ELECTROMOTIVE FORCE TABLE and OXIDATIONREDUCTION REACTIONS H. L. LOCHTE University of Texas, Austin, Texas
This faper presents a non-mathematical plan of using theelectromotiweforcetableasa guidein teaching oxidationreduction reactions. Its use in predicting whether substances should react and what products should form in either electrolytic or spontaneous oxidation-reduction reactions is stressed.
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HE electron-transfer method of writing oxidationreduction reactions has become more or less general during the last decade. Neither it nor any other method of deriving these reactions makes possible the correct writing of equations unless the products formed are known. The two all-important questions that confront the student are: 1. Will these substances react? 2. What products will be formed?
If he can answer these he can use any one of a number of schemes of deriving equations and obtain correct results. While simplified forms of electromotive force or displacement tables have long been used in most texts on general chemistry and while advanced courses quite generally make use of the Nernst formula and some complete form of table, the author knows of only a few schools in which, in first-year chemistry, continuous and consistent use is made of electromotive force tables that show the complete half-cells and their potentials a t specified concentrations. The intelligent use of such a table is of such great value and can so readily be taught that it seems worth while to point out the uses made of this table during the last fifteen years in first-year classes a t the University of Texas. The table used is that given by Schoch and Felsing.' This table is a slight modification of the one given in the "Handbook of Chemical Engineering" by Lidde112and is less complete than that given by Latimer and Hildebrand.3 For convenience in following the discussion given below, a few typical half-cells with their voltages against an arbitrary zero pole are listed here.
Scnocn AND FBLSING, "General chemistry," Von Bwckmann-Jones, Austin, Texas, 1931, p. 344. LIDDELL, "Handbook of chemical engineering." MeGrawHill Book Co.. New York City, 1922, p. 680. S L AND HILDEBRAND, ~ ~ "Reference ~ book ~ of inorganic ~ chemistry," The Macmillan Co., New York City, 1929, p. 367.
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TABLE I (1) Hao (pas) (1 atmosphere) ( 2 ) ZnD (metal)
3.828
+
2 H+( M in OH-) 2 (r) 2 (4 S' (element) -C 2 id
+
3.758 3.51
Zot+ ( M )
3.07
so(element)
2.26
Feit+( M )
++(d2 (4
The following facts about the table are stressed in lectures and quizzes: 1. Every complete horizontal line constitutes a halfcell whose potential against an arbitrary zero is given by the voltage shown. 2. In any horizontal line the left-hand member is richer in electrons than the right-hand one. 3. Any two of these when properly connected yield battery with a voltage equal to the difference between the corresponding voltages. 4. Any pair of these half-cells can be used in an electrolytic cell provided a potential greater than the difference between the two voltages is applied. 5 . The chemical action obtained in both of the preceding cases constitutes an oxidation-reduction reaction. 6. The reaction taking place in the battery takes place equally well in a test-tube if the proper ingredients are mixed. 7. Battery action or spontaneous oxidation-reduction takes place when the left-hand member of an upper and the right-hand member of a lower half-cell are properly connected or mixed. Any spontaneous oxidation-reduction reaction thus involves clockwise rotation of members of half-cells if they are arranged as in the table. 8. If the right-hand member of an upper and the left-hand member of a lower half-cell are present without the other members no spontaneous reaction takes place but the two can be used in an electrolytic cell if sufficient potential is applied; e. g., Zn++ and Fe0. 9. Two right-hand or two left-hand members of halfcells do not react; e. g., ZnDand Fe" or Zn++ and Fef+. 10. If more than two complete half-cells are present battery action will take place first between those farthest apart in the table; i . e., those yielding ~ the greatest voltage. If one of the half-cells
a
is exhausted the potential will drop to that obtainable with the next pair of half-cells and the reaction continues between this new pair. 11. If more than two complete half-cells are presenr. in electrolytic cells that reaction will take place first which requires the least applied voltage. 12. A decrease in concentration of a right-hand member shifts the voltage of that half-cell upward by 60 mv. per electron change per tenfold deaease in coucentratiou; e . g., No. 1and No. 5. 13. A decrease in concentration of a left-hand member of a half-cell shifts its potential downward by 60 mv. per electron change per tenfold decrease in concentration; e. g., No. 3 and No. 6. 14. Since formation of an insoluble compound or of a complex ion may produce a large decrease in coucentratiou of ions involved, the formation of such compounds may result in placing the half-cell containing this ion species far from a related one listed in the table. Such shifts must be considered in making predictions in regard to reactions. 15. A complex half-cell yielding a series of compounds like those obtained with concentrated sulfuric acid and with nitric acid will, in general, yield the same products with all half-cells of approximately equal potentials. If the products obtained with a few representative half-cells are known those to be expected with others in the table can be intelligently predicted. The student who is familiar with these rules and has been shown instances of their application by reference to the table can in nearly all cases predict whether a reaction will take place and what products will be formed. In doubtful cases he can a t least make an
intelligent guess. The advanced student can, if he has complete data available, calculate the exact potential and extent of reaction by substituting in the Nernst formula. Unfortunately complete data are rarely available and in such cases the mathematical formula can yield no more information than the table. Many persons to whom a mathematical formula is bewildering have no difficulty in using the table intelligently, and even the well-prepared student very often fails to get as much information out of a formula as out of a table. The fact that in some cases reactions that are predicted from the table fail to take place because their rate is too slow is pointed out as they are met in lecture, but this complication is not stressed in the discussion of the table because it tends to confuse the student at that stage. It is self-evident that much and varied drilling and quizzing rather than formal lecturing is required if the average student is to master a table containing as much and as varied information as this one. But the results in terms of systematic chemistry learned, in self-reliance, and in enthusiasm of students are well worth the time spent on the electromotive force table. For years the grades of freshmen at the University of Texas have been higher during the second semester whiie studying oxidation-reduction reactions than during the first semester while metathetical reactions are discussed. The highest grades of the year are made in quizzes dealing with the use of the electromotive force table and the derivation of oxidation-reduction reactions. For these reasons the author feels that the consistent use of a complete table of this type is of fundamental importance in teaching general and analytical chemistry.
