Electrochemical Measurements in General Chemistry Lab Using a Student-Constructed Ag-AgCI Reference Electrode M. K. Ahn, D. J. Reuland, and K. D. Chadd Indiana State University, Terre Haute, IN 47809
Electrochemistry is an important part of introductory chemistry courses ( I , 2). Thus, it is highly desirable to include experiments that involve the systematic measurements of cell potentials for a number of half cells in the accompanying laboratory. However, the availability of an inexpensive, reliable reference electrode is essential for incorporating these experiments into the freshman laboratory. Commercial standard reference electrodes are expensive and fragile, making them unsuitable for freshman use. Two recent papers ( 3 , 4 )in this Journal describe construction of standard calomel electrodes. Calomel electrodes involve the use of mercury metal and its compounds, and they are too complicated to be made in a general chemistry laboratory In this paper we describe a simple method of making a reproducible and durable Ag l AgCl reference electrode that is used by many of our freshman chemistry students. The student-constructed reference electrode is used to measure
A,
wire
4
a. The reduction potentials of a number of half cells b. The instability constant of Cu(NII8)? c. The solubility product of CuS The cost of the electrode is determined by the Ag wire, which costs about $4.50 per electrode for the wire used in our laboratories. However, the wires can be recycled each semester. After the initial expense during the first semester of use, the electrodes can be prepared for pennies per electrode. A commercial Ag IAgC1 electrode costs between $60 and $130. Experimental Construction of Reference Electrode
One end of a 13-em length of 1.0-mm-diameterpure Ag wire or 90110 AgICu alloy wire is swirled briefly in a 1:l:l mixture of H20:concd HN03:concdHC1 to a depth of about 2-3 cm. The surface becomes uniformly tarnished with a thin layer of AgC1. This electrode is rinsed quickly and thoroughly with distilled water. One end of a 10-cm section of 6-mm-0.d. soda glass tubing is rotated in the hot flame of a laboratory burner until the opening is almost sealed. The other end is fire-polished. Asmall glass wool or cotton plug is pushed to the nearly sealed end with a glass rod. Similarly a small wad of wet filter paper is tightly packed on the glass wool plug. A saturated solution of KC1 is prepared, and 1-2 drops of 0.1 M &NO3 are added. This solution is added to the tubing to a height of about 7-8 an.The tarnished end of the wire is inserted next with the tip slightly above the filter paper. The wire is secured to the glass tubing a t the top with tape over a parafilm seal. This air-tight seal minimizes leakage ofKCI solution through the pinhole. See the figure. If the electrode is immersed in the test solution to the same level as the meniscus of the KC1 solution, any residual leakage would be diffusion-controlled and negligible. 74
Journal of Chemical Education
D agram of stbdenl-conslr.cted Ag AgCl reference eleclroae. The In chness of the AgC waling on the Ag w re s great y exaggeralea.
Cell Potential Measurements Several inexpensive pure metal wires or foils are cut to an appropriate size and polished with emery cloth to serve as electrodes. A 0.10 M solution of each of their salts is prepared. Acell is set up that consists ofthe reference electrode dipped in 100 mL of a salt solution and its appropriate cathode. The cell notation given by Ag I AgCI I Cl- (saturated)lMZi I M
Its cell potential, E, is measured with a high-impedance digital voltmeter (DVM). The anode is connected to the common and the cathode to the dc-voltage input side of the DVM. Approximate Determination of ~ i , t f o r ~ u ( ~ ~ 3 ) 4 ' * and Kspfor CUS
A cell is c o n s t ~ c t e dwith the reference electrode dipped in 100 mL of a 0.010 M solution of Cu2+with the tip of a polished Cu wire immersed as the cathode. The cell potential is measured. For the determination of the instability constant Kin., 10 mL of 6 M NHdaq) are added, and the solution is stirred. Then the cell potential is measured. For the determination of the solubilityproduct constant K,, 2 mL of 2.0 M NaHS solution are added to this solution containing NH3(aq). Then the solution is stirred, and the cell potential is measured.
Calculations The reduction potential of any half cell, M"+I M, can be determined from eq 1. E,d = Ed1
Summary of Standard Reduction Potential Measurements Half Cell
From this, the standard reduction potential for
can be calculated a t room temperature with the Nernst equation.
Eo old
F alloy
F pure Ag
P Lit. Value Average 0.80 f 0.03 V 0.34 f 0.03 V-0.34 f 0.04 V-0.12
alloy
(1)
+ Eref e l d e
AgtlAg
0.78 V
0.80 V
0.81 V
CU'+~CU
0.32 V
0.34 V
0.34 V
~e'+/Fe -
-0.35 V
4.34 V
-
-0.09 V
4.16V
Fe3+lFe
0.80 V 0.34 V 4.44 V 4.04V
To determine Kin.,for CU(NH~)~'+ a t this mH31, the overall reaction equilibrium is assumed to be a s below.
various half cells studied. The literature value for the reduction potential of the references electrode (6)is given by where
E ref eleetmde = + 0.197 V
The reduction half-cell reaction for this reference electrode is and 0.0010
I C ~ N H , )=, ~ ~ ]
Its half-cell potential is constant a s long a s the chloride ion concentration remains a t the solubility value. For the prepared electrodes
and [cu"] is obtained from the Nernst Equation. (4)
To determine K., for CuS CUS(S)
Z? cua(aq) + ~ ~ ( a q )
the formula is KSp= [ C U ~ + I [ S ~ I
(5)
.
where ICu2+1 ~~~~~. . is obtained from ea 4..in which the value o f E used is obtaincd from the solution containing 11% from the S a H S addition The sulfide ion concentration. 3 x 10 l o M. is determined from the equilibrium ~~
Erefelfftmde = 0.17* 0.01 V
determined from a cell that consists of the prepared electrode and a commercial Ag I AgCl reference electrode as the two half cells. The literature value was assumed for the commercial reference electrode. The discrepancy is probably related to the thickness of the AgCl layer. This value of 0.17 V was used throughout for all calculations. In some cases, such as Zn2+I Zn, the nature of the counter ion influences the measured cell potential. The salts used in this study were Also, it is important to prepare fresh solutions for the half cells Fez+I Fe and Fe3+I Fe. The agreement between the measured and the literature
NH3(aq)+ HS-(aq) 2 2 NH4+ + s2-(aq) with Keq =
Kb[NH3]
xK2[H2SI
= 2,0
10-10
K,
-
usine the more recent value of 1.0 x 10"' for K2.(HIS) . (5). In calculating the NH4+concentration, one must consider the amount coming from the reaction + HS7aq) + CuS(s) + NH4+(aq) + 3NH3(aq) Cu(NH3)42f(aq) Results and Discussion Thls experiment takes one 3-h laboratory period and has been used successfullv at Indiana State Universitv for several years both in tge general chemistry laboratory and in the high school honors program. This year, one student was selected to run replicate experiments, and these data are reported. Six electrodes were prepared a s described above-three with a 90110 AgICu alloy wire and three with a pure Ag wire. In addition, an electrode from a previous class, which was a t least two years old, was refilled with saturated KC1 solution and used. The results from all electrodes used were consistent. They are summarized in the Table for the
values shown in the table is generally good with the exception of the M P I Mg and the A13+I Al half cells. For these reactive metals, Mg and Al, there are good reasons for these discrepancies, including the formation of a tenacious oxide coating on their surfaces (7) and the interference from the leveling effect due to the possibility of reducing the solvent, water.
These two half cells are included in the experiment merely to indicate to the student some of the oossible ~roblems. The half cell potentials determined-with the 2-year-old electrode, Eoldalloy, are also in good agreement with the average values for all electrodes. This indicates that this reference electrode is durable as well a s inexoensive and convenient to use. The averaee concentraon of CuZt determined in the ammonia solutik was 2 x 10." M, yielding a n average instaVolume 69 Number l
January 1992
75
bility constant of 1x 10-13 for Cu(NH3).,%. This is in good agreement with the literature value of 1.5 x lo-" (8). The average concentration of Cu2+determined after the sulfide M, yielding a solubility product precipitation was 1x K , of 3.0 x 1e5 for CuS, which is a reasonable value. Conclusion Electrochemistry is an important area of general chemistry. Many aspects of this subject are illustrated in this experiment: reduction potentials, concentration dependence of cell potentials, and aqueous equilibria. The reference electrode prepared by the student is reliable, rugged,
76
Journal of Chemical Education
and inexpensive.Additionally, it has been used by students in an experiment involving an automated potentiometric redox titratiou. Literature Cited 1. MGuarrie, D.A.;Rock,P.A.Oeneml Chemlsty,Znded.: Reeman:NeuYork, 1984. 2. Chang, R. Chemistry, 3rd ed.:Random House: N e r York,1988. 3. Ku8ude.K J. Chem. Educ 1989.66.53. 4. Rsndle. T H.; Kelly. P. J. J Cham.Edm. 1084. 61, 721. 5. Myem Rollie J J. C k m E d u c . 1985,63,687. Chemufy;KolUlaffandEiving,~ds.: 1 Wiiey:N ~ W 6. Tanaka,N.in ~ a t i ~ o n A n o l y t M York, 1963:Vol. 4 , p 2411. 7. Cottm2,FA.; Wilk"aon, G.AduanrodlnagonicChemiioy,5 t h d ; W i l e y : NewYork, 19R8.
