Miniware for Galvanic Cell Experiments Norman C. Craig, Martin N. Ackermann, and William 9. Renfrow Oberlin College, Oberlin, OH 44074 Currentlv. miniware is beine introduced rapidly into the organic chemistry laboratory. ~ e m i m i c r oquaiitative analysis is also making a comeback in general chemistry laboratories. There is interest in developing other experiments for general chemistry that use miniware. The reasons for doing so are several and compelling. Disposal problems are simplified, and risks to students are minimized. When expensive chemical materials are desirable to make experiments more comprehensive and current, the small amounts needed for use with miniware are affordable. Students mav be encouraged to try their own modifications of an experiment without undue risk. The use of small auuaratus and small amounts of materials is usually a timesaver for students and support staff alike. For some time in our generalchemistry laboratorywe have been using a simple miniware design of a galvanic cell that bas the advantages outlined in the previous paragraph. With such cells and an inexpensive digital multimeter students have been able to explore some simple galvanic cells and go on to measure equilibrium constants for several reactions in a sinele afternoon. As shown in the fieure. .. . the minicell consistsof a small rial and the stems of two medicine droppers. Half-filled with ammonium nitrate, thevialser\~euasthesalt bridge. The dropper tubes contain a few drops of electroactive solutions and small metal wires or strips as electrodes. Cotton plugs are pushed into the bottoms of the dropper tubes to keep the electrode solutions from flowing out into the salt bridge solution. The compact design of this minicell also reduces the contribution of internal cell resistance to the measured voltage, and hence the measured voltage is close to the desired emf. The systems we have included in our experiment and the expected emf s are: (1)CuICu2+(0.1M)IIAg+(O.lM)IAg
second dissociation constant of sulfuric acid. Quinhydrone is a natural, equimolar mixture of quinone (Q) and hydroquinone ($HZ) which gives equimolar concentrations of Q and QH2upon dissolution. The half reaction for this couple and the standard emf are
+
Q(aq) 2H,0t(aq)
+ 2e- = QH,(aq) + 2H,O
Eo = 0.70 V.
The anode compartment contains a buffer solution of equal concentrations of HSOI- and S04Z- and a tiny amount of solid quinhydrone. [H30+] comes from the measured emf and the Nernst equation. Student observations for the emfs for cells 1 through 3 are within 0.02 V of the expected values in almost every case. Most values for the three equilibrium constants are within a power of 10 of the accepted values, which is quite good. A discrepancy of afew hundreths of avolt in the emf translates into such an error in the equilibrium constants because of the logarithmic relationship between concentration and emf in the Nernst equation. For performing this experiment students are supplied with compact kits containing electrode metals, glassware, other materials, and reagents in dropper bottles. Wooden cotton swab sticks are useful for inserting the cotton plugs. Solutions of known concentration for the half cells are pre-
E = 0.43 V
The expected emf s are computed with activity coefficients of unity, an approximation that is used throughout the experiment. As a preliminary experiment students measure emf s of cells 1and 2 and use them t o predict the emf of cell 3. They also check the measured emf of cell 3 against the value computed from the accepted En with the aid of the Nernst equation. Q = [ZnZ+]/[Ag+]Zis not unity for 0.1 M solutions. From cell 4 students determine the equilibrium constant for the dissociation of the tetraamminezinc(I1) complex ion after computing [Zn2+]from the measured emf and the Nernst equation and [NH3] from stoichiometry. From cell 5 they find the solubility product constant for silver chloride after computing [Agi] from the measured emf and the Nernst equation. From cell 6, which uses quinhydrone in the anode compartment and a platinum electrode, they determine the
Assembly of a minigalvanic cell lrom a Zdram viai (17 mm X 60 mm)and two medicine dropper tubes. The viai contains 2 M ammonium nitrate. The medicine dropper tubes contain cotton plugs, wires (or thin strips) ol metal, and eieCtrOa~tive~OIutionS.
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85
pared by adding appropriate numbers of drops. Solution levels in the medicine dropper electrode compartments must be higher than the level of ammonium nitrate in the vial. Otherwise, the flow of ammonium nitrate into the tiny electrode comoartments would chanee concentrations sienificantly. ~ a k i c u l a half r cells can b; put aside and reused in constructine other cells. Of course. reolacement of the cotton plugs a d careful rinsing of the droppers are essential if a half cell is to be remade. Also, care must he exercised to clean the electrode wires with scouring powder if there is any doubt about their condition. Sometimes students mix up an electrode or electrode solution. An untoward consequence can be the plating out of another metal on the surface of an electrode.. ex.. - . c&er .. on zinc. Solutions of coooer sulfate. silver nitrate, zinc sulfate, sodium sulfate, and sodium by: drogen sulfate are all 0.1 M. Ammonium nitrate is 2 M, potassium chloride is 0.2 M, and ammonia water is 7 M. For observine voltaees we use a Radio Shack Micronta Digital ~ u l t i m e i e (cat. r No. 22-191). This inexpensive m u ~ t i ~ e t e r is battery powered and has a liquid crystal display. I t has a 10-Mil input resistance and gives voltage readings to 0.001 V on the 2-V scale. Students must he reminded of the importance of observing the electrical sign of the electrodes as well
86
Journal of Chemical Education
as the magnitude of the voltage. Unsteady voltage readings are usually due to a cotton plug being so tight in an electrode comoartment that a circui
