Verification of the Form James S. Evans Lawrence U n ~ v e r s ~ t y
Appleton, W~scons~n 54911
of the Nernst Equation An experiment for introductory chemistry
The use of the Nernst equation in introductory chemistry courses is becoming more widespread, if one judges by the trend of recently published textbooks. Evidently a parallel development of laboratory experiments has not yet taken place. It occurred to the author that students might gain a better understanding of electrochemical cells if they could demonstrate for themselves the empirical validity of the Nernst equation. Several laboratory manuals (1-4) contain semiquantitative demonstrations of the concentration dependence of electrode potentials. Typically, the potential of a particular metal-metal ion electrode is first measured against a reference electrode. Then the potential is remeasured after the free metal ion concentration has been drastically reduced by formation of a complex ion, or a precipitate, or both. The student is seldom encouraged to make any use of the emf data that he has obtained. Only one laboratory manual (6) on this author's shelf presents a quantitative verification of the Nernst equation in a manner suitable for the introductory chemistry course. The potential of a copper-cupric ion electrode is measured a t four known values of cupric ion concentration (in the range 0.000,54.5M) using a zinc-zinc ion refcrence electrode. The student is directed to plot emf against the logarithm of the cupric ion concentration. He is then asked how much, on the average, the potential of the copper electrode changes for each tenfold change in cupric ion concentration. 0f.course there is no such dearth of electrochemical cell experiments that are suitable for advanced courses. I n fact, a physical chemistry experiment that is somewhat analogous to the one described here has been in print for many years (6). I n the experiment outlined here, groups of students record data for the concentration dependence of the ferrous-ferric half-cell potential at a platinum electrode, using a silver-silver ion reference electrode, a salt bridge, and a vacuum tube voltmeter. This particular system has several appealing advantages. The voltage differences are maximal because the electrode reaction involves a one-electron transfer. Since there are two concentration variables, one or more '%ctitious Nernst equations" can be investigated and eliminated in favor of the correct one. The reversal of the sign of the potential differencethat occurs during the series of measurements has extraordinary pedagogic value in reinforcing the understanding of sign conventions and t,he concept of spontaneity. The data can be obtained in 30-60 min with simple apparatus and small quantities of reagents. 532
/ Journal of Chemical Education
Experimental Details
The silver reference electrode consists of a piece of silver metal (wire or sheet form) dipping into about 25 ml of 1.0 M AgNO8 solution in a 50-ml beaker. A platinum wire, about 5 em in length and platinized on one end, is dipped into the solution containing ferrous and ferric ions. The electrodes are connected to the voltmeter by insulated leads terminating in alligator clips. The circuit is completed by a salt bridge, consisting of an 11-cm circle of filter paper that has been soaked in saturated I
