R. H. PeRuccil and P. C. Moews, Jr. Western Reserve University
Cleveland, Ohio
A Simple Quantitative Electrolysis Experiment for First Year Chemistry
M o s t freshman chemistry laboratory courses include work in electrochemistry, usually an experiment in electrode potentials and one in electrolysis. The commonly employed electrolysis experiments tend to be of two types. Some are intended simply to illustrate the nature of electrode reactions through the appearance of electrode products, e.g., Ie which is detected by its color, OH- which is detected with phenolphthalein indicator, and H1and O2 which are collected as gases (1-6). Other experiments are designed as quantitative illustrations of Faraday's laws and usually involve a comparison of the electrochemical change occurring a t an electrode with the corresponding change occurring in a coulometer, these comparisons being made on a weight basis (6-8). Recently a number of electrolysis experiments illustrating the principles of coulometry have been described (9,10). The requirements we have imposed in our selection of an electrolysis experiment for freshman chemistry is that it be simple to perform, that the equipment required be inexpensive (less than $10 per student setup), and that a quantitative application of Faraday's laws should be possible in the experiment, but without the need for analytical balances. The Experiment
The electrolysis employed is that of 0.5 M Na804 solution. Two 400-ml beakem serve as half-cell compartments. To provide maximum contact between electrodes and solution, and to immobilize the electrodes, we have found it advantageous to use metal foil electrodes which are contour-fitted to the bottom and walls of the beakers. To each compartment is added 200 ml of 0.5 M Na2S04. A 10.0-ml sample of an HzS04solution of about 0.001 to 0.01 M is pipetted into the cathode compartment. The acid solution is provided as an "unknown." Four drops of an appropriate indicator are added to the cathode compartment. The electrodes are connected to a milliammeter and two 6-v dry cell batteries connected in series. The circuit is closed by joining the two half-cell compartments with a U-tube filled with agar gel which is impregnated with 0.5 M Na2SOt. The experiment is timed by using either a stop watch or the sweep second hand of a wrist watch or wall clock. The current is measured as a function of time, and the length of time required to produce a color change in the cathode compartment is noted. The end point is best approached by interrupting the current by removing the salt bridge, stirring the solution in the cathode compartment and noting the indicator color. The salt bridge can then be replaced for short
' Preaent address: California State College, San Bemardino, Cdifornia. 552
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Journal of Chemical Education
periods (about 5 sec) until the exact end point is reached. We have found that good results are obtained with the student set-ups if the solution in the cathode compartment is stirred intermittently. Stirring is helpful in locating the end point of the neutralization reaction. However, if a more precise millia~nmeteris used in the experiment the best results are obtained when the solution in the cathode compartment is not stirred until just prior to the end point. This is because of the tendency for stirring to partially depolarize the eleetrodes. The current remains practically constant if the electrodes are polarized during the electrolysis. Details
The Half-Cell Reactions. If the electrodes used in the electrolysis are inert the overall reaction is simply the electrolysis of water, occurring as the two halfreactions, Anode: 2Hz0 Cathode: 2Hn0
-
+
0, 4H+ + &
+ 2e-
+ 4e-
+ 20H-
The OH- produced a t the cathode neutralizes the H+ in the acid sample added to the cathode compartment. A Typical Run. Some typical data obtained in the coulometric analysis of 10.0 ml of an unknown H2S0, solution are presented in Table 1. The total amount of electrical charge involved in the electrolysis is determined by summing together a number of increments. If desired this charge can be determined by plotting current as a function of time and evaluating the area under the curve. We have found the summation of increments to be quite satisfactory. The final result has been rounded off to suggest an accuracy of analysis of about 1%. The accuracy is limited primarily by the accuracy of the milliammeter. Indicator and Indicator Blanks. In 0.5 M Na80, solutions prepared from ordinary distilled water, an appreciable quantity of dissolved COe is present. This means that in an analysis not only must the unknown acid be neutralized in the cathode compartment, but also carbonic acid. There are two ways to make allowances for dissolved COr. In one method an indicator blank is run on 200 ml of 0.5 M NapSol in the cathode compartment. The blank must then be subtracted from the experimental value obtained with the unknown acid (see Table 1). In the second method the 200 ml of 0.5 M N d O P is pre-electrolyzed until the indicator in the cathode compartment assumes its basic color. The unknown acid is then added and electrolysis resumed until this same color change is produced for a second time. Many different indicators have been used in this experiment and a number of them have proved satisfactory. Phenol red, phenolphthalein, and bromthymol blue are all quite suitable. It is instructive to
determine indicator blanks for several indicators and to show how the values obtained can be related to the pH range over which the indicator color change occurs. For example with bromthymol blue (pH 6.&7.6), blank = 1.26 coulombs; phenol red (pH 6.6-8.0), blank = 1.42; phenolphthalein (pH 8.3-10.0); blank = 1.84. Table 1.
Time
x
Time (min) (seo)
Current
10
16
150 X 2 1 . 2 = 3180
{!
22
Total
Blank
30 45
21.3 21.2 21.2 21.9
= 770
n nfifi
i
Remarks Solution yellow
I n ~ ; r u Do&3nt. t stir Solution yellolv Yellow
= 11.666 coulombs = --1.50
10.17 coulombs 1 mole e-
[Hi]= 10.17 coulombs X 96,500 -- X ooulamba
1 mole
OA-
mole e.
r n o l e ~ '
1 mole OH1Hil. determined bv direot titration = 0.0105 M
. 1 = o.o,06
0.01 1
Salt Bridge. The following procedure works well for preparing salt bridges. Soak 12.5 g of USP grade agar in a small volume of cold water. Add to this agar 500 ml of a boiling water solution containing 0.25 moles Na9S04. Carefully pour the agar solution into 6-in. U-tubes maintained in an upright position. Allow the agar to gel. Upon gelation there will be a slight contraction in volume. Add a small volume cf hot agar solution to each arm of the U-tubes to fill them completely. The quantities given here are sufficient to fill about 8 U-tubes. The gels can he preserved for a rather long period of time by inverting the U-tubes in a solution of 0.5 M NazSOa. Choice of Electrodes. The electrodes chosen for the experiment should be inert; a t least they should he such that the only significant cathode reaction is the reduction of water. I n Table 2 are presented the results of the analysis of an HzSOnsolution using four different sets of electrodes. The results with all the electrodes are within the 1-2yo accuracy anticipated with the milliammeter used. It was noted that with copper and silver electrodes some oxidation of the metal to the metal ion did occur a t the anode. The presence of Agf was determined with C1- and that of CuZ+ using PAN indicator (l-(2-pyridyl-azo)-2-napthol). We have used copper Table 2.
In+] found 0.0029 0.0031 0.0032 0.062
- ...
rent (ma)
,120 X 2 1 . 3 = 2560
Tv~icolStudent Results
Student group 1 2 3 4
Typicol Data in the Analysis of An HzS04 Unknowna
15 X 2 1 . 8 = 327
35 X 22
Table 3.
Com~arisonof Electrodes'
electrodes in the student experiment, although from the standpoint of accuracy and cost, nickel electrodes would be equally satisfactory. Choice of Meters. The accuracy with which the experiment can he performed is determined almost exclusively by the quality of the milliammeter used. The results presented in Tables 1 and 2 were ohtained using a S i p s o n 1329C, 0-25 ma, dc meter. This meter, which costs about $15, could he read to the nearest tenth of a milliampere and had an accuracy rated a t 2%. In the student set-ups an EMICO, 0-50 ma, dc meter, which costs less than $2, was used. With this meter, current can he estimated to 0.5 ma, a t best, and the rated accuracy is only 5%. EMICO meters were chosen for the student experiment because of their lower cost. Comnarison of the two different milliammeters indicated that most of the EMICO meters give low readings, by as much as 10%. For use in student experiments one may either calibrate each of the low-cost meters against a more precise meter or apply the correction factor of roughly 10%. Typical Student Results. Table 3 summarizes the results obtained in some typical student experiments. These results are not nearly as accurate as those p r e sented in Tables 1and 2 but almost without exception the inaccuracies are a result of the poorer quality meters. ~~
Acknowledgment
The typical student results presented in this paper were ohtained by the participants in the 1963-64 NSF In-Service Institute for High School Chemistry Teachers a t Western Reserve University. Literature Cited ~ I N B A O., C ~J., CHEM.EDUC.,20,303 (1943). D. R.,J. &EM. EDUC.,25, 495 (1948). MARTIN, TEICEIMAN, L., J. CHEM.EDUC.,34, 291 (1957). TIMM,J. A., AND NEAL,P. E., "Laboratory Exercises in General Chemistry," McGrsw-Hill Book Co., New York, 1956, p. 177. J. F., "A Basic Laboratory Colme in College Chem(5) HAZEL, i s t ~ . "John Wiley and Sons, Inc., New York, 1956, p. 12d. (6) Z n r r ~ ~S., n ,VERNON, A. A,, AND LUDER,W. F.,"A Lab* ratnrv ~" Mxmml of General Chemiatrv." " , W. B. Saunders Co., Philadelphia, 1955, p. 235. (7) GARRETT,A. B., et al., "Chemistry for the Laboratory," 2nd ed., Ginn and Co., New York, 1957, p. 127. M. J., AND PLANE,R.A., "Experirnentd Chemie(8) SIENKO, trv." 2nd ed., McGraw-Hill Book Co., New Yark, 1961, p. 125. (9) I~EILLEY, C. N., J. CHEM. EDUC., 31,543 (1954). H.H., (10) VANLENTE,K. A., VANATTA,R. E., AND WILLARD, J. CHEM. EDUC.,36, 576 (1959). (1) (2) (3) (4)
~
Electrodes
iH+l found f M )
[H+], determined by titration
0.00738
Copper Silver Nickel Platinum Conditions: Two 4 F dry cells (12 v), Simpson l329C, O25 ma meter, phenol red indicator (0.02%).
Volume 41, Number 10, October 1964
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