A DEMONSTRATION OF CHANGES IN IONIC EQUILIBRIA DURING CHEMICAL REACTIONS
Vivid, interesting experiments, presenting some of the fundamental concepts of the ionization theory for beginning students, are described in this paper. A striking demonstration is given of (1) the removal of ions by the formation of the weak electrolyte, water, in the reaction H+
+ -el+
Na+
+ -OH +HzO + Na+ + -CIS
(2) the removal of ions by the formation of an insoluble precipitate in the reaction Ba++
+ 2-OH + 2H+ + -SO4 -+BaSOl + 2H20,
(3) the formation of ions by the reaction of slightly ionized substances to form a highly ionized substance in the reaction NH40H
+ HAc +NHI+ + -Ac + H20.
A recent article by McClelland (1)describes a method of demonstrating one of the reactions given above. The author (2) has made use of the reactions described in the following paper but with a much less satisfactory apparatus and technic. Iron Electrodes
--
No. 10 Rubber StopperNo. 1 Rubber Stopper 150-Cc. Add SolutionRubber Disk
--40-Watt Lamp
150-Cc. Basic Solution.-with Phenolphthalein
110 VoltA:C. Circuit
The accompanying drawing indicates the apparatus used in all three of these experiments. The electrodes are iron rods cut from ordinary ring stand supports. Any reaction between the iron rods and the acids is not noticeable since the solutions are dilute. The carbon rods formerly used gave trouble by breaking. The rubber disk a t the center of the cylinder is a section cut from a No. 10 rubber stopper. The No. 10 stopper a t the top of the cylmder and the rubber disk both have holes for the passage of the electrodes and a hole in the center which is closed by a No. 1 stopper. 2722
VOL.7, No. 11
CHANGES IN IONIC J3QUILIBRIA
2723
The No. 1 stopper is inserted in the disk from below and in the No. 10 stopper from above. The first experiment offers support for the assumption that the reaction between sodium hydroxide solution and hydrochloric acid solution reduces the total number of ions by half. This reaction as written in the first paragraph expresses that idea. Tenth normal solutions of hydrochloric acid and sodium hydroxide are titrated, and then diluted to solutionsof approximately 0.0006 N concentration, being made as nearly equivalent as possible by the usual quantitative procedure. The electrode assembly is removed from the cylinder and 150 cc. of the basic solution containing phenolphthalein are introduced from a pipet. The rubber disk is moistened with water to make it slide easily, and the electrode assembly inserted in the cylinder. The switch is now closed and the attention of the students directed to the intensity of the glow of the lamp. This glow is due to the conductivity of the sodium and hydroxide ions alone. The hydrochloric acid solution is now added and is supported by the rubber disk so that none gets down into the sodium hydroxide. The increased intensity of the glow of the lamp should be noted at this point. Now the solutions may be mixed by pushing out the stopper in the disk with a stirring rod, stoppering the opening in the No. 10 stopper, and inverting the cylinder and shaking. After the thorough mixing it will be seen that the lamp gives the same faint glow from the sodium and chlorine ions remaining that it did with the sodium'and hydroxide ions in the original basic solution. The disappearance of thevink color of the phenolphthalein indicates the removal of hydroxide ions. Thus the experiment has furnished evidence for the theory that neutralization of such strong electrolytes is simply a reaction between hydrogen ions and hydroxide ions to produce undissodated water. Since these ions furnish half the condnctance of the solution before mixing it is quite reasonable that the sodium and chloride ions remaining should conduct only as well as the original sodium and hydroxide ions in the basic solntion alone. The second experiment makes use of the same apparatus but substitutes 0.02 N barium hydroxide for the sodium hydroxide and 0.02 N sulfuric acid for the hydrochloric acid solntion. I n this case it should be noted that the light bums with almost full brilliance with the basic solution alone because of the greater concentration of the ions. On mixing, the disappearance of the pink of the phenolphthalein indicates the removal of the hydroxide ions as before. The formation of the cloudy precipitate of insoluble barium sulfate and the complete extinction of all glow from the light indicates the removal of the rest of the ions. Thus it is shown that ionic equilibria may be shifted toward completion by the formation of an insoluble electrolyte as well as by the formation of a slightly ionized substance.
2724
JOURNAL OF CHEMICAL EDUCATION
NOVEMBER, 1930
The last experiment involves the use of 0.003 N ammonium hydroxide for the basic solution and 0.003 N acetic acid for the acid solution. With the basic solution in the cylinder alone or with both the basic solution and the acid solution in the separate compartments; the light should barely show any visible glow. After the solutions are mixed, the light should glow brilliantly. The application to the ionic theory can be made by pointing out that two weakly ionized electrolytes have reacted to form one strongly dissociated electrolyte and water. The form in which the equations are written in the first paragraph is intended to emphasize the theory which the experiments are designed to demonstrate. Equations of this type which indicate the original reactants and the final products are spoken of as the equations for the main reactions by Bray and Latimer (3). In contrast with this form of equation, users of the ionic theory sometimes make use of equations designed to show the mechanism of the reaction. In such cases the ammonium hydroxide and the acetic acid would be shown as ionizing before any reaction could take place. However, there is nothing about this demonstration to indicate any such mechanism. In conclusion the author wishes to acknowledge the kindness of G. Ross Robertson and Leon Robinson in making certain suggestions used in the experimental work described in this paper.
Literature Cited (1) MCCLELUND,"Experiment t o Show the Removal of Ions in Double Decomposition," J. CHEM.EDUC.,7, 1579 (July, 1930). (2) STONE and D ~ N "Experiments . in General Chemistry," McGraw-Hill Book Co., Inc., New Yark City, 1929, pp. 424. (3) BRAYand LATIMER,"A Course in General Chemistry," Macmillan Co.. New York City, 1926, p. 6.
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Persons Drinkhe Lemonade Made in Enamel Pails Get Antimony Poison. Icecold lemonade made in enameled pails bought especially for the purpose caused a severe thoueh of a 5rm a t Newcastle-upon- brief outbreak of illness amone the 70 employees Tyne, England. The actual cause of the illness was antimony poisoning. The pails in which the drink was .re . pared were not made for coaking purposes and the cheap enamel glaze on the inside contained antimony oxide equivalent to 5 per cent of metallic antimony, .. which the acid of the drink dissolved. The drink was made from prepared "fruit crystals" which contained tartaric instead of citric acid, and after preparation was allowed t o stand in the pails overnight, to be iced, sweetened, and served the next day. The drink had been prepared by the firm for the benefit of its employees during the extreme hot weather. The investigators comment an the unfortunate ending of this good intention and state that enamel containers intended for cooking, which the pails in question were not, are given a glaze which is u, adjusted that the ordinary acid solutions used in cooking will only remove inappreciable amounts of antimony, and are thus safe t o use.-Science Sem'ce ~
