V04.4. NO. 2 METHOD OR TEACHING OXIDATION
AND REDUCTION
REACTIONS223
THE POLE REACTION METHOD OF TEACHING OXIDATION AND REDUCTION REACTIONS' H. L.LOCHTE, UNIVERSITY
OP
TEXAS. AUSTIN,TEXAS
In 1910 Professor E. P. Schoch, a t the University of Texas, began teaching oxidation and reduction reactions as electron exchange reactions. From the &st the method has been in use in classes of first-year college chemistry. Both Professor Schoch and his successors made many changes in detail in the method since 1910 but the fundamental ideas have not been changed. A somewhat similar method has been in use a t the University of California for a number of years and, recently, Brinkley2 described a similar method developed a t Yale. The teaching of oxidation and reduction reactions is prefaced by a series of lectures on the periodic system as developed and, especially, as explained by means of modern theories of atomic structure. Only generally accepted ideas of structure are presented, but there is enough agreement in these to permit the use of atomic structure lectures as a valuable introduction to electron exchange reactions of all kinds. This topic is followed by a chapter on electrolysis which naturally leads to consideration of back electromotive force of electrolytic cells or to the electromotive force of battery cells. Although both electrolytic and battery action are oxidation and reduction reactions, oxidation and reduction reactions, as commonly discussed, result only when the constituents of the two battery poles are mixed. Even in the case of electrolytic reactions the equations for all pole reactions are written as encountered, and balanced so as to make the number of electrons lost equal to those gained. The first step in the study of oxidation and reduction reactions involves predicting whether a reaction will take place and knowing what products will result. From data derived during demonstration, a simple electromotive force table is drawn up and explained before the fairly complete e. m. f. table is presented (Table I).3 The student is expected to become thoroughly familiar with this table without actually memorizing it. Each line represents a complete reversible pole or half cell, i. e., the combination may be used as pole of either an electrolytic or battery cell. The left-hand member of each pole has the greater; theright-hand member the lesser number of electrons, so that in changing from the condition on the left to that on the right-hand side the atom or ion loses one or more 'Paper read before the Division of Chemical Education, American Chemical Society, Tulsa, Okla., April 7, 1926. 2, 576 (July, 1925). 'S. R. Brinkley, THISJOURNAL, For the original table see Liddell, "Handbook of Chemical Engineering," McGraw-Hill Book Co., 1922, vol. 11, p. 690.
electrons. The strongest reducing agents then are found in the upper left-handcorner; the strongest oxidizing agentsin thelowerright-hand corner of the table. Whenever the left-hand member of one pole is electrically connected to or mixed with the right-hand member of any pole below it in the table, a battery or oxidation and reduction reaction takes place. If the ions or atoms are both right-hand or left-hand members there can be no electron exchange; if they are so situated that counter clockwise rotation would result, only electrolytic and no battery action can take place. Any electron exchange in this case would have to proceed with
~
changing
Elements in reduced state
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21
Potassium metal Sodium metal Calcium metal Magnesium metal Aluminum metal Hydrogen gas Zinc metal Sulfide ion (sodium sulfide sol. normal in S--) Iron metal Hl ( I atm.) Cadmium metal Sulfide ion (sat. sol. of hydrogen sulfide in pure water) Lead metal Nickel metal Sulfide ion (sat. sol. of HzS in normal HCI) Leadmetal Tin metal Hydrogen gas Bismuth metal (same for Sb) Stannous ion (stannous chloride sol ) Mercury metal
22 Copper metal 23 Iodide ion (normal in I-) 24 Ferrous ion (sol. normal in
Fe++) 25 26 27 28
Silver metal Mercury metal Oxygen ion (in neutral sol.) Bromide ion (sol. normal in Br-)
tendency in volts
Elements in oxidized state
+ 1 (-) + 1 (-) Ca++ salt sol. + 2 (-) Mg++ salt sol. + 2 (-) Al+++ salt sol. + 3 (-1 H+ ions in N/1 NaOH sol. + 1 (-) Zn++ salt sol. + 2 (-) K+ salt sol.
Na+ salt sol.
+
S o element 2 (-) Fe++ salt sol. 2 (-) H + ion (neutral water) Cd++ salt sol. 2 (-)
+ +
+
S o element 2 (-1 Pb++in (sat. sol, of PbS04) Ni++ salt sol. 2 (-)
+
+
+ 2 (-)
S o element 2 (-) N/1 Pb++ salt sol. 2 (-1 Sn+* salt sol. 2 (-) N/1 H+ ion sol. 1 (-1 Bi+++ salt sol. 3 (-) Sn++++ salt sol. 2 (-) (stannic chloride sol.) Hg+ salt sol. 1 (-) (HglCI1 in normal C1- sol.) 2 (-) CuC+salt sol. I"element -k . 1 (-) . . (sat. . sol. of iodine) FeC++ salt sol. 1 (-) (normal in Fe+++) Ag+ s a l t k l . 1 (-) Hg+ salt sol. 1 (-) Oxygen gas (sat. sol.) 2 (-)
+ + +
+
+
+ +
+
+ +
Bro element (sat, sol.)
+
+ 1 (-)
VOL.
4, NO. 2 METHOD O F T E A C ~OXIDAT~ON NG AND REDUCTION REACTIONS225
TABLE I (Concluded) ELECTROMOTIVE FORCE TABLE Elements in reduced state
29 Sulfur in compounds, with valence less than 6 (+) 30 Nitrogen in com~ounds,with valence less than 5 (+) 31 Cri++ compounds 32 Chloride ion (sol. normal in Cl -) 33 Mu++ compounds 34 Cl" element
35 Oxygen ions (in any acid sol. having very few 0-- ions)
Relative changing tendency in volts
Elements in oxidized state
+
-0.80 (?) So+ compounds free (-) (conc. HL3OJ - 1.OO t o -0.72 N 6 +compounds (HNOs) free (-) -0.7 C++ compounds free (-) (sol. of chromic acid)
+
+
+
C1" element (sat. sol.) 1 (-) Mn7+compounds 5 (-) (sol. of permanganate) -0.40 (?) Chlorine in compounds where i t has positive valence +free (-) (NaOCI, KCIOI) -0.30 f?) Oxveen .- -eas liberated from a Dlatinum pole by electrolysis of salts of oxyacids such as nitrates. sulfates. free (-) phosphates -0.34 Ph4+ (from PbOr-solid in dil. H&OJ free (-) -0.10 F"element 1 (-) 0.00 Theoretical zero pole -0.65 -0.49
+
+
36 PbC+ ion (in sat. sol. of PbSO, in dil H d O S 37 Fluoride ion sol. 38 Theoretical zero pole
+
+
Solutions are normal unless stated otherwise.
energy supplied from without. A number of atom-ion combinations are listed in several positions in the table to show the effect of concentration of the ion concerned and the subject of this type of shift in position is stressed during the last few lectures on electron exchange reactions when reactions of nitric and sulfuric acid are considered in detail. The subject of corrosion of iron brings out other cases of such a shift of voltage. Familiarity with the periodic and electromotive force tables enables the student to predict when an electron exchange reaction should take place and what products will be formed. Even when he encounters border line cases he can make an intelligent guess as to the products probably formed. As chemical equations form a very valuable means of concise and accurate expression, a simple, consistent, and logical method of deriving the complex forms sometimes met in oxidation and reduction reactions is highly important and, from the standpoint of the teacher, very desirable. Here, again, the electron is treated as an entity or unit of matter which is as much a part of the oxidation and reduction reaction as the ion is of the metathetical reaction. Even though the student is not expected, in later years, to derive equa-
tions by the following method when he can write them by inspection, the beginning student is required to derive every oxidation and reduction equation used regardless of degree of simplicity. This requirement is insisted upon to make the student thoroughly familiar with the mechanical steps involved before the more difficult equations are encountered in his work. The steps required are: (1) Writing of the negative pole equation, (2) Writing of the positive pole equation, (3) Balancing of above and adding, (4) Adding to both sides of all ions accompanying any ions of step (1) or (2) and of any salifying acid required. (5) Combining of (3) and (4) to form the final complete equation. The following sample derivations illustrate the use of the 5 steps of the derivation: 1. Equation for reaction between metallic zinc and hydrochloric acid.
(5) Zn"
+ 2HC1 G ZnCL + Ha2
2. Equation for reaction between KMnOa with FeSOa in f-12SOasolution.
3. Equation for reaction between aluminum metal and sodium nitrate with excess sodium hydroxide.
VOL. 4. NO. 2
METHOD OR
TEACHING OXIDATION AND I~EDuCTIONREACTIONS
227
This scheme of deriving the final equation by simple and fundamental steps seems, a t first hand, to require too much time and energy, but is well worthwhile as i t eliminates all memorizing of complete equations or even of key equations for each oxidizing agent and i t enables the student to derive the equation for any reaction with which he is sufficiently familiar to predict the probable products formed in the first two steps. Long experience with the method leads us to believe this method is the most satisfactory one in use not only as far as immediate results are concerned, but also because it is based on fundamental principles that are constantly used in advanced work.
