JOURNAL O F CHEMICAL EDUCATION
THE STANDARDIZATION OF SODIUM HYDROXIDE SOLUTION A Laboratory Experiment i n General Chemistry PHILIP R. FEHLANDT College of Puget Sound, Tacoma, Washington
T m s paper is the description of a simple laboratory experiment for students in the first year course in college chemistry. No originality is claimed for the experiment, but it does not seem to be well known. A survey of a number of the more commonly used manuals in general chemistry fails to disclose any experiment like this one, which the author has used in his own teaching with considerable success. For an elementary introduction to standard bases and acids, the usual laboratory manual directions are either (1)to prepare an acid solution of approximately known value by dilution of concentrated reagent of known value, and then use this approximate value as the "standard value" in a titration experiment against an unknown base; or (2) an acid (or base) is evaluated by titration against a "stock solution of known normality" taken from the reagent shelf. Students want to know how the "known normality" is ultimately determined. The standardization used in this method of the author is simple and direct. In brief, it consists in actually weighing the NaCl formed (after drying, of course) when a definite volume of unknown concentration
NaOH is neutralized with HC1. The unknown base is prepared in large volume by the stockroom force and the student measures out approximately 20 ml. in a graduated cylinder. Whatever volume he takes, he estimates that volume to the tenth milliliter. Previous instruction has been given on: the importance of a clean glass surface, how to take a readmg on a meniscus, and how to drain the graduate for a definite time period. The measured volume is emptied into a previously cleaned, dried, and weighed porcelain evaporating dish. Phenolphthalein indicator is added, and then HC1 until the indicator color is discharged. The volume and concentration of the acid used are immaterial, and no care need be taken to get exact neutralization, since any excess of acid will boil off in the next step of evaporation. The evaporation must be done very carefully to avoid loss by spattering; gentle simmering is advised. Particular care is needed when the residue becomes pasty, because local superheating and the consequent steam explosion can ruin the experiment. It is recommended that when the student thinks the residue is completely dry the heat be increased and the salt residue be thor-
JUNE, 1949 oughly baked. After cooling and weighing, the salt formed is determined by difference. The student then calculates: (1) the number of mols of NaCl in this weight of residue, and from the equation he can see that this number is also the same as (2) the number of mols of NaOH in the original volume of base used; he then calculates (3) the number of mols of NaOH that would have been in a whole liter of the unknown base. Even
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the slowest student seems to grasp the idea that he has then determined the normality of the NaOH solution. It is not claimed that the method is a quantitative one, hut nevertheless it is surprising how close a careful student can come. Perhaps the limiting factor in the precision that can be expected is the estimation of one part in 200 when reading the approximately 20-ml. volume of NaOH taken in the graduated cylinder.
