edited by GEORGEL.
GILBER;
Denison Unikrsily Granville. Ohio 43023
Half Cell Reactions: Do Students Ever See Them? SUBMITTEDBY
Joseph D. Clparlck Norman Thanas Hlgh School 111 Ea.l33rd Street New York. NY 10016 CHECKED BY
Gordon Parker Unlverslly ol Toledo Toledo, OH 43606 The usual demonstrations and labs of electrochemical cells use two different metals in their 1 M solutions. Or else, there is a lab that merely uses the two metals in a dilute electrolyte. A meter indicates the direction of the electronic current. If a more nrecise lab uses a salt bridee or a porous cup, the observatio& are usually limited to what can be seen i i a lab session. The mieration of ions through the salt bridge or cup is seldom seen. If a cell is left to operate overnight, there might be some evidence for the migration of ions through the salt bridge and the reduction and oxidation that takes place in the two half cells. If a Zn/Cu cell is used, the end result after several hours is not very convincing. The Zn electrode does oxidize, but the redudion of the copper is never that obvious. In order to show that there are actually two real half reactions, I have found i t useful to use a FeC131KI cell, which in a lab period (or perhaps a little longer) can show two separate reactions that go to completion. Apparatus Small test tubes, two large test tubes, graphite electrodes, connecting wire, atring or yam (far salt bridge), center-point galvanometer or milliameter. Materlal Iron(II1)chloride (0.1 M),potassium iodide (0.1 M)plus starch solution. (The FeCL3need not be prepared with HCI.) Sodium nitrate solution (dilute) for salt bridge Procedure Have the students mix s few milliliters of the two solutions in a amall test tube. Then set up the two half cells with the graphite electrodes in the large test tubes. The string or yarn can be soaked in a potassium nitrate solution. Have the students connect the galvanometer to determine the current produced. In some cases a milliameter may be more sensitive. The meter is disconnected, and the two electrodes are connected and the half cells left for the remainder of the period. Results and Dlscusslon The first reaction between the FeCL and the KI with starch results in a dark blue solution as the iodide is oxidized to iodine. But the reduction of the Fe" to the lighter FeL' is not evident. As the half cells read, the blue can be seen around the electrode in the KI cell, and the FeC4 cell will get noticeably lighter in color. If left longer, the two half reactions can be
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Metcalfe; Williams; Castka. Exercises andExperimenh in Chem isby: Hoit. Rineharl and Winston: New Yo*. 1970: pp 267-269.
seen in the two test tubes. E = 0 can also be observed when equilibrium is reached. This is a standard lab. I know1. but i t is usuallv used iust to show a meter reading. 1'n the us"al redox reactfons stidents see the end ~ r o d u c t sall mixed toeether and find it difficult to see that there are two distinct products. In the half cells, the ~ r o d u c t are s auite evident. 1tis also a goodway to introduce students to cells that do not involve metal electrodes that are oxidized and reduced. More and more cells are being made that do not involve metal electrodes, but inert electrodes that enable electrons to he exchanged. This is especially true of the hydrogen cell, which is used as the standard for electrode potentials. Metals losing electrons are easier to understand.
Ammonia Bottle SUBMITTED BY
MIchaeI Sheets Arkansas HI^ School Texarkana. AR 75502 Ronald DlStefano Northampton Area Community College Bethlehem. PA 18017 One impressive demonstration of the solubility of a gas in water is the ammonia fountain described by, among others, Shakhashiri' and Summerlin and E n l ~In. ~these demonstrations, a dry, round-bottom flask is charged with ammonia gas, A d &r is added. Ammonia dissolves in the water, lowering the pressure inside of the flask, "pulling" more water un a tuhe from a reservoir (there is a chance of the flask im'ploding if i t is cracked, or if the demonstrator uses a flask of a different desien). I t has been sueeested bv J. M. Manion (Cniversity of central Arkansas, donway, A R ) that a drv flask and ammonia is not neressarv. Addine 15-20 ml. of concentrated aqueous ammonia to-a flask,swirling it about, and pouring the liquid out leaves enough ammonia in the flask to produce an acceptable fountain. If the demonstration is to be repeated, there is no need to dry the flask. Simply add more concentrated ammonia solution and repeat. I suggest that another variation of this demonstration is possible. A plastic, 2-L soft-drink bottle and its cap should be rinsed out and the bottk fitted with a one-hole rubber stopper (#3 or # 4 ) . The stopper has a short piece of glass tubing through it and is connected to a30-mLplastic syringe by a short (2-3 in.) piece of rubber tubing. Fill the syringe with water. Working in a hood or another well-ventilated area, add 15-20 mL of concentrated aqueous ammonia to the soft-drink bottle. Cap the bottle and agitate. Remove the cap carefully-gas pressure can cause some of the ammonia solution to spray out of the top. Quickly pour out the excess ammonia solution and place the stoppersyringe assemhly into the top of the bottle. Be careful not to
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Shakhashiri. B. 2. Chemical Demonstrations; University of Wisconsin: Madison. WS, 1980; Voi. 2. Summerlin, L. R.; Ealy, J. L., Jr. Chemical Demonstrations, A Sourcebook for Teachers; American Chemical Society: 1985. Volume 68 Number 3
March 1991
247
inject a n y water into t h e bottle until you are ready t o perform t h e demonstration. Students expect the hottle t o "bulge" when t h e syringe is pressed, since you are adding material to a n already "full" hottle. Hold t h e hottle upside down with the syringe a t the bottom. Press t h e syringe. T h e hottle will collapse immediately a s t h e ammonia dissolves into t h e water, creating a lower Dressure inside. T h e deeree of collaose. of course. dependLon several factors, iniluding the a m o u n t of water iniected and t h e concentration of ammonia in t h e bottle. If a feadrops of phenolphthalein is added t o t h e water, the basic pink can be seen in the hottle. C a u t i o n should be used in handling the concentrated ammonia solution. Proper eye protection should be worn and good ventilation used i n the preparation. Any of the concentrated ammonia solution t h a t comes in contact with t h e skin should he washed off a s quickly a s possible. The excess ammonia solution should be flushed down the drain. With reasonable care, this demonstration poses n o great hazards.
Redox Demonstrations and Descriptive Chemistry: Part 3. Copper(l1-Copper(l1) Equilibria S U B M ~ EBY: D
Charles E. Ophardt Elmhurn College Elmhurn, IL 60126 CHECKED BY:
George Wollaslon Clarion unlverslty Clarlon, PA 16214 T h e descriptive chemistries of copper(1) and copper(I1) ions are compared hy observing some rather unusual redox properties involving precipitation and complex ion reactions. T h e redox principles have been explained and illustrated in two previous demonstrations.l~2
Safety Prepare the solutions containing ammonium hydroxide and ethylenediamine in a ventilated hood. 1L
500mL 500 mL 15mL 250 mL 20 g 20 g 600mL 1OOmL
0.01 M copper(I1) chloride 0.1 M potassium iodide 6 M ammonium hydroxide cone. ammonium hydroxide 25% ethylenediamine potassium chloride copper thin wire turnings 0.01 M copper(1)chloride (see below) 0.01 M diamminecopper(1)complex (see below)
Preparation of Coppefll) Chloride Dissolve 60 g ammonium chloride in 600 mL distilled water, and boil to expel oxygen. When cool, add 5-10 g copper wire turnings, and dissolve 0.6 g of capper(1) chloride with a minimum of stirring. (It may not all dissolve.)
Preparation of Diamminecoppertl) Complex Do this in o hood. Into a 125 Erlenmeyer flask, put 100mL of 0.01 M copper(I1) chloride and 5-10 g thin copper turnings. Bring the solution to a boil to expel oxygen. Then add 15 mL of concentrated ammonium hydroxide, which will produce the deep blue colored copper complex (A). The solution should turn nearly colorless after about 48 h.
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Ophardt, C. J. Chern. Educ. 1987, 64, 716. 'Ophardt, C. J. Chern. Educ. 1987,64,807. Driscoll, J. A. J. Chem. Educ. 1973, 50, AS9.
248
Journal of Chemical Education
Experimental Procedure Demonstrations Starting with Coppeflll) Ions Pour 200 mL of 0.01 M eopper(I1) chloride into each of four 600mL beakers. Safety: Have watch glasses available to cover the beakers containing ammonium hydroxide and ethylenediamine. Procedure I. Add 100 mL of 6 M ammonium hydroxide into the first beaker to form a blue eopper(I1)complex ion (A). Proeedure2. Into the second beaker, pour 50 mL of 258ethylenediamine to form a purple copper(I1) complex ion (B). Procedure 3. Pour 100 mL of 0.1 M potassium iodide into the third beaker, and observe the formation of a yellow solution (C)and precipitate (Dl from a redox reaction. Next add 100 mL of 6 M ammonium hydroxide, and note the rapid disappearance of the precipitate and the formation of a blue color (E). Procedure 4. Finally, dissolve about 20 g of potassium chloride into the solution in the fourth beaker and then add 5-10 g of copper turnings. Observe theslow formation of a barely visible white cloudy precipitate (F).(Go on to Proeedure 5 while waiting.) Procedure 5. (Adapted from DriscolL3)At least two days prior to the demonstration, make up the diammineeopper(1)complex (G)as indicated above to establish the equilihrium. At the time of the demonstration describe howthe flask was prepared, and refer to the color before equilibrium as the same as solution (A). At the time of the demonstration simply pour the colorless solution into a beaker, . and watch the color gradually change to a deeper blue (H). Demonstrations Starting with Coppefll) Pour 200 mL of copper(1) chloride solution with a minimum of splashing into three 600-mL beakers. Safety: Have watch glasses available to cover the beakers cantainine ammonium hvdroxide and ethvlenediamine, ~ r w e d u r e 6 Into . the first heakeradd 100mL of 6 M ammonium hydroxide. The resulting sdution may be slightly blue \I1 but nut as blue as the previous coppertll, solution with ammonia (A).As the solution stands for awhile, a darker blue oulw develops in the surfare laser due to a reaction with air. P m c p d u r ~7. Add 5u mLof25'bethylenedian1ine to the roppertl) solutiunintherecund beaker.A purplecoloredsolution rJJ that first develops, gradually darkens, add within a few minutes produces a copper metal color on the surface of the beaker (K). Procedure 8. To the third beaker add 100 mL of 0.1 M potassium iodide solution to produce a white precipitate (L). Pour half of the solution into another beaker, and add 50 mL 6 M ammonium hydroxide toone and 25%ethylenediamine to the other. In both cases, the precipitate dissolves to farm colored solutions, (A) and (J,K), respectively. Dlscusslon T h e reactions in this demonstration are designed to contrast and compare the relative stabilities of the copper(1) and c o ~ o e r ( I 1 oxidation ) states with different anion -~ or - - li-~ gand enGironments. It is important to note t h e relative order of the reduction ~ o t e n t i a l for s the c o m e r ions in a n aoueous environment. &.
Rsductlon Potentlalsa Oxidizing Agent6
Reducing Agents
[Cu(en)*12++ e- =
fCu(en)P
CUI
C"
+
8-
=
[Cu(NHs)2]++ e- = [Cu(enkli + e- = [Cu(NHoh12++ e- = CUCI e- =
+ cuz++ e- = cuz++ 2e- = O2 + 2H20+ k-= cut + e- = I, + 2s- = cu2++ CI- + e- = cue+ + I- + e- =
Cu Cu
+ 1-
+ 2NH3 + 2en
[Cu(NH&l+ + 2NH3
+ c1-
CU C"+ C"
40H-
cu
21C"C1
cu1
