W. C. VOSBURGH, PELHAM WILDER, JR., and J. H. SAYLOR Duke University, Durham, North Carolina
QunmmTIvE analysis was osiginally taught as a part of the analytical chemistry course. As it has declined in practical importance, the conventional qualitative analysis rourse has been dropped from some curricula and in others its emphasis has been changed. The main concern of the theoretical part is now the physical chemistry of water solutions. It is recognized that the laboratory work of qualitative analysis give? experience in a large number and variety of reactions which illustrate some important theoretical principles. Over t,en years ago one of the authors devised and tested with a small class a set of experiments that he hoped would bett,er illustrate solution chemistry. The greatest change was to make some of the experiments quantitative. The revised program seemed successful and it was adopted for all students in the first semester of the second year of chemistry. Entitled "Fundamentals of Analytical Chemistry," this course has heen in operation since then; the instructors concerned have found it to he superior to a laboratory nroeram ronsistine whollv of oualitative analvsis.
. ~-
GENERAL ORGANIZATION
Fundamentals of Analyt,ical Chemistry follows one year of general chemistry, in which tmo-thirds of the second semest,erls laboratory work is devoted to qualitative analysis. The course carries four semester hours credit and consist,^ of one lecture, one recitation period, and two three-hour laboratory periods per meek. The lectures t,reat the standard material and prnblems in elementary solution theory and are illustrated where practicable by lecture demonstrations. The recitation sections are devoted t,o informal discussions of the previous lert,ure and the material and problems assigned from the text; written quizzes are given frequently. About 160 st,udents enroll for the course each year. An average class is romposed of 15 per cent chemistry majors, from six to eight per cent majors in other scientific disciplines, and 75 per cent premedical students. LABORATORY PROGRAM
The objective of the laboratory program is to illustrate theories and reactioue of solutions of electrolytes. A number of the experiments are quantitative because it is felt t,hat a11 understanding of t,he quantitative aspect,s of t,hc theories is important,. These experiment,~constit,ut,e all intsoduction to the niet.hods of
' From a paper g~.esentedat-the Southeastern Regional Meeting of the American Chemical Society, Birmingham, ,uahama, October, 1954.
quantitative analysis. Three experiments are of the conventional qualitative analysie type, siuce these procedures and reactions contain many good illustrations of electrolyte theory. The 11 experiments completed during a semester are described below. ( 1 ) Weighing. The student is taught how to use aud care for the analytical balance and to weight to the nearest milligram. Here as in the other experiments, the principle is stressed and precision is sacrificed to speed and convenience. The maximum precision required in any experiment is one part per hundred (one per cent). (2) Preparation and Standardization qf a 0.5 M HCl Solution. A solutiori of approximately 0.5 M HCI is standardized gravimetrieally by weighing silver chloride. The use of a pipet, the technique of weighing a precipit,ate, and the significance of a standard solution are taught. (3) Determination of the Concentration of a n U n known N a H C O Solution. The fundamental concepts of volumetrio analysis are illust,rated by titration of a measured volume of unknown iYaHC08 solution with the standard HCI solut,ion prepared in Experiment 2. The definition of roncentration in terms of both grams per liter and moles per liter is stressed. I n the titration the buret is read to the nearest 0.1 ml. (4) Properties of Electrolytes. This consists of a number of qualitative, test tube experiments, directed toward an understanding of the chemical nature of solutions of electrolytes. ( 5 ) Colorimetric Determination of Copper. This experiment illustrat,es the relationship between the i d o r of a solution and t,he coucent,rat,ion cf the colored solute, in preparation for use of t,he colorimetric method in latcr experiments. The ropper-ion ronrentration in an unknown solution is found by the addition of et,hyleuediamine t,o a measured sample and determination of how much of a known solution treated in the same way is required to match the color. The color comparison is made visually with matched test tubes hut without an instrument. (6) A m m o n i u m S u l j d e Group of Qualitative Analysis.
The analysis of this group is carried out by the conventional srheme for the separat,io~~ of the cations, aud on a se~nimicrosealr. The emphasis is pla~wlon chemical urinci~leb xud s t w i m i ~ mterhniirue rather thau OII the tcwhil~gof an immediately practiral analytical method. The ammonium sulfidegroup is laced before the hydrogen sulfide group in t,he lahoratorv schedule
642
because of the greater experimental difficulties generally experienced by the student in the analysis of the hydrogen sulfide group. (7) Ionization Constants of Picric Acid and Dinitrophenol. In order to help the student to connect the chemical equilibrium equation with the actual behavior of a system in equilibrium, the ionization constants of two weak acids are measured. The acids are chosen so that the fraction ionized can be measured by a simple colorimetric procedure. The student first observes that the addition of a large excess of hydrochloric acid to picric acid destroys the yellow color by suppression of the ionization. I n contrast, sodium hydroxide deve!ops the full color by salt formation. The color is obviously that of the anion. Then solutions containing moderate excesses of strong acid are matched in color by alkaline solutions to which the required amounts of a known picric acid solution are added. The picrate-ion concentrations of the acid solutions are thus determined calorimetrically. The nonionized picric acid can be calculated by difference and the hydronium-ion concentration is set equal to the hydrochloric acid concentration. The ionization constant can then be calculated. The ionization constant of 2,Cdinitrophenol or of 2,6dinitrophenol is determined similarly. Because these are much weaker acids than is picric acid, the hydrochloric acid concentration required to give the common-ion effect is much smaller, and the hydroniumion concentration must he taken as equal to the sum of the hydrochloric and ionized dinitrophenol concentrations. (8) Determination of the Ionization Constant of Acetic Acid by Colorimetric pH Measurements. Thia experiment introduces buffer solutions, pH, indicators,
JOURNAL OF CHEMICAL EDUCATION
and the colorimetric measurement of pH, as well as some new quantitative analysis methods. A sodium hydroxide solution is prepared and standardized against the standard acid prepared in Experiment 2, and also against a 0.1 M potassium acid phthalate solution prepared by weighing the pure solid. Then an acetic acid solution is standardized against the sodium hydroxide solution and some of the latter is diluted to 0.1 M. T n o acetate buffers are prepared from the acetic acid and 0.1 M sodium hydroxide solutions, and a series of phthalate buffers from the 0.1 M phthalate and sodium hydroxide solutions. By means of methyl red indicator, the acetate buffers are compared colorimetrically with the phthalate buffers, of which the p H values are known. The ionization constant of acetic acid can be calculated from the compositions and p H values of the acetate buffers. ( 9 ) Hydrogen Sulfide Group. This is similar to Experiment 6. (10) Solubility Product of Lead Iodide. A further exercise in the application of the principles of chemical equilibrium to a real system is furnished by the precipitation of lead iodide, when solutionsof leadnitrate and potassium iodide of known concentrations are brought together in varying proportione. By a simple experimental method it is found that; the solubility of lead iodide is affected considerably by the presence of either lead nitrate or potassium iodide, but that in spite of variations in the separate ion concentrations the value of the solubility-product constant is approximately the same in all trials. (11) Qualitative Analysis of a n Alloy. A commercial alloy composed primarily of metals of the first three groups of the qualitative analysis scheme is analyzed for members of those three groups only.
