Silver Equilibria via Cell Measurements
Richard W. Ramette
Carleton College Northfield, Minnesota 55057
A quantitative laboratory experiment
Textbook discussions to the contrary, there are few metal electrode-metal ion systems which readily conform to the expectations of thermodynamic^. and the Kernst equation. Bockrisl has recently discussed the frustrations inherent in attempts to rely on classical electrochemistry when dealing with cell processes influenced by the widely-unappreciated phenomenon of overpotential. Nevertheless the classical galvanic cell can serve very well as a teaching device and in a few instances the measurement of potentials is quite reliable as a means for determining metalligand equilibrium constants. A case in point is the tremendously effective use of the mercury electrode for the study of metal-chelate equilibria,%and many teachers have used the silver metal electrode for potentiometric titrations of halides and for simple cell studies of silver equilibria. The following experiment has worked well in my sophomore quantitative chemistry course as well as in summer institutes for high school teachers, and is an example of how to get a fair amount of instructive information out of a short laboratory period using a silver electrode. The purpose of this experiment is to determine the equilibrium constants for the chloride, ammonia, and bromide complexes, and to deduce the formula for the ammonia complex. The fundamental basis for the experiment is both direct and simple: the observed potential of a silver wire electrode is used to calculate the activity of silver ion which is in equilibrium with a series of solutions having known concentrations of the ligands to be studied. The galvanic cell set-up is shown in Vigure 1. At all time3 the millivoltmeter registerr the potential of the silver electrode with re,?pect to the particular reference electrode used. This potenti:d depends on the activitv of the silver ion in the nolution. and the
' B o c s n ~ s.I., O M . , J. CHFX.Eonc., 48, 352 (1971).
S ~ F ~ M1:.I D W., 5513 (10.56). a
*NI)
I ~ ~ . : I L L I : YC. , N.,
J.
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silver ion activity will be governed by the extent to which the silver is complexed ("removed") by the ligands which can be added to the solution in known quantities. According to the Nernst equation the voltage of the galvanic cell used in this experiment is E
=
E,@'
1 + Ej - 0.0592 log lAg+lf~.
where Eredois the standard reduction potential for the silver ion-silver metal couple compared to the potential of the reference electrode and E, is the unknown but small liquid junction potential. [Ag+]is the equilibrium molarity of silver ion in contact with the silver metal electrode, and f*. is the activity coefficient of the silver ion a t the particular ionic strength used. Initial Conditions
To a 100-ml volumetric flask, pipet 10 ml of a stock solution of 0.00100 M silver nitrate, 0.100 M ammonium nitrate. Dilute
Figure 1. Apporatvs sel-up for cell. A, p H meter used on millivolt scale. 8, reference electrode 0 s supplied with p H meter. C, solution of 0.1 M KNOs. D, agar-KNOa salt bridge to prevent chloride contomination of the silver half-cell. The agar gel i. prepared by water-bath heating of 3 a of mwdered ..or with 1 0 0 ml of roturoted KNOa until the m a r is fully dispemed. silver ion solution in which the e q h b r i o are &loblirhed. F, silver wire electrode (available from Sargent-Welch Co.). The silver wires should be cleaned b y touching them to o piece of oluminum metol immersed in boiling sodium bicarbonate solution. G, mognetic stirrer.
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