edited by GEORGEL. GILBERT Denisan University Granville, OH 43023
Redox Demonstrations and Descriptive Chemistry: Part 1. Metals SUSMITTEDBY
Charles E. Ophardl Elmhurrt College Elmhurrt. IL 60126
CHECKED BY
Donald R. Paulson Calllornia slate Unlrerslly. Lor Angeles Los Angeles. CA 90032 At a time of increasing interest in the teaching of descriptive chemistry, new approaches and methods may be helpful. One such approach to introduce the elements in various oxidation states is to use a table of reduction potentials to predict redox reactions in conjunction with dramatic chemical demonstrations. This can then provide the vehicle and motivation to use descriptive text references to look up further information such as nomenclature, molecular geometry, properties, and reactivities for various ions and molecules. A wealth of descriptive chemistry is neatly summarized in a reduction potential series and allows the convenient prediction of literally thousands of redox reactions. Fairly extensive listings of oxidation or reduction potential series are found in Huheeyl, Day and Selbin2, or the CRC Handbook3. The spontaneity and products of a redox reaction are predicted by using the following generalization: "A reducing aeent will react with an oxidizine aeent below it in the sera reduction potential i&" This generalization is applGd series with the potentials for half reactions listed in ascending order. A necessary cautionary remark is that this rule predicts only thermodynamic spontaneity and does not address itself to kinetic considerations, possible mechanisms, or overvoltage. Descriptive chemistry of various oxidation states for a particular element depends upon the half reaction potential for a pair of oxidation states. T h e half reaction reduction potential for any two oxidation states is influenced by several factors includine thermodvnamic considerations. pH. and complexation or precipitation equilibria4. T h e descriptive chemistrv for a oarticular oxidation state also deoends upon the stabiiity of ihe state toward water, oxygen from theair, and disoro~ortionation. . . Drpendiny upon the situation, these demonstrations can he used at either the introductory or advanced levels. At the introductory level the demonstration results are used t o introduce or explain various redox principles and descriptive chemistry facts. At the more advanced level, the student should use the demonstration results, the reduction potential series, and a descriptive chemistry textbook to elucidate the equations to represent the reactions.
Experimental Procedure Procedure I (adapted from Chen5).Into a l-L Erlenmeyer flask, pour 50 mL 0.1 M Fe(N03)~.At first, pour a few milliliters of 0.1 M NaZS2O3down the side of the flask and note the formation of an immediate deep purple color, which indicates the formation of an iran(II1) complex. Then quickly add the remaining Na&03 until a total of 70 mL has been added. The color slowly changes to Light yellow in 3-5 min due to a redox reaction. Procedure 2. Next add 5 mL concentrated HC1 to acidify the solution in preparation for the next reaction. Slowly pour 25 mL of 0.01 M KMn04 (purple) dawn the side of the flask which turns colorlessas it makes contact with the iron solution. Next add 15 mL of 0.1 M KSCN to indicate the final iron oxidation state. Ploeedure 3. Slowly add 10 drops of SnClz and watch the deep orange iron complex solution slowly turn to colorless in less than a minute. Finally, quickly add 10 mL saturated HgCB and note the which mav either maduallv initial aooearanee of a white orecioitate .. . . or quickly turn gray. Reduction potentials for iron, tin, and mercury are Fez++ 2eSn2++ 2eSn4++ 2eS4O& + 2eHg~C12+ 2eFe3++ eHgz2++ 2e2Hg2++ 2eMnOn- + 8Ht
-----
Fe Sn SnZC 25~03~2Hg + 2C1-Fez+ 2Hg Hgzz+ + 5e- MuZ++ 4H20
EO = -0.44 V EO=-0.18V EO= +0.15V EO = +0.17v EO= +0.21V EO = +0.77 V EO = +0.78V EO = +0.92 V EO = +IS1 V
Discusslon Procedure I. The first reaction between iron(II1) ions and sodium thiosulfate is the immediate formation of a purple complex, bisthiosulfatoferrate(II1) ion. Fe3++ 2Sz0,2-
-
[Fe(SZO3),]-
The second reaction is the kinetic controlled eradual color change to colorless caused by the reduction of &on(111) ions to iron(I1) ions bv thiosulfate. This reaction follows the rule that the reducingagent (S~OJ?-) must beabove theoxidizing agent (Fe3'). The reaction using half reactions 4 and 6 in 2Fe3++ 2S,03'-
=2FeZt + Sn02-
T h e structures for Sz032-and S4Os2- may he given if desired. Procedure 2. T h e addition of ootassium nermaneanate is done in the manner of a classical permanganate redox titration.Theacidicconditions insure that thecolorless Mn" ion and not M n 0 2 is the reduction product of Mn04-. (The chemistrv of Mn is treated'in a separate demonstration.) The ironi11) ions as a reducing agent &e oxidized back to the iron(II1) state as indicated bv the formation of the deep orange from thiocyanate complex. The reaction using haif reactions 6 and 9 is
Oxldation States of Iron, Tln, and Mercury Chemicals and Equipment 50 mL 0.1 M Fe(NO& 70 mL 0.1 M NazSz03 5 mL 11.5 M HC1 15 mL 0.1 M KSCN 104 mL SnClz(Dissolve 11.2 g SnClr2HzOin 70 mL cone. HCl, filter or decant any insoluble material, and dilute to 100mL with Hs0.) 10 mL saturated HgCB 25 mL 0.01 M KMnO4 Magnetic stirrer (optional)
716
Journal of Chemical Education
'
Huheey. J. E. Inorganic Chemistry, 2nd ed.; Harper and Row: New York, 1978; p 312f. Day, M. C.: Selbin, J. Theoretical Inorganic Chemistry, 2nd ed.; Reinhold: New York. 1969; p 344f. Weast, R. C., Ed. HandbookofChemistryandPhysics;TheChemical Rubber Co.: Cleveland, OH; Section D. Sanderson, R. T. J. Chem. Educ. 1966,43,584. Chen, P. S. J. Chem. Educ. 1970,47, A784,
'
The iron(II1) thiocyanate complex is the classical test for the detection of iron(II1) ions. Procedure 3. The final series of reactions are sometimes used in the preparation of iron samples for titration as the iron(I1) ion. Tin(I1) chloride is used to reduce iron(II1) to iron(I1) which is responsible for the disappearance of the deep orange color. Using half reactions 3 and 6, the following equation is obtained
+
-
+
Sn2+ 2Fe3+ Sn4+ +Fez+ In this situation i t might be worthwhile to note how half reactions 3 and 6 were selected. Half reaction 6 is easy since i t is the only one containing Fe3+, the oxidizing agent. SnZ+ appears in both half reactions 2 and 3. In half reaction 2 it is an oxidizing agent, whereas in half reaction 3 it is a reducing agent which is needed to react with Fe3+,an oxidizing agent. The final reaction and appearance of the solution depends on how much excess tin(I1) is left in the solution. The final reaction is a good example of where a careful inspection of six half reactions gives the proper ones to explain the reaction. Half reactions 2, 3, 5, 6, 7, and 8 should be examined. Since Hg2+ is used in this final reaction and is in its highest oxidation state, it is the oxidizing agent in reaction 8 to produce Hg22+. The question is what does Hg2f react with since all of the following are now present in solution: Sn4+, Fez+, Mn2+, S4Ofi2-,C1-, SCN-, and excess Sn2+. Careful inspection quickly eliminates Sn4+ and S40fiZ-as oxidizing agents. Mn2+ is a reducing agent hut below the Hg2+ half
reaction. I t is possible for Fez+ as a reducing agent in reaction 6 toreact with Hg2+but the solution should turn back to red-orange as soon as any Fe3+ ion is formed as in reactions 10 and 11.This reaction with Fez+and Hg2+may be too slow to be observed. As a consequence SnZ+must he a reducing agent as in half reaction 3 to produce Sn4+ in the reaction,
Perhaps a quick search in a descriptive chemistry text will yield information about the properties of Sn2+, Sn4+, Hg2+, and HgzZ+.Why is mercury(1) written as HgZ2+?The white precipitate in the reaction is Hg2C12. The gradual darkening of the white HgzClz precipitate is due to the formation of Hg metal. Since Hg22+is an intermediate oxidation state, it can behave as both an oxidizing agent in half reaction 7 and a reducing agent in half reaction 8. As soon as Hgz2+is formed, it can also react with Sn2+using reactions 3 and 5 or 7.
A comparison of half reactions 5 and 7 shows how the presence of a precipitating agent can affect the half reaction potential. The Hg2C12 is more stable toward reduction than the simple Hg22+ ion, although Sn2+ is still a sufficiently strong reducing agent to form the mercury metal. The author wirhes to acknowledge the help of Tony Gryzemski, a student, in workina out the ~roceduraldetails for this demonstration.
Volume 64
Number 8
August 1987
717
