio thus obtained is given in Table 111. After boiling the neutralized mixture for the nearly an hour, the ratio 1.066 was obtained-nearly theoretical value. Experiment 3b. Barium hydroxide was added a t room temperature to potassium sulfate as in Experiment 3a, but when a temporary end point had been obtained in the cold by titration with sulfuric acid, 1 gram of acid was added in excess and the liquid boiled for 10 minutes; the titration was finished as always at room temperature.
January 15, 1931
INDUSTRIAL AND ENGINEERING CHEMISTRY
117
Experiment 4a. Fifty cc. of 0.2 N potassium sulfate were used. The barium hydroxide was added slowly and with constant shaking of the hot potassium sulfate solution. After cooling, the first end point obtained with sulfuric acid gave the ratio recorded in Table 111. On boiling the mixture, barium hydroxide was slowly liberated. Thus, by boiling for 45 minutes and adding sulfuric acid whenever the liquid became alkaline, the value of the ratio finally obtained was 1.071. Experiment 4b. As in Experiment 4a, but when the end point had been reached at room temperature, a gram of 0.2 N sulfuric acid was added in excess and the mixture boiled 10 minutes; the titration was then finished in the cold with barium hydroxide. Experiment 5a. Fifty cc. of 0.5 N potassium sulfate a t room temperature were slowly added t o the cold barium hydroxide solution and the mixture titrated cold with sulfuric acid. Experiment 5b. Same conditions as Experiment 5a, but when the end point was reached in the cold a slight excess of 0.2 N sulfuric acid was added and the mixture boiled for 10 minutes; the titration was finished, after cooling, with barium hydroxide.
as a result of growth of the small particles of the barium sulfate to larger crystals. However, as shown by Experiment 3c, Table 11,heating had no effect. If a hot solution of sulfuric acid is precipitated with a warm solution of barium chloride, the error due to occlusion is smaller though it still amounts to about 0.4 per cent. It seems that two factors are responsible for the deviating results : adsorption and occlusion. Occluded sulfuric acid or barium hydroxide cannot be removed by boiling or by treatment with excess of base or of acid. The adsorbed acid can be removed fairly rapidly by boiling the mixture near the equivalence point for 5 minutes. By experiment, it could be shown that freshly precipitated barium sulfate adsorbs sulfuric acid. An excess of a hot solution of barium chloride was added to a warm solution of sulfuric acid. After 5 minutes’ standing the mixture was filtered, and the precipitate washed with water until the reaction of the filtrate Discussion of Results was neutral. The total filtrate was titrated with barium TITRATION OF SULFURIC ACIDWITH BARIUM HYDROXIDE- hydroxide. The ratio found was 1.037, thus indicating an Under no conditions are highly accurate results obtained; adsorption of about 3 per cent sulfuric acid by the precipitate. moreover, they are not so well reproducible as in ordinary TITRATIONOF BARIUM HYDROYIDE WITH SULFURIC acidimetric titrations where no precipitate is formed. There- Acm-In the titration of a cold or hot solution of barium fore, the titration is not suitable for work of precision. The hydroxide with sulfuric acid a strong co-precipitation of the best results are obtained if the acid is titrated at room tem- base by the precipitate occurs, the error amounting to about perature with barium hydroxide using phenolphthalein, 2 per cent (Experiment la, Table 111). This adsorbed base is phenol red, or methyl red as an indicator. After the first only slowly removed by boiling the mixture near the equivacolor change, the mixture is heated to boiling, more base lence point, It is better to add a small excess of sulfuric acid, is added until the color persists for 15 to 30 seconds. The boil for 5 minutes, and titrate back. The error then is still boiling is continued for 5 minutes and the titration finished about 0.8 to 1per cent and is caused by an occlusion of barium after cooling to room temperature (protection with soda-lime hydroxide by the precipitate (Experiment l b ) . tube). It seems that a trace of barium hydroxide can be From theExperiments 3a to 4b (Table 111) it may be incarried down by the barium sulfate (see Experiment 1, Table ferred that if barium hydroxide is added to an excess of neutral 11); accuracy 0.1 to 0.2 per cent. For most practical work, potassium sulfate solution, an adsorption of the base by the these results are satisfactory. precipitate takes place. This adsorbed base can be removed If a hot solution of sulfuric acid is titrated, the deviation by boiling the mixture near the equivalence point for a few is much larger, amounting to about 0.6 per cent (Experiments minutes with a small excess of sulfuric acid. After back titra2a and 2b, Table 11). Unexpectedly it is found that this tion a t room temperature the error is 0.0 to 0.2 per cent. By error is due to co-precipitation of barium hydroxide with the boiling the mixture near the equivalence point without adding barium sulfate although an acid solution is titrated. If an excess of acid, the adsorbed base is very slowly removed. the base is run in with great speed, a local excess is formed; If the potassium sulfate solution is added to the barium hybetter results may be expected if the barium hydroxide is droxide, an adsorption and occlusion of the base takes place. added drop by drop with continuous shaking of the flask. The error by occlusion is about 1per cent. That this is really the case is shown by Experiment 2b, though Literature Cited the deviation is still 0.4 per cent. If the sulfuric acid is precipitated at room temperature with an excess of barium chloride, a co-precipitation of sulfuric (1) Bruhns, G., 2. anal. Ckem., 66, 23 (1916). Drather, E. von., Ckcm.-Ztg., 62, 518 (1928). acid takes place. The error is about 0.9 per cent (Experiment (2) (3)rKolthoff, I. M.,“Volumetric Analysis,” Vol. 11, translated by N. H. 3a). It was expected that heating the mixture after precipiFurman, Wiley, 1929. tation a t room temperature might decrease the deviation ( 4 ) Schmitt, K. O., 2. anal. Ckcm., 71, 284 (1927).
NOTES A N D CORRESPONDENCE Colorimetric Determination of Silica Editor of Industrial and Engineering Chemistry: In a recent issue [IND. ENG.CHEM.,Anal. Ed., 2, 276 (1930)l Thayer has made an exhaustive criticism of the well-known Dienert and Wandenbulcke [Compt. rend., 176, 1478 (1923)l method for the estimation of silica in natural waters. He has shown that in the presence of phosphate and iron the yellow color, which develops on adding ammonium molybdate and sulfuric acid to a dilute solution of silica, is considerably enhanced. In the case of sea water and many other natural waters, where phosphorus and
iron are present only in minute amounts, the method of Dienert and Wandenbulcke can be used without modification. But where the waters contain considerable amounts of these elements, it is necessary t o precipitate the phosphate and the iron and determine the silica in the filtrate. Thayer has shown that phosphate can be precipitated satisfactorily as calcium triphosphate a t a weakly alkaline reaction but that in the presence of iron, in amounts of 25 mg. and more per liter, part of the silica is precipitated along with the phosphate and iron (Table IX). H e has found, however, that iron can be precipitated in acetic acid solution as ferric phosphate without loss of silica, and the excess of phosphate then precipitated as calcium triphosphate.
