an equilibrium experiment for an introductory laboratory course in

The result obtained may be checked by a measurement of thee. m. f. of the cell. Ag, AgN03, KNOa (sat. sol.), Fe(NO& Fe(NO&, Pt. The temperature coefic...
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The need for a satisfactory equilibrium experiment for class use has been felt for some time. The most suitable reaction appears to be that of silver with ferric ion to give silver ion and ferrous ion. Equilibrium may readily be approached from either side, and the time required i s such that experiments by different classes may be carried out on successive days using the same apparatus. The result obtained may be checked by a measurement of thee. m. f. of the cell. Ag, AgN03, KNOa (sat. sol.), Fe(NO& Fe(NO&, Pt. The temperature coeficient of this cell has been determined so that the results obtained by the two methods may be compared at the same temperature.

. . . . . .

Physical chemistry stresses the importance of equilibrium from a theoretical viewpoint. However, none of the equilibrium experiments developed up to the present time, has fulfilled the requirements imposed a t this laboratory. There is a dearth of reactions whose eqailibrium concentrations are such as to permit accurate analytical determinations, and which come to equilibrium within a reasonable time. The reaction, Ag

+ Fe(NOs)ss=*


+ Fe(NO&,

investigated by Noyes and Brann (I), appears to fulfil the necessary requirements. Popoff, Fleharty, and Hanson ( 2 ) have recently determined the equilibrium of the reaction 2Hg

+ 2Fe(C104).


Hg:(C10~)~ 2Fe(CI04)-;

this reaction, though perhaps more satisfactory because of the absence of interfering reactions, cannot be adapted for class use on account of the length of time required to reach equilibrium. The equilibrium constant was determined both analytically and electrometrically. In the analytical method the end-point was approached from both sides of the reaction. Experimental The experiments were carried out in 125-cc. glass-stoppered bottles sealed with beeswax, which was found more satisfactory than paraffin. In each bottle was placed 2.5 g. of finely divided silver and 100 cc. of ferric nitrate of the desired concentration, three concentrations being used. I n order to obtain equilibrium from the ferrous-silver ion side, one set of bottles was heated to 5 5 T . in a water-bath for one half hour. Then all the bottles were placed in a thermostat which was kept at 24.7 * O.l°C. and This communication is a portion of the thesis submitted by Miss Hyde in partial fulfilment of the requirements for the degree of Bachelor of Science in Chemistry at the University of Ulinois. 2065



Ocrossn, 1931

provided with means of agitation. The bottles were kept in the thermostat for twenty hours. At the end of this time they were taken out and analyzed and the equilibrium constant determined. The silver for the experiment was prepared in the following way: 85 per cent formic acid, to which about one-third of its volume of water had been added, was neutralized with ammonium hydroxide. The ammonium formate was then heated to boiling. A fairly concentrated silver nitrate solution was also heated t o boiling. The ammonium formate was then slowly added in small amounts, great care being necessary because of the violence of the reaction. After the reaction was complete the silver was filtered off, washed thoroughly, and dried by suction.


.o' Ionic strength. Mallinckrodt C.P. Fe(NOa)3.9He0was used in the preparation of the ferric nitrate solutions. The ferric nitrate was dissolved in 0.0125 N nitric acid solution in order to prevent hydrolysis. All of the analyses were carried out volumetrically. Total iron was determined by converting the nitrate into the sulfate and proceeding in the usual manner. The equilibrium mixture was titrated with potassium permanganate solution to determine the ferrous iron and with potassium thiocyanate solution to determine the concentration of the silver ion, the iron present acting as an indicator. Slight acidification with sulfuric acid and heating prevented the obscuring of the end-point due to precipitation of ferric hydroxide. The same sample was used in the determination of

VOL.8, No. 10



both ferrous ion and silver ion concentrations. The ferric iron was determined by difference. Table I shows the change in the equilibrium constant with a change in co?centration. In order to obtain a true equilibrium constant i t is necessary to extrapolate to zero concentration. This may best be done by plotting the equilibrium constant against the ionic strength of the solution and extrapolating to zero ionic strength (3). Figure 1 shows the results. From Figure 1 the extrapolated value of the equilibrium constant is 0.130.

TABLE I Told Fc. m




TololFe, m









,33915 ,125 ,126 ,0335 ,0135 ,1223 -1223 ,0506 ,0256 ,115 ,113 .I027 .00915 ,1255 ,1265 ,0335 .0135 .I223 ,1223 ,0506 .02X ,115 ,114 ,1024

KO,Ka represent the constant from the silver-ferric iron and the silver ionferrous iron sides, respectively. At this university the class was given instructions and placed the silver in the bottles and labeled them during one laboratory period. Twenty hours before the beginning of the succeeding laboratory period the bottles were placed in the thermostat. During the usual laboratory period the titrations were carried out. For the electrometric work the cell, Ag, AgNO1. KNOa (sat.), Fe(NO&, Fe(NO&, Pt

was used. I t was assumed that the saturated potassium nitrate solution reduced the liquid junction potential to a negligible amount. The ferric nitrate solution for this part of the experiment was prepared as in the analytical part. The silver nitrate solution was prepared by direct weighing using a good grade of silver nitrate. The ferrous nitrate solution was prepared by adding a boiling solution of barium nitrate to a boiling solution of ferrous sulfate until the reaction was complete, and then the barium sulfate was rapidly filtered off. The ferrous nitrate was protected from the air by a layer of stanolind. The determination was carried out as follows. The solutions of ferric and ferrous nitrate were mixed in equal amounts and the mixture was placed in a half-cell similar to that used for the calomel half-cell. In another half-cell was placed the silver nitrate in contact with a polished silver electrode. The side arms of both half-cells dipped into a beaker containing a saturated potassium nitrate solution. Popoff and Kunz (4) found, in determining the potential of the ferric-ferrous electrode, that it required some time for an electrode of this type to reach equilibrium. It was found that from about one and a half to two and a half hours after the cell was set up it gave a steady potential and this value was taken as the potential of the cell a t that particular temperature. Table I1 shows a characteristic change of the potential of the cell with time.



TABLE I1 Tima (Mi".)


0 15 60 75 90 105 120 130 150 270 330

Cell set up 0.W55 0.0490 0.0490 0.0495 0.0495 0.0495 0.0495 0.0495 0.0336 0.0320

From the observed value of the electromotive force of the cell, by means of the equation, E - E " = RTIn


+ Ang+,

A PI***

E n for the cell may be obtained. Assuming the activity coefficients given in Lewis and Randall's "Thermodynamics" ( 5 ) , the equilibrium constant may be determined. As is seen from the table the ionic strength of the solution must be 0.1 M or less. The average value of the equilibrium constant obtained electrometrically is 0.114. Popoff's recent Eo value for the ferric-ferrous electrode, together with the E o for the silver-silver ion electrode gives a value of 0.1306 for the equilibrium constant. Determinations of the electromotive force of the cell a t 14', 25', and 31°C. give the average temperature coefficient of the cell as -0.000409 volt per degree. Using this value, the observed electromotive force may be corrected to the temperature a t which the analytical determination was camed out. Conclusions A procedure has been devised whereby a class meeting one three-hour laboratory period per week may perform an equilibrium experiment in two class periods and check the remlts electrometrically in a third. Literature Cited NOYES, A. A,, AND BRANN, "Equilibrium of the Reaction between MetallicSilver and Ferric Nitrate," I. Am. Chem. Soc., 34, 1016-27 (1912). Popom, S., FLEHARTY, AND HANSON, "Oxidation-Reduction Potentials. IV. The Determination from Equilibrium Data. B. Ferric-Ferrous Electrode," ibid.; 53, 1643-51 (1931). LEWISAND RANDALL, "Thermodynamics," McGraw-Hill Book Co., Inc., New York City. 1923, p. 374. POPOFF. S.. AND KUNZ, "Oxidation-Reduction Potentials. I. Ferric-Ferrous ~lectrode,"3. Am. Chem. Soc., 51,382-94 (1929) LEWISAND RANDALL, 106. ~ d .P. , 332.