INDUSTRIAL A N D ENGINEERING CHEMISTRY
April, 1926
389
Analysis of Lime' By John C. Bailar THE GREATWESTERNSUGAR Co., DENVER, COLO.
IT0 methods for the determination of calcium oxide in the presence of calcium carbonate have been developed in this laboratory and used with satisfactory results. These methods are (1) titration with a solution of iodine, and (2) titration with a solution of zinc chloride.
T
Iodine Method
SOLUTIOSS--4solution of iodine made by dissolving 45.27 grams of iodine and 90 grams of potassium iodide in 200 cc. of water and diluting to 1 liter. One cubic centimeter of this solution is equivalent to 0.01 gram of calcium oxide. A solution of sodium thiosulfate of such strength that 2 cc. are equivalent to 1 cc. of the iodine solution. This requires 44.27 grams per liter. STAh-DfiDIZATION-TheSe solutions may be standardized by any of the well-known methods. Calcium and iodine react in the ratio of 1 mol of CaO to 2 mols of I in this determination. DETER11IINATION-\veigh 1 gram of lime into a dry liter flask and hydrate by adding 500 cc. of boiling water. Add an excess of iodine, shake, and allow to stand 15 minutes, then titrate the excess of iodine with the thiosulfate solution using starch indicator. Generally this lime will be hydrated in less than 10 minutes. NOTES-In order to tell when a sufficient e x c w of iodine has been added put 3 cc. of the iodine into 250 cc. of water. This will give the color of the solution when a proper excess has been added. When adding the iodine the color is sometimes quite dark a t first, but on shaking for a few seconds the proper color will appear. If the lime is weighed into a wet flask it will react with the small amount of water present and produce a high temperature. This will cause the lime acted upon to become somewhat inert. With active limes this determination has been made in 10 minutes but with ordinary limes 30 minutes is a fair time. Use only freshly boiled distilled water and add it to the lime as hot as possible. An objection to this method is the excessive cost of reagents where many determinations are to be made. Zinc Chloride Method
SOLUTIONS-A solution of zinc chloride made by dissolving 25 grams of zinc chloride in a liter of water and filtering. One cubic centimeter of this solution is equivalent to 0.01 gram of calcium oxide. A solution of sodium hydroxide of such strength that 1 cc. is equivalent to 1 cc. of the zinc chloride solution. An indicator solution made by dissolving 5 grams of phenolphthalein in 500 cc. of alcohol, dissolving 2 grams of alizarin cyanine green C in 500 cc. of water, and mixing the two solutions. Use from 0.5 to 1 cc. for each determination but after deciding on the amount that gives the most satisfactory results always use the same amount. A difference in the amount of indicator makes a difference in the end point. STANDARDIZATION-Standardize the sodium hydroxide against a standard acid. Then standardize the zinc chloride against the sodium hydroxide as follows: Put into a liter 1
Received March 9, 1925
flask 500 cc. of boiling water. 4 d d 25 to 50 cc. of the sodium hydroxide and the indicator. Titrate this with the zinc chloride until a fairly permanent green color is obtained, then add a t least 10 cc. more, shake vigorously, and allow to stand 15 minutes. Titrate the excess of zinc chloride with sodium hydroxide to a lavender (not pink) color. The lavender will change to pink on standing. Shake vigorously after the lavender appears and if it disappears add more sodium hydroxide. DETERMINATION-weigh 1 gram of lime into a dry liter flask and hydrate by adding 500 cc. of boiling water. When the lime is hydrated titrate as the sodium hydroxide was titrated in the standardization. NOTES-DO not boil after adding the zinc chloride, as calcium carbonate is decomposed by boiling zinc chloride. The proportion of phenolphthalein and green dye may be varied a little, but if too much green is used the end point will come too soon. The success of these, and most other methods, depends upon getting the lime completely hydrated. Lime cannot be suitably hydrated in cold water. With most limes there is a small amount that hydrates very slowly, even when using a large excess of boiling water. Sucrose Method The sucrose method as practiced in this laboratory is as follows: Weigh 5 grams of lime into a dry 200-cc. Kohlrausch flask and add about 150 cc. of boiling water. Shake occasionally until t h e lime is hydrated, then add 40 grams of granulated sugar. Shake occasionally until the lime is in solution. Cool to room temperature, fill t o mark, mix thoroughly, and filter through a dry filter. Discard the first 15 or 20 cc. Titrate 100 cc. with standard nitric acid of such strength t h a t 1 cc. equals 0.05 gram of lime, using phenolphthalein indicator. The buret reading multiplied by 2 gives the per cent available lime.
Experimental Results Sample A Iodine method: Average five determinations Difference between highest and lowest Zinc chloride: Average twelve determinations Difference between highest and lowest
Per cent
89.11 0.10
89.08 0.20
Sample B (Very active lime made by special process) Zinc chloride Sucrose Iodine Per cent Per cent Per cent
Av.
91.0 91.3 91.0 91.0 91.0 91.06
92.6 92.6 92.8 92.8
93.20 93.25 93.35 93.25
92.70
93.26
..
...
Samfiles of Limr from Some of Sleffen Houses Sucrose Per cent
No.
88.0 .88.0 79.0 79.2 84.6 84.6
Av.
82.6 82.4 84.0 84.0 80.6 80.6 87.0 87.0 83.4 83.4 83.65
-
Zinc chloride Per cent
Iodine P u cent
87.5 87.5 79.0 79.0 85.5 85.6 84.5
88.00 87.95 80.00 79.90 85.25 85.00
84.37
84.45
84.10
INDUSTRIAL A N D ENGINEERING CHEMISTRY
390
Discussion of Results “Available lime” means nothing until it is known for what purpose it is available. In the ammonia industry available lime is determined by boiling the sample with an excess of ammonium chloride and titrating the liberated ammonia. A sample of Portland cement, which may be considered an impure overburned lime, gave by this method 58.0 per cent, by the iodine method 16.5, by zinc chloride method 12.5, and by the sucrose method 9.8 per cent. The writer believes that the method is yet to be devised that will give the calcium oxide in lime without giving some of the silicates and other compounds present. Even pure water breaks down some silicates and the amount broken down depends upon the amount of water, the time of con; tact, and the temperature. It is possible that a method might be devised by using an-
Vol. 18, No. 4
hydrous reagents. The writer did some work along this line with promising results, but the methods are not suited to factory work. There is always difficulty in getting, and keeping, anhydrous reagents. Of the three methods given in this paper the iodine and zinc chloride are quicker than the sucrose. The sucrose method requires one filtration; the others do not. The sucrose method requires standard flasks and pipets, and it is necessary to cool it to room temperature before making up to volume. The other methods require no standard flasks or pipets and should be run hot, as soon as the lime is hydrated. Standard flasks and pipets give possibilities of error which are avoided in the other methods. Which of these methods gives the lime available in the sugar industry has not yet been determined. The writer believes that one of these, probably the zinc chloride method, will give satisfaction.
The Choice of Indicators for Alkaloidal Titrations’ By H. Wales DRUGCOXTROL LABOUTORY, BUREAUOB CHEXISTRU,WASHINGTON, D. C.
H E customary procedure in the determination of alkaloids is to dissolve the base in an excess of standard acid and titrate this excess with standard alkali, usually with methyl red as an indicator. The accuracy of this method has not been extensively investigated. Keblerllx and later Kippenberger2 titrated a number of alkaloids with various indicators and chose the ones that gave the best results. On account of their broad range these have been largely discarded for the newer and more sensitive indicators. ever^,^ who investigated the hydrogen-ion concentrations of a few alkaloidal salts by colorimetric methods, recommended certain of the newer indicators for their titrations. The hydrogen-ion concentrations of some alkaloidal salts have been determined by M ~ G i l l ,Masucci ~ and MoffatJ6 and Krantz’ by potentiometer methods, using the hydrogen electrode. As the hydrogen electrode is known to reduce alkaloids,* Rasmussen and Schou,I1 and Wagener and McGil15have used the quinhydrone electrode for determining hydrogen-ion concentrations. Kolthoff ,9.t Treadwell and Janett,12 and Dutoit and Meyer-Levy13 have investigated the titration of alkaloids by conductivity methods, and Popoff and McHenry14 have used a platinum electrode in titrating alkaloids. The discrepancies in the results obtained by many of these investigators are too great to be explained on the basis of the reduction of the alkaloid by-the hydrogen electrode. Masucci and MoffatG found a wide variation in the hydrogenion concentrations of commercial samples used with no preliminary treatment. Some explanation may be afforded by the work of Tutin,lb who showed-that quinine sulfate crystallized from slightly acid or alkaline solutions required several recrystallizations to make it neutral again. Rasmussen and Schoull dissolved alkaloids in an excess of acid and titrated back, electrolytically, with alkali. The hydrogen-ion values computed from their titration curves
T
1 Received
February 6, 1926.
* Numbers in text refer to bibliography at end of article.
t During the course of this investigation Kolthoff [Pharm. Wcckblad, 61, 1287 (1925)l published the pH values of a large number of alkaloidal salts, computed from the dissociation constants, and recommended indicators in several cases. The fallacies arising by choosing indicators from the pH of the salt are discussed in this paper.
agreed very closely, while those obtained by “direct measurement” often showed appreciable variations. These authors, however, do not state how their direct measurement values were obtained. The quinhydrone electrode has been investigated by Granger and NelsonI1GBiilmann,’’ Kolthoff,Bl10 La Mer and Parsons,l* and others. It has been used for titrations by Harris,lg Wagener and M ~ G i l l and , ~ Rasmussen and Schou.ll Procedure The procedure followed in the investigation here reported was essentially that of Rasmussen and Schou.ll The alkaloids used were commercial products obtained from wellknown manufacturers. One hundred milligrams of the alkaloid were dissolved in a slight excess of acid and the volume diluted to 50 cc. This excess of acid was then titrated with alkali, using quinhydrone and saturated calomel electrodes, and the voltages a t definite intervals were recorded. The variation in voltage per cubic centimeter of alkali was plotted against the volume of alkali used and the end point (center of break) of the titration determined. The hydrogen-ion values were computed from the formulazo pH =
-
-
0.4661 0.00014t 0.0541 0.0002c
+
T
where ?r is equivalent to the observed voltage plus a correction for the calomel cell, obtained by determining the voltage of known buffer solutions, and t is the temperature in degrees Centigrade. This formula is applicable between 15” and 30’ C . Wagener and McGillS have shown that the small quantity of neutral salt produced during the titration has a negligible effect on the hydrogen-ion values as determined by the quinhydrone electrode. Sulfuric and hydrochloric acid have been used in several cases. No differences greater than those due to experimental errors were obtained. Both 0.1 N and 0.02 N sodium hydroxide (carbonate-free) have been used with identical results. When an acid is titrated with an alkali the hydrogen-ion value of the solution changes slowly until all but a very small part is neutralized. With the neutralization of
