3. The titration must be made in the cold with carbonate-free alkali and the necessary precautions taken to prevent the access of carbon dioxide from the air during the titration. Phenolphthalein is the most satisfactory indicator.

4. In the calculation of results, the acid value of the sample is to be taken as 0.9990 that of pure benzoic acid.

For full directions regarding the use of benzoic acid, see articles by George W. Morey, J. A. C. S. 34, 1027 (1912), and E. R. Weaver, ibid. 35, 1309 (1913).

Remarks. 25 c.c. 0.1 M NaOH = 0.3051 g. C2H5COOH. In using alcohol to dissolve the benzoic acid, it is preferable, though not absolutely necessary, to use alcohol which has been redistilled from potassium hydroxide, because such treatment of the alcohol eliminates small amounts of nitrogenous and acid principles ordinarily present in alcohol, thereby reducing the size of the blank. It will be found that some of the alcohol will soon evaporate during the first part of the titration leaving a slight ring of benzoic acid around the walls of the titrating vessel. Therefore, it becomes necessary toward the end of the titration to wash down the walls of the titrating vessel carefully to make sure that the deposited benzoic acid is included in the titration.

135. Preparation and Standardization of Approximately 0.1 Molar Sodium Hydroxide. Pour into a clean liter flask about 600-700 c.c. distilled water, then add 4.5 g. stick sodium hydroxide, taking care that the hydroxide is free from any visible encrusted carbonate; make up to the mark with water and thoroughly mix the solution by shaking; transfer the solution to a clean stock bottle. Select three beakers of about 300 c.c. capacity each, and then weigh out upon a tared watch-glass the benzoic acid in successive portions of about 0.300 g., 0.310 g., and 0.320 g. to the nearest tenth of a milligram and transfer the several portions to the respective beakers, observing the precautions of handling the watch-glass as mentioned in § 130. After the several portions have been weighed out, add 10 c.c. alcohol and 3-5 drops phenolphthalein solution to each. A standard 50 c.c. burette is now filled to the zero mark with the sodium hydroxide solution, and the titration of the benzoic acid portions conducted with due observance of the precautions mentioned in § 134. The first faint

pink is taken as the end point; a blank is run on the same volume of alcohol, and correction applied accordingly. The average deviation of the three titrations should not be greater than 1 part per 1000.

If a supply of benzoic acid is not at hand, the sodium hydroxide solution may be standardized by means of standard acid as follows: pipette out by means of a standard pipette three 25 c.c. portions of the standard acid, approx. 0.1 molar, delivering each portion into a separate beaker; add 100 c.c. distilled water and 3-5 drops phenolphthalein solution respectively, and then conduct the titration by adding the sodium hydroxide solution from the burette until the end point is reached. The average deviation of the three titrations should not be greater than 2 parts per 1000.

136. Exercise No. 3. Prepare and standardize 1 liter of approx. 0.1 molar hydrochloric acid.

137. Exercise No. 4. Prepare and standardize 1 liter of approx. 0.1 molar sodium hydroxide solution.

138. Exercise No. 5.

-

Determine the amount of hydrochloric acid in an unknown solution of same furnished by the instructor. 139. Exercise No. 6. Determine the amount of sodium carbonate in an unknown sample furnished by the instructor.

140. Examples.

1. In standardizing an approx. 0.1 molar solution of hydrochloric acid by titration against sodium carbonate, using methyl orange as indicator, the following results were obtained:

0.1317 g. Na2CO3 required 22.83 c.c. HCl

What was the molarity of the hydrochloric acid? What was the average deviation? Should any of the results be rejected?

Ans. 0.1089 molar. 1 part in 2000. No

2. How many c.c. of a 0.1039 molar HCl solution would be required to titrate 0.2905 g. K2CO3, using methyl orange as the indicator?

Ans. 40.46 c.c.

3. 0.2047 g. Na2CO, required 34.45 c.c. of a certain hydrochloric acid solution to neutralize it when methyl orange was used as indicator. 27.54 c.c. of a sodium carbonate solution required 33.46 c.c. of the same hydrochloric acid for neutralization with methyl orange as indicator. What is the molar strength of the sodium carbonate solution?

Ans. 0.0681 molar

4. 30.00 c.c. of an approximately 0.2 molar HCl solution gave 0.8042 g. of AgCl. 32.00 c.c. of this acid were found to neutralize 34.62 c.c. of a sodium carbonate solution, using methyl orange as indicator. Find the value of the sodium carbonate solution. Ans. 0.0865 molar

5. 0.2560 g. of a mixture containing Na2CO3 and NaOH was titrated with 0.1 molar HCI. It took 24.3 c.c., using phenolphthalein as indicator, and then 19.17 c.c. more, using methyl orange as indicator. How many g. each of NaOH and of Na2CO3 were present? Ans. 0.0205 g. NaOH. 0.2032 g. Na2CO;

6. 20.0 c.c. H3PO4 solution, diluted to 150 c.c., required 29.75 c.c. of 0.1 molar NaOH to reduce the concentration of H+ to 10-4-5 and 29.5 c.c. more to reduce it to 10-8-0. What is the molar strength of the phosphoric acid, and what indicators should be used? Ans. 0.1481 molar

7. 2.00 c.c. (by burette) of a sample of "potash liquor" were diluted to about 300 c.c. Using phenolphthalein as the indicator, 22.99 c.c. of 0.9523 molar HCl were required, and upon adding methyl orange 0.52 c.c. more of the HCl were required. The density of the original solution was 1.444 at 20°. What was the percentage by weight of the KOH and K2CO2?

8. In order to render citric acid (CO2H CH2)2 • C(OH)·CO2H + H2O anhydrous, Sörensen16 recommends heating the acid at 70° under 20-30 mm. pressure until constant weight is obtained. In testing a lot of citric acid, which had been so treated, 0.4004 g. of sample was dissolved in water and titrated with barium hydroxide solution and found to require 25.82 c.c., using phenolphthalein as indicator. 20.00 c.c. of the barium hydroxide solution were the equivalent of 48.42 c.c. of 0.1000 molar HCl solution. What percentage of anhydrous citric acid did the sample contain? Ans. 99.95%

9. Verify the fact that when a weight of sample is taken which is one hundred times the value of 1 c.c. of the standard solution expressed in grams of the constituent being determined, the burette readings in c.c. give percentages direct.

10. A sample of soda weighing 25.00 g. is dissolved in water and made up to 250.0 c.c., and one-fifth of this solution is taken for titration. What must be the normality of the standard acid (assuming the alkalinity to be due wholly to Na2CO3) in order that twice the number of cubic centimeters of acid used shall indicate the percentage of Na2CO3 in the sample? (G. McPhail Smith, loc. cit., § 13, p. 197.) Ans. 1.887 N

11. In the estimation of small quantities of free CO2 in water, an excess of Ba (OH)2 is added to precipitate the CO2 as barium carbonate. The excess of Ba(OH)2 is then titrated with oxalic acid, using phenolphthalein as indicator. How many c.c. of 0.4875 molar oxalic acid must be diluted to 1 liter in order to give a solution, 1 c.c. of which would correspond to 1 mg. CO2 in the above method? Ans. 46.62 c.c.

16 Biochem. Ztg. 21, 171 (1909).

CHAPTER IX

SOLUBILITY PRODUCT PRINCIPLE

THEORY. ADSORPTION. RÔLE AND REGULATION OF HYDROGEN ION CONCENTRATION

141. The solubility product principle is a relationship which is at the basis of all gravimetric procedure in that it generalizes the behavior of difficultly-soluble salts1 in their saturated solutions.2 The importance of the principle lies in the fact that whenever we are concerned in bringing about precipitation or in preventing it, or in effecting solution, we are dealing with the very conditions of which this principle treats.

Rigorously stated, the principle is as follows: in a saturated solution of a difficultly-soluble salt, the product of the concentrations of the constituent ions for any given temperature is sensibly a constant, each concentration being raised to a power equal to the relative number of ions furnished by one mole of the salt upon dissociating.

3

Thus for the general case of the difficultly-soluble salt AmBn we would have as the relative number of ions which it would furnish upon dissociating, m cathions and n anions according to the scheme:

(CA+...) × (CB-...)”

=

Solubility Product AmB2

142. As examples of the foregoing we give herewith a list of the solubility products of the more common difficultly-soluble salts met with in gravimetric analysis:

1 By the term difficultly-soluble, we shall mean not that the salt is difficult to get into solution but that the salt is only very slightly soluble. The term is rather a loose one and there is really no exact line of division as to which salts are to be considered difficultly-soluble and which are not, but broadly speaking most of the salts considered under this category have a solubility less than 50 milligrams per liter.

2 A saturated solution of a substance is one in which the solid phase of the salt is in equilibrium with the solution.

3 The solubility product for any given substance varies slightly with temperature. For temperatures between 15° and 35°, however, the variation is so slight that for most difficultlysoluble salts we can ordinarily regard the value of the solubility product as independent of the temperature.