inject any water into the bottle until you are ready t o per- form the demonstration.

Students expect the hottle t o "bulge" when the syringe is pressed, since you are adding material to an already "full" hottle. Hold the hottle upside down with the syringe a t the bottom. Press the syringe. The hottle will collapse immedi- ately a s the ammonia dissolves into the water, creating a lower Dressure inside. The deeree of collaose. of course. de- pendLon several factors, iniluding the amoun t of water iniected and the concentration of ammonia in the bottle. If a feadrops of phenolphthalein is added t o the water, the basic pink can be seen in the hottle.

Caut ion 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 ha t comes in contact with the skin should he washed off as quickly as possible. The excess ammonia solution should be flushed down the drain. With reasonable care, this demonstration poses no great hazards.

Redox Demonstrations and Descriptive Chemistry: Part 3. Copper(l1-Copper(l1) Equilibria S U B M ~ E D BY:

Charles E. Ophardt Elmhurn College Elmhurn, IL 60126

CHECKED BY:

George Wollaslon Clarion unlverslty Clarlon, PA 16214

The descriptive chemistries of copper(1) and copper(I1) ions are compared hy observing some rather unusual redox properties involving precipitation and complex ion reac- tions. The redox principles have been explained and illus- trated in two previous demonstrations.l~2

Safety Prepare the solutions containing ammonium hydroxide and

ethylenediamine in a ventilated hood.

1 L 0.01 M copper(I1) chloride 500mL 0.1 M potassium iodide 500 mL 6 M ammonium hydroxide 15mL cone. ammonium hydroxide

250 mL 25% ethylenediamine 20 g potassium chloride 20 g copper thin wire turnings

600mL 0.01 M copper(1) chloride (see below) 1OOmL 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 100 mL 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.

' 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 600-

mL beakers. Safety: Have watch glasses available to cover the bea- kers 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 258ethylene- diamine 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,

~rwedure6 . 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 sur- fare 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 hy- droxide toone and 25% ethylenediamine to the other. In both cases, the precipitate dissolves to farm colored solutions, (A) and (J,K), respectively.

Dlscusslon

The reactions in this demonstration are designed to con- trast and compare the relative stabilities of the copper(1) and co~oer( I1) oxidation states with different anion or li- -~ - - -~

gand enGironments. It is important to note the relative order of the reduction ~o t en t i a l s for the comer ions in an aoueous

& .

environment.

Rsductlon Potentlalsa

Oxidizing Agent6

[Cu(en)*12+ + e- = CUI + 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- =

Reducing Agents

fCu(en)P C" + 1- Cu + 2NH3 Cu + 2en [Cu(NH&l+ + 2NH3 CU + c1- C"+ C" 40H- cu 21- C"C1 cu1

This relative order of reduction potentials where Cuf is a stronger oxidizing agent than Cu2+ is opposite to that found for most other metals. From the table the combination of half reactions 7 and 10 gives the overall equilibrium:

2CuXt - Cu + CuXZC; X = H20, K = lo6

Under these standard aqueous conditions, the equilibri- um favors the formation of cop~er(l1) ions. This equilibrium may he shifted to the left if the concentration of free cop- per(1) ions ia decreased and shifted to the right if the concen- tration of free copper(I1) ions is decreased. By changing the environment around copper by replacing water ligands with others, the position of equilibrium is changed. If X = NH3, then K = 2 X lo-%; if X = en (ethylenediamine), then K = 105.4

Procedures 1 and 2. The first two reactions are designed to show the colors of the co~oer(I1) com~lexes for later com- -. . . parisons.

Cu2+ + 4NH8 - CU(NHJ?~ (blue) (A) tetraamminecopper(I1) ions (14)

Cult + Zen - Cu(en)P (purple) (B) bis(ethylenediamine)eopper(II) ions (15)

Procedure 3. In this reaction, the iodide ion behaves as a reducing agent to act upon Cu2+ ions by combining half reactions 11 and 13 from the table. In addition, the Cu+ ions are immediately precipitated with the excess iodide ions to form CuI. Notice the drastic change in the Cu2+/Cu+ reduc- tion potentials, influenced by the change in the aqueous environment (eq 7) to iodide ions (eq 13).

2Cu2+ + 41- - 2CuI (Dl + I, (C) yellowlwhite yellow precipitate solution (16)

The addition of ammonium hydroxide causes an immedi- ate shift in eauilibrium and a redox reaction. At first dance this reaction may be hard tounderstand hecause thereareno apparent oxidizing agents below the CuI potential (eq 13). The ammonia molecules from the ammonium hydroxide react with free Cut ion in equilibrium with CuI to form the diamminecopper(1) complex ion.

CuI - cut + I-

The diamminecopper(1) ion, in turn acting as a reducing agent, immediately reads with elemental iodine formed in reaction 16 using reduction potentials 5 and 11. The final result is that the precipitate and the yellow color disappear to form the characteristic blue color of the tetraamminecop- per(I1) complex observed in (A).

Procedure 4. This reaction is another example of how the

'Cotton, F. A,; Wilklnson. G. W. Advanced Inorganic Chemishy. 4thed.; Wlley: New York. 1980: p 801.

formation of an insoluble precipitate changes the Cuf/Cu reduction potential. The reaction using half reactions 6 and 12 rather than 7 is:

Cu2+ + 2CI- + CU- 2CuCl (F) white precipitate (19)

Procedure 5. The equilibrium established after several days depends on the presence of Cu, Cu2+, and NH3 in the absence of oxygen. In addition, the relative stability of the ammine complex determines the relative position of the reduction potentials so that the equilibrium favors the Cu(NH&+ complex. The equation using half reactions 3 and 5 is:

Apparently, the linear diamminecopper(1) complex is more stable than the square planar geometry in the cop- per(I1) complex. However, the diammine complex is unsta- ble toward oxidation by oxygen in the air, (during the pour- ingoperation), as shown by half reactions 5 and 9.

(colorless) blue complex (21)

Procedure 6. In the last series of reactions, the starting material is copper(1) as the dichlorocopper(1) complex. The addition of ammonium hydroxide to the copper(1) ions should be compared with the results in reaction 14 and the final equilibrium in reaction 20. The development of a light blue color indicates that only a small amount of Cu2+ is present as CU(NH&~+, while the main species is the color- less Cu(NH3)zf ion (I). The formation of a blue layer on the surface is the same as reaction 21.

Procedure 7. The reaction of ethylenediamine with Cu(1) ions in reaction 22 should be compared with that of reaction 15 and reaction 21. Compare the relative positions of the reduction potentials with ammonia as a ligand (eq 3 and eq 5) with those for ethylenediamine (eq 1 and eq 4).

The reaction of copper(1) with ethylenediamine illustrates the chelate effect where the square planar geometry of cop- per(I1) stabilized by two ethylenediamine is more stable than the tetrahedral geometry that would have to be as- sumed by copper(I), which originally preferred a linear ge- ometry. As indicated by the half reactions l and 4, the Cu(en)zf complex disproportionates to Cu(en)z2+ and Cu metal:

Cu(en),+* Cu(en),+ (J) + Cu (K) + 2 en (22)

Procedure 8. Finally, the last series of reactions provides a second look at some previous reactions. In the firsi reaction, KI precipitates copper(1) as Cul.

Cui +I- - CUI (L) white precipitate (23)

The reaction of ammonia. dissolvine the ~ r e c i ~ i t a t e and producing a blue solution, is a little hard to understand since it Derhaps should have stayed colorless as in reaction 17.

~maliamounts of oxygen or Cu2+ may cause the redox formation of blue Cu(NH3).,2+ (M) as discussed earlier. The addition of ethylenediamine to CuI causes the same dispro- portionation reaction a3 in eq 22.

Volume 68 Number 3 March 1991 249