JOURNAL OF CHEMICAL EDUCATION
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DEMONSTRATIONS USING A DIVIDED PROJECTION CELL WILLIAM B. MELDRUM Haverford College, Haverford, Pennsylvania
Tm
divided projection cell provides a means of demonstration which is unusually effective, particularly in the illustration of electrolytic effects and of oxidation-reduction reactions which take place spontaneously. Its use with large classes is especially to be recommended. The cell1 in question is illustrated in Figure 1. It is made from clear optical glass, with plane-parallel walls, and with a glass partition extending from the top of the cell to within about 5 mm. of the base. A convenient size for this type of demon st ratio^ has the dimensions 50 mm. high by 40 mm. wide and 15 mm. deep. Placed before the condensing lens of the proj&tion lantern a clear image may be obtained on the screen. Unless a rectifying prism is used the image will be inverted but this is not of importance. A few experiments, representative of the many which are suitable, will be briefly described. These will fall into two categories: Electrolyses and oxidation-reduction reactions.
pendent as required by Faraday's law-and that the writing of an over-all equation covering the two effects has little significance. Experiment 1. Prepare a neutral solution of potassium nitrate containing litmus 0r.a "universal" indicator. Insert platinum wire electrodesZand electrolyze with a direct current of suitable magnitude. The evolution of hydrogen and oxygen is observed and the indicator turns to the acidic color in the anode half of the cell and to the alkaline color in the cathode part. The reactions are: Anode:
Cathode :
..
ELECTROLYSES
There are two general advantages'df using the divided &I in such cases. One is that reaction products a t one electrode are definitelyisolated fromthose a t the other. The other is that the student can realize more clearly that the reactions a t the two electrodes are qualitatively independent of each other-though quantietively de-
Ezperiment 3. Put into the cell a solution of ferric chloride containing a few drops of potassium ferricyanide; this solution shdnld show no green-blue color. To remove any ferrous ion that may be present initially a drop of dilute bromine solution may be added to the ferric chloride solution. Insert platinum wire elec-
These divided projection cells were available prior to the war from Frits K6hler, in Leipzig. The Central Scientific Company, of Chicago, is considering their manufacture and invites correspondence from those interested.
*Convenient electrodes, utilisable in all of the experiments described in this paper, may be made by fusing a 20-mm, length of No. 22 platinum wire into a 6-mm. glass tube, putting a little mercury into the tube, and inserting s. copper wire to make contact with external connections.
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SEPTEMBER, 1948
trodes and electrolyze. Around the cathode the greenblue color of ferrous ferricyanide appears, showing that the ferric is being reduced:
The reverse change a t the anode may be shown by electrolyzing a solution of ferrous ammonium sulfate containingafew drops of potassium thiocyanate solution. The appearance of the deep red ferric thiocyanate complexion indicates that the ferrrous is being oxidized to ferric. Experiment 3. Put snfEcient clean mercury into the cell to seal the eau (Fieure 1). Into the anode oart -out sodium chloride solution and into the cathode part put distilled water containing phenolphthalein. Use a platinum wire as anode but on the cathode side put the platinum wire down into the mercuryso that the mercury, becomes the cathode. Upon electrolysis, chlorine is evoli.ed a t the anode. At the cathode the sodium ion discharges, the sodium so formed dissolving in the mercury, forming sodium amalgam. This latter diffuses though the mercury comes in contact with the water in the cathode part of the cell and reacts, liberating hydrogen from the mercury surfacean dforrning, in effect, sodium hydroxide solution, as shown by the red color of the phenolphthalein. The cathode changes may he written: -
A
,
-
&
F i ~ ~ 1. r e The Divided Projartion C.11
table galvanometer of low sensitivity is substituted for the voltmeter in the circuit the deflection of the pointer indicates the passage of electrons from one electrode to the other through the external circuit. The galvanic cell is now operating in closed circuit, oxidation occurring a t one electrode and reduction a t the other. But of greater interest in the teaching of inorganic chemistry is the fact that we have here a means of illustrating the actual mechanism in a n oxidation-reduction reaction. h many cases it is possible to observe not only the evidence of the passage of electrons but also the chemical Na+ + l e xHg NaxHg (Sodium Amalgam) action in each part of the cell. The student is brought to realize clearly that the step-equations used so widely 2 N a z H g 2 x 2 0 -Hat +2NaC 20H2xHg in representing oxidation-reduction reactions-the oxiThis electrolysis illustrates the Castner process for the dation step and the reduction step, respectively-are not manufacture of sodium hydroxide. merely a device for obtaining a balanced over-all equaWhen ammonium chloride is similarly electrolyzed tion but represent actual reactions which may take essentially the same observations are made although place separately and without putting the oxidizing it is preferable to substitute another indicator, ,e. g., agent and the reducing agent into intimate contact. methyl purple, for the less sensitive phenolphthalein. Experiment 4. The gap beheath the partition is The very interesting product, ammonium amalgam, is sealed with salted agar gel. In one compartment put a obtained which may swell and threaten a short circuit. solution of silver nitrate with a platinum wire electrode As is well known, ammonium amalgam is not stable and inserted in it. On the other side put a piece of zinc wire decomposes thus: or sheet dipping into a solution of zinc sulfate. A galvanometer in the circuit indicates the passage of electrons; a descending current of solution from the zinc OXIDATION-REDUCTIONREACTIONS electrode shows that zinc is going into solution; crystals Galvanic cells of various types may be constructed of silver are observed growing on the other electrode. using this projection cell. The gap beneath the par- For these effectswe write the equations: tition is sealed by means of a 1.5 per ceut agar gel containing about 4 per ceut of sodium sulfate or some other suitable electrolyte. This conducts electrolytically just as an aqueous solution does and functions as a salt Experiment 6. The gap is sealed with salted agar gel. bridge. The solution and metallic conductor cornposing one electrode are put in one side of the cell and some I n one compartment put potassium iodide solution and other electrode similarly set up in the other. The in the other, potassium dichromate solution. Insert electromotive force may be shown to the class on a large platinum wire electrodesin the two solutions and connect through the galvanometer. No deflection is observed scale voltmeter. The same arrangement may he used with only a slight and no effect in either solution. Now add dilute snlmodification to illustrate a different, although closely furic acid to the dichromate solution: The brown color related property. If a microammeter or a lecture- of iodine immediately begins to appear along the wire
+
+
-
+
+
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JOURNAL OF CHEMICAL EDUCATION
in the iodide solution; the solution near the wire in the other compartment begins to turn brown; the galvanometer needle is deflected. Thus, we have: 216e
- -+ In
+ c&-- + 14H80+
2e
(1) (2)
2 c r + + + + 21H20
In order to obtain the over-all equation it is necessary to multiply the first equation through by 3 so as to equalize the number of electrons gained by the dichromate ion and lost by the iodide. Thus we have 61-
-
+ Cr&-- + 14HaO+
2Cr+++
+ 31%+ 21H10
Experiment 6 . The reaction between ferrous ion and ~ermaneauateion in dilute sulfuric acid solution is one of the bYest for demonstration purposes. Freshly made ferrous ammonium sulfate solution to which has been added a little potassium thiocyanate solution is put in one compart&ent and potassium permangauate solution in the other. Platinum wire electrodes are iuserted in both. The ealvanometer ~ o i n t e ris slizhtlv " deflected but the deflection becomes markedly greater when dilute sulfuric acid is added t o t h e permanganate solution. The characteristic red color of the ferric thiocyanate complex ion appears around the electrode in the ferrous solution showing that oxidation to ferric is taking place. Here we have
the deflection increases; on adding a very little sulfuric acid to the peroxide solution the deflection decreases; on adding excess sodium hydroxide to the peroxide solution the deflection greatly increases and simultaneously, as is to be expected, the evolution of oxygen on the wire electrode becomes copious. Also it could be noted that the hydrogen peroxide solution, slightly acidic at first, becomes more strongly so as the reaction proceeds, The experimental observations are in accord with the mechanisms indicated by the following equations: H*Oz HOz-
++ H,O * HO1- + H,O+ H*O = on--+ HIOt 0,--
A
-
-
-
0, t
(4 (b) (4
+ 2e
These lead to the over-all oxidation step in the reaction: H102 2Hn0 Onf 2Hp0+ + 2e (1)
+
The reduction step is:
-
5e
+
-
+ MILO,- + 8H,0t
MuC+
+ 12H20
So that the over-all equation will he: 5HLh + 2Mn06- + 6HaOf 502 f 2Mn++
+
(2) .,
+ 14HzO
hi^ type of reaction, which may be demonstrated using the divided projection cell as described, may also be carried out using a tube filled with salted agar gel Fe++ F e + + + + le (1) and forming a salt bridge of any desired length. We FeC++ + 6SCNFe(SCN).--have, for example, in this laboratory, reduced ferric (red) 5e + Mn0.- + 8HaOi Mu+++ 12H20 (2) chloride solution in one beaker by means of stannous chloride solution in another fifteen feet away. Although BY multiplying the first equation through by 5, adding, the effectcannotso easily be made visible to large and striking out what occurs on both sides, we the divided projection cell it affords class as by the over-all equation: some justification for the old term "reaction a t a dis5Fe++ + Mn0,- + 8H,O+- W e + + + MuC+ 12&0 tance." The formation of zinc bromide from zinc and bromine can be effected in this way just as truly as by Experiment 7 . A demonstration of particular inter- immersing the zinc in bromine solution; in either case est because it offers support of a plausible explanation of the products are zinc and bromide ions. the mechanism involved is the reaction between potasStudents have often carried out experiments using sium permanganate and hydrogen peroxide, in which the method and have thereby learned something of the the hydrogen peroxide is oxidized to water and oxygen. mechanisms of various reactions. Theory has been It is one of those demonstrations that will give satisfac- checked by a quantitative 'comparison of substance tory results only if the preparations are carefully made reduced or oxidized and the change in hydrogen ion and a rehearsal run under the conditions set up. The content. The reaction between cupric ion and iodide demonstration is "temperamental" because of the in- ion has been examined and it has been shown that acid stability of hydrogen peroxide and the catalysis.of its is not essential for the reaction nor does it specifically decomposition by so many different things-including catalyze the reaction as some texts aver; sodium chlothe colloidal agar in the gel. ride or some other electrolyte speeds up the reaction In one compartment put potassium permanganate just as well and, under the conditions of our experiment, solution; in the other put a solution of hydrogen per- probably because it increases the electrolytic conducoxide in dilute sodium sulfate, to render it conducting, tance of the solution. floated on a layer of concentrated sodium sulfate soluIn conclusion the writer wishes to acknowledge his tion, which in turn makes contact with the salted agar great indebtedness to the late Dr. Charles W. Moulton, gel bridge. Platinum wire electrodes are used. On of Vassar College, a master demonstrator, who inclosing the circuit the galvanometer needle is slightly troduced to him some of the possibilities of this useful deflected; on adding sulfuric acid to the permanganate demonstration technique.
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+
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