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.
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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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