A SEMI-QUANTITATIVE VISUAL METHOD for COMPARING ELECTROLYTIC CONDUCTIVITIES in LECTURE DEMONSTRATIONS* F. E. BROWN AND W. G. BICKFORD Iowa State College, Ames, Iowa
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LMOST every laboratory manual for general chemistry describes an apparatus for comparing electrolytic conductivities. The apparatus usually consists of a dipping electrode in series with an incandescent light bulb and a source of electric current. Our modification of this simple apparatus consists (1) in adding a second circuit in parallel with the single circuit commonly employed, and (2) in using multiple electrodes, electrodes so arranged that the sum of the conductivities of two solutions may be measured, as well as the conductivity of any one solution. The added circuit consists of a variable resistance and an incandescent light bulb, the reference lamp. The reference lamp should be as nearly like the test lamp (the incandescent light bulb used in series with the solutions) as possible. With a 110-volt circuit, 220-volt lamps or ordinary carbon-filament bulbs are quite satisfactory. The maximum resistance of the rheostat should be large enough almost or entirely to prevent the glowing of the reference lamp when the current is applied. The maximum resistance should not be less than 200 ohms, and 500 ohms is desirable. The resistance should be calibrated so that the resistance in series with-the reference lamp may be determined within about 10 ohms. Resistances may be made by winding any wire having a fairly high resistance and melting point on a non-conducting frame. The wire may be obtained from old electric heaters. The multiple electrodes may consist of two pieces of iron or nichrome wire bent into U-shapes and mounted rigidly a t a convenient distance apart on a block of some non-conductor, of metal flags mounted rigidly on insulated wires, or of graphite rods suitably connected. The different pairs of electrodes should be similar, and if they consist of wires or rods, provision must be made for equal depths of immersion in the
electrolytes when conductivities are to be compared. The presence of the second parallel circuit containing
the reference lamp and the calibrated variable resistance enables the demonstrator: (1) to keep before his class a measure of the conductivity of any one solution while that solution is modified in any way or another solution is substituted for it; (2) to recall the brightness of the test lamp under any previous condition under which i t was matched by the reference lamp and the resistance in series with the reference lamp was recorded; (3) to determine approximately the actual resistance between the electrodes as they stand in any particular electrolyte (when the two similar lamps are equally bright the resistances in the two circuits are approximately equal) ; (4) to determine the relative specific conductivities or the relative equivalent conductivities or relative molar conductivities of any two solutions as accurately as the lamps can be matched; and ( 5 ) to determine the absolute specific, absolute equivalent, or absolute molar conducti;ity of any of the * Presented before the Division of Chemical Education at the ~ ~ l u t i o nass accurately as the comparisons are made, Meeting of the A. C.S., held Thirteenth Midwest Regional if the conductivity of one of the solutions is known. Louisville. Kentucky, October 31-November 2, 1935. 384
The apparatus in use in our lecture room consists of two pieces of board and the common electrical devices attached to them. The larger board, A (Figure I), is 60 X 20 X 1.3 cm. (8 X 24 X 0.5 in.) To one of its faces are attached: two electric light sockets, a, in which is mounted the referencelamp, andb, in which is mounted the test lamp; a 400-ohm. variable resistance, c; and two binding posts, d, and e, which are connected by about six feet of drop cord, I, to the source of electrical current, to the reference lamp on one side by insulated wire, 2, and on the other side through the variable resistance by insulated wires, 3 and 4, and to the test lamp on one side by insulated wire, 5, and on the other side through the electrodes mounted on the other board, B, by means of insulated wires 6 and 7. The second board, B, is about 18 X 4 X 0.6 cm. (7 X 1.75 X 0.25 in.). On it are mounted four electrodesf,g, h, and i. The cathodes, f and g, were made by bending a piece of No. 8 B&S gage nichrome wire 30 cm. in length to form three sides of a square each 10 cm.long. The anodes, h and i, were made in exactly the same way. To one angle of the cathode was soldered one end of insulated wire, 7, and to an angle of the anode one end of insulated wire, 6. The middle third of each bent wire was covered by a piece of close-fitting rubber tubing. In board B, 4 cm. from each end and 0.75 em. from each side were bored snugly fitting holes to receive the electrodes. Any movement of the electrodes is further hindered by slipping a one-hole No. 00 rubber stopper on each electrode and pressing i t firmly against the board, B. The solutions to be used in the demonstrations are placed in ordinary 150-cc. lipped Pyrex beakers. When the conductivity of a single prepared solution is to be tested, only one anode and one cathode are used. If a solution is to be diluted, the solution is placed in one beaker and the water to be used for &lution in a second beaker. One pair of electrodes is dipped into each beaker, and the sum of the conductivities of the two separate liquids is recorded by the reference lamp. Then the two are mixed, and the conductivity of the mixture is compared to the sum by comparing the brightness of the test lamp with that of the reference lamp. Using this apparatus and carbon lamps, differences in conductivity can be recognized by large classes: (a) when 50 cc. of molar acetic acid is diluted to 100 cc. of 0.5 molar, (b) when 50 cc. of 0.5 molar ammonium hydroxide is diluted to 100 cc. of 0.25 molar, (c) when 0.02 molar solutions of sodium chloride and potassium chloride are compared (because of greater hydration of the sodium ion, in spite of its smaller mass, the sodium ion is a slower carrier of electricity than the potassium ion), and (d) when the temperature of a 0.02-molar solution potassium chloride is raised 20° within the range between 0' and 80'.
The multiple electrode enables the demonstrator to show: (1) the lack of conductivity simultaneously of the dry salt and pure water before they are mixed, (2) the sum of the conductivities of the water and a solution to be diluted as well as the conductivity of the original solution and the diluted solution, and (3) the sums of the conductivities of separate solutions as well as the conductivities of the mixtures of these solutions. Some specific cases are interesting. (1) When multiple electrodes are dipped simultaneously into 50 cc. of 0.25-molar solutions of each of ammonium bydroxide and acetic acid in separate beakers, the test lamp glows feebly. When the solutions are mixed and the electrodes are again inserted, the lamp becomes brilliant. (2) If 0.02-molar solutions of sodium hydroxide and hydrochloric acid are used, the lamp is about as bright as for the separate solutions of ammonium hydroxide and acetic acid but almost ceases to glow when the solutions are mixed. (3) The sum of the separate conductivities of 0.01 molar solutions of ammonium acetate and hydrochloric acid is greater than their combined conductivity. In the case of ammonium hydroxide and acetic acid, two weak electrolytes are replaced by one strong electrolyte and water. The strong electrolyte, ammonium acetate, furnishes so many more ions than the sum of the ions furnished by the two weak ones, acetic acid and ammonium hydroxide, that much greater conductivity results even though the best conducting of all ions, the hydrogen and the hydroxyl ions, are those lost. In the case of sodium hydroxide and hydrochloric acid, two very strong electrolytes are replaced by one strong electrolyte and water. There is no increase in the number of sodium ions or chloride ions, and almost all of the hydrogen and hydroxyl ions unite to form water. The case of ammonium acetate and hydrochloric acid shows how the presence of ions other than hydroxyl ions can appreciably decrease the concentration of the hydrogen ions. In this case two strong electrolytes are replaced by a strong electrolyte and a weak electrolyte, and the weak electrolyte, acetic acid, forms many molecules and in that way removes many hydrogen ions and acetate ions from solutions. The dimensions of the apparatus mentioned in this paper are merely those we happened to use. The concentrations and volumes were convenient for our apparatus. Any other dimensions within rather wide limits and then necessarily modified volumes, or concentrations, of solutions would do equally well. An ammeter visible to the whole class could be substituted for the two lamps. For small classes the lamps could be compared more exactly by a grease-spot photometer. However, the results will always be approximate and are designed for demonstration to large classes, not for exact work.
