Richard W. Zuehlke Lawrence College Appleton, Wisconsin
Laboratory Group Exercises in Acid-Base Theory
O n e topic which is universally taught in the introductory chemistry course is the chemistry and equilibrium of acids and bases. The two experiments here described have helped to integrate the lecture and laboratory approaches to this topic. A definite effort was made to modify the experiments in such a way as to concentrate students' attention directly on the basic principles, rather than on a manifestation of them. The first involves the determination of the pH transition interval of some indicators, treated as simple acids and bases. A student is issued small quantities of one or two indicators in the solid form, identified by the organic name for the compound. He is then asked to find a suitable solvent for the indicator (those available are water, ethanol, and dilute NaOH solution), prepare a stock solution of the indicator, determine the transition interval by dropping some of the indicator solution into solutions of "known" pH previously prepared by him by dilution of 0.1 M HC1 or NaOH, and find the common name of the indicator. On comparing the results of a whole class, one finds that experimental transition intervals located a t the extremes of the pH scale compare favorably with tabulated values (discrepancies amount to those caused by activity effects). However, in the mid-pH range, quite large discrepancies are often noted. Consistent repetition of this pattern with 16 different indicators soon convinces a student that the discrepancy here is due not to a variability of response of the indicator, but rather by his lack of knowledge of the true pH of the solution. This leads into a discussion of the need for and the effect of buffer solutions from a practical point of view. Somewhat aside from the principles of acid-base chemistry, the student finds that he has had some thorough practice in pH-[H30+] conversion exercises as he tries to narrow down the transition interval to nonintegral values of pH. Occasionally, a student receives an indicator such as thymolsulfouephthalein (thymol blue) and gets a first-hand look a t the role of the solvent in acid-base chemistry when he sees the different colors of stock solutions prepared in alcohol and in water. Also with this indicator, the nature of dibasic acids is clearly seen. After the student has become adept at handling the pH concept by using and understanding indicators, he and a small group of other students prepare a series of indicator standards by the procedure of Malm and F r a n t ~ . ' . ~Students as individuals are then assigned several increments of 0.01 M NaOH solution ranging from zero to 50.00 ml to be added to 25.00-ml portions of 0.01 M HC1 or acetic acid. When these two solutions have been mixed, the student takes portions of the mixture, adds appropriate indicator solutions to them, 354
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and compares the resulting solutions with the indicator standards previously prepared. Having thus determined the pH at several stages of an acid-base titration, the student secures several map tacks and indicates his results on an appropriate laboratory chart shoving "pH" versus "ml of added base." The results of this experiment are s h o ~ nin Figures 1 and 2.
Figure 1. C l o a dat.: 0.01 M NoOH.
"tiIralion curve" of 25.00 rnl of 0.01 M HCI with
Figure 2. Class data; "titrotion curve" of with 0.01 M NOOH.
25.00 rnl
of
0.01 M HCnHaO*
The experimental curves correspond very closely to the typical titration curves for the two acids. I t is Presented at the 140th Meeting of the ACS, Chicago, Illinois, September, 1961. 1 The best indicators found by experience are methyl violet, thymol blue, bromophenol blue, bromcresol green, methyl red, bromothymol blue, thymolphthalein, indigo carmine, and 2,4,6 trinitrotoluene. That this selection might be improved is evidenced by the unusual scatter of the points in several areas of the titration curves-much of this is caused by the use of indicators whwe color changes lack strong contrast. 2 MALM,L. E., AND FRANTZ, H. W., "College Chemistry in the Laboratory, 2," W. H. Freeman & Co., San Francisco, 1959, Erp. 34.
gratifying to note that even the slight sharp rise at the start of the acetic acid titration curve is barely detectable; this feature can be brought out more clearly by assigning more small volume increments of the base, and suggesting that the indicator standards be prepared in 0.2 or 0.5-pH unit intervals rather than in intervals of integral values of pH (see Fig. 3 ) .
Figure 3. Clots doto: with 0.01 M NoOH.
"titration curve" of 25.00 rnl of 0.01 M HC8HaOr
In the figures shown, one notes a relative scarcity of points in the vertical region of the titration curves. This is caused by a combination of the nature of the very sharp change in pH around the equivalence point and the lack of buffering in the indicator standards when prepared as indicated above. The latter difficulty can be remedied by explaining the need for buffer solutions after the first experiment is completed, and incorporating buffer solutions in the indicator standards. If the
instructor prefers to discuss buffer solutions in regard to the titration curves, he can have the stockroom prepare the indicator standards ahead of time, using the buffer solutions of Clark and Lubs3 or commercially available solutions.4 The above experiments have been used by the author as described, first to introduce the topic of acid-base titrimetry, and then as an illustration of the properties of acids and bases; students are asked to elaborate on the latter aspect through preparation of a written report. Several improvements and variations on these experiments are immediately obvious. With a little ingenuity, one can alter the experiment on indicators to illustrate acid-base chemistry in non-aqueous solvents. Also, one can use the indicators to illustrate some problems in molecular structure, such as the matters of resonance and the effect of molecular structure on p K values. The application of the overhead projector to further instruction in the use and theory of indicators is apparel~t.~The indicator method of Stedman6 may also offer an avenue for effective modification. The use of newly available, low cost pH meters in this experiment would seem to defeat the whole purpose of providing familiarization with acidbase theory through the use of indicators. The author would like to acknowledge the help of the Photography Laboratory of the Kimberly-Clark Corporation in reproducing the figures shown in this paper. 'See taMe in KOLTROFF,I. M.,AND LAITENEN,It A,, "pH and Electro Titrations," 2nd ed., John Wiley & Sons, Inc., New York, 1941, p. 34. ' Obtainable from Hatmsnn-Leddon Ca., Philadelphia, Pa. SPIEGLER, K. S., ET AL., J. CHEM. EDUC.,39, 87 (1962). ' STEDMAN, D. F., J . CHEM.EDUC., 35, 456 (1958).
Magnetic Stirring Promotes Smooth Boiling Use of a seamless Teflon-covered magnetic stirring bar in conjunction with a magnetic stirrer effectively prevents violent bumping and has a tendency to deorease the foaming prevalent in so many reactions. A particular advantage of the Teflon bar is that vacuum distillations can be conducted without having to introduce extraneous gases. Distillations can be oonducted a t much lower pressures with consequent deorease in decomposition of thenndly sensitive campounda. Another (though minor) advantage is that ordinary still head8 and one-neck Basks o m be used. Stimng bars are effective when either oil baths or glass fabric heating mantles are used. They stand rough usage and are inert to almost all reagents and extremes of temperature. Magnetic stirring is also recommended in oontinuous liquid-liquid extractions-both in the extractor, t o insure better contact between the two liquids, and in the pot, to prevent bumping and discourage foaming. The bumping or foaming that may occur during a Soxhlet extraction, where the effectiveness of ordinary ebullators decreases as the extractina- solution becomes concentrated, can he prevented by magnetic stirring. Another advantage of the magnetic stirrer is that its effectiveness is not marred by a temporary increase in internal pregsure which farces liquid into capillaries and sometimes clogs them. In consequence, the temperature can fall without the danger that subsequent reheating will cause bumping.
JORDAN J. BLOOMFIELD UNIVERS~TY OF ARIZONA, TUCSON Volume 39, Number 7, July 1962
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