Spot Reaction Experiments Part VII: Chemical Paradoxes FRITZ FEIGL Laboratorio Central d a Produpio Mineral, Ministerio da Agricultura, Rio de Janeiro, Brazil
(Translated by Ralph E. Oesper, Uniuersity of Cincinnati)
A
PARADOX is a completely unexpected contradiction . m . properties, behavior, tenets, definitions, etc., the reverse of those that have gained acceptance by virtue of frequent trials, whose results have been consistently alike in type and sense. Certain compounds exhibit particular properties and reactions, which are so characteristic and typical that the mere possibility of a change in these attributes seems extremely remote. If, then, under certain circumstances, such materials behave in a manner opposed to these expected regularities, a chemical paradox is a t hand. Such surprises occur and can be demonstrated experimentally. However, a rigorous examination of these cases shows that chemical paradoxes seldom present genuine contradictions or real exceptions to the accepted regularities. In most instances a seeming paradox arises because properties and changes in materials are regarded or defined too schematically, and insufficient emphasis is placed on certain factors that seem insignificant but which actually are of prime importance. Consequently, an accurate and critical review often will clear up chemical paradoxes, and provide an acceptable explanation. The contradiction or opposition to the norm then can no longer be maintained. This paper deals with several chemical paradoxes, which, in tmth, are merely apparent rather than real paradoxes. 26. Hydrogen Peroxide as a Reducing Agent (Oxidizing Agents Undergo Mutual Reduction). Hydrogen peroxide furnishes oxygen readily (Hz02 + H 2 0 0). It is a strong oxidant and consequently is often used in classroom discussions as an example of an oxidizing agent. Similarly, potassium permanganate, sodium hypochlorite, and potassium ferricyanide are cited as typical, widely used oxidants. Hydrogen peroxide reacts a t once with solutions of these materials:
+
-
ZKMnOl+ 3H101-2KOH
+ H1O, NaCl + HzO + On + HIOB+ 20H- -2Fe(CN)a' + 2H80 + O1
NaOCl 2Fe(CN)sC
+ 2Mn0. + 2 H 2 0+ 30.
These equations present instances of the mutual reduction of two oxidants. Hydrogen peroxide, in an alkaline solution, reduces auric solutions to free gold: Au+++
+ 3H20. + 60H-
-
6H.O
+ 3 0 +~ 2Au
Logically the three foregoing reactions can also be considered as instances in which the peroxide acts as reductant. This is contrary to its usual oxidizing action, and a t first glance seems paradoxical.
An explanation of this arpphoteric behavior with regard to oxidations and reductions is obtained if it is noted that all oxidative actions of hydrogen peroxide result in its being changed into water. The change HzOz-+ H 2 0 represents a reduction of hydrogen peroxide. This is in agreement with the invariable rule that in every redox reaction the oxidant is reduced and the reductant is oxidized. The foregoing equations, which show hydrogen peroxide seemingly acting as a reducing agent, also show that the peroxide is converted into water. Consequently, the peroxide acts not as a reductant but rather, in the strictest sense, as an oxidant. The position taken here is that every true rednctant must be oxidized. This cannot be the case in the present instances, because no oxidation product of hydrogen peroxide exists. Accordingly, while hydrogen peroxide is not a reductant, nonetheless it undoubtedly brings about reductions. The reason for the reducing action, just as for its oxidizing action, lies in the special way in which one of the oxygen atoms is bound. It is this particularly mobile and consequently easily removable oxygen atom that accomplishes both oxidation and reduction. When peroxide acts as an oxidant, this oxygen atom combines directly witk the reducing agent, or is changed to water; when reductions occur, i t combines with other oxygen atoms that are furnished by the real oxidants. Procedure: All these reactions can be best carried out on a spot plate. (a) One drop of dilute potassium ferricyanide solution is mixed with one drop of very dilute femc chloride, and the mixture is then treated with one drop of hydrogen peroxide. The initial brown color (ferric ferricyanide) changes through green to blue (ferric ferrocyanide). (b) One drop of 0.1 N permanganate is mixed with one drop of dilute sulfuric acid, and the mixture is treated with one drop of peroxide solution. The red color is discharged within a few minutes. If permangauate is mixed with peroxide and allowed to stand in the absence of acid, manganese dioxide precipitates. (c) One drop of gold chloride solution is treated with one drop of an alkaline solution of peroxide. A dark brown precipitate of gold appears a t once. A violet colloidal suspension of gold results if an extremely dilute gold chloride solution is used. 27. Sulfurous Acid Brings about Oxidations. Sulfurous acid is one of the commonest reducing agents. In aqueous solution, i t abstracts oxygen from other compounds and is oxidized to sulfate. However, it is
easy to demonstrate, in a t least two cases, that this acid (or its anhydride) can, directly or indirectly, bring about oxidation. This effect is in such contrast to the usual reducing action that a chemical paradox appears to be involved. A direct oxidizing action occurs when hydrogen sulfide reacts with sulfur dioxide:
not the external effect, but rather the chemism of the induced oxidation of nickelous oxide. Procedure: (a) One drop of sulfurous acid and one drop of saturated hydrogen sulfide water are mixed in a depression of a black spot plate. Sulfur precipitates. (b) Freshly precipitated, well-washed nickelous hydroxide is streaked on filter paper. The paper is laid across the mouth of a small test tube. A few milliHlSOa 2Hb 3H10 35 liters of a water solution of sulfur dioxide are run into the test tube from a pipet. The green nickelous oxide turns gray to black within a few minutes. 28. Nitric Acid Is Not an Oxidant. Nitric acid is The sulfur of both the participants is liberated. Since the change of S(4)- S(0) is obviously a reduction, constantly called on when a strong oxidizing agent is while the change of S(-2)+S(0) is patently an oxida- needed. I t has been found that the oxidizing action is tion, and since these valence changes occur con- due in reality to nitrite or nitrogen oxides. These are currently, a redox reaction is taking place. The hy- present in nitric acid, either before the action starts, drogen sulfide is the reductant, and the sulfur dioxide, or are formed and constantly replenished while the nitric acid acts. If the latter is perfectly free of nitrous the indubitable oxidant.' An indirect oxidation brought about by sulfur acid (or nitrogen oxides), or if the nitrous acid (origidioxide is presented if this gas is allowed to react, under nally present or produced by the reaction) is removed, proper conditions, with moist nickel hydroxide. The the nitric acid becomes practically ineffective as an hydrous nickelous oxide is converted into hydrous oxidant. Sodium azide, or hydrazoic acid, destroys uickelic oxide, an oxidation which usually is accom- nitrous acid immediately: plished only by strong oxidizing agents such as bromine, persulfate, etc. This singular reaction was discovered by ha be^.^ It has been made the basis of a sensitive It has no effect on nitric acid. Many materials, which test for sulfur d i o ~ i d e . ~The mechanism of the reaction normally are oxidized almost instantaneously by nitric has not been wholly elucidated. Probably, basic acid, no longer are attacked if sodium azide is added t o the acid beforehand. Experiments illustrating this nickel sulfate is formed initially: inhibiting effect are described in detail in Part I of this series of papers.' 29. Aluminwm Forms Hydrated Aluminum Oxide i n the Air. Metallic aluminum, as is well known, is oxidized when i t is heated or burned in the air. On the other hand, the possibility of converting considerable quantities of the metal into its hydrated oxide without heating, and merely by the action of air and This product absorbs oxygen from the air, forming moisture, seems paradoxical a t first sight. This renickel sulfate and nickel tetrahydroxide: action is easily accomplished if alnminum foil is treated with a drop of dilute mercuric chloride solution, allowed to dry, and then rubbed. Free mercury is produced 3Hg) and amalgamates (3HgCl~ 2A1 -3 2AIC13 the aluminum. The latter is activated in this interAnother explanation is that SO2 (or HzSOa) forms an metallic compound; in other words, i t is much more readdition compound, SOrO1, by taking oxygen from active than the compact metal. Consequently i t is the air. This oxygen is assumed to be activated so that autoxidized by the combined action of the air and moisture. Hydrated alnminum oxide results; the mercury the following reaction occurs: is regenerated, and amalgamates more of the unchanged ZNi(0H)e SOz.02 NiSO* + Ni(OH), aluminum. This successive formation and decomposiAccording to both explanations sulfurous acid does tion of aluminum amalgam makes it possible for even not directly bring about the oxidation of Ni(0H)z to minute amounts of mercury to activate considerable Ni(OH)r, but only induces it. The action is indirect quantities of aluminum and thus indirectly effect the and results from an autoxidation to which the H2S03is production of large amounts of hydrated aluminum subjected. The real oxidant is the oxygen of the at- oxide. mosphere, and, in truth, the SO2 or (HzS03) is oxidized This seeming paradox of the oxidation of aluminum along with the nickelous oxide. Thus the paradoxical a t room temperature can thus be satisfactorily exbehavior of SO2 is clarified, if the focus of attention is plained as due to the catalytic action of mercury. Procedure: Aluminum foil is freed of grease by washSee FEIGL, THISJOURNAL. 21, 294 (1944). HABERAND BRAN,2. physik. Chem., 35, 81, 608 (1900). ing the metal with ether. One drop of 0.01 N mercuric
+
-
+
+
+
-
TEIGL, "Laboratory Manual of Spot Tests," translated by OESPRR,Academic Press, New York, 1943, p. 162.
+
chloride solution is placed on the clean foil. After one or two minutes, the solution is wiped off with filter paper. Excresences, which consist of hydrated aluminum oxide, appear in a short time. The same effect is obtained if mercuric iodide or sulfide is placed on the foil and rubbed with a cloth. The reactions, which occur in the solid condition, are due to the liberation of mercury:
-
3H& 3HgS
+ 2AI + 2A1-
2A1h
AllSs
+ 3Hg + 3Hg
30. Nonvolatile Oxides of Tin and Antimony Are Made to Disappear by Heating Them with a Volatile Compound. Tin dioxide and antimony trioxide withstand high temperatures. In contrast, solid ammonium iodide volatilizes when even gently ignited, because i t thermally dissociates into the gases, ammonia and hydrogen iodide. However, if a mixture of one of these oxides with an excess of this salt is ignited, the volatilization of the ammonium iodide is accompanied by the disappearance of the oxide. Consequently, this presents an apparent paradox, in that a heat-resistant material is made to disappear by subjecting i t to temperatures which normally do not affect it. The explanation lies in the fact that the mixture reacts chemically when heated, and all the products are volatile. At high temperatures the hydrogen iodide furnished by the dissociated ammonium iodide attacks the tin (or antimony) oxide: SnOt
+ ?HI
-
SnIl
+ 2Hn0
Stannic iodide boils a t about 340' C. and consequently the tin is volatilized, not as oxide but as iodide. Since Sb13 and As13 boil in the neighborhood of 400°C., SbzOaand A%OJcan likewise be made to disappear by igniting them with an excess of ammonium iodide. The decomposition of SnOz by conversion into Surd, which is then volatilized, can be used to test the purity of S~OZ.~ Procedure: A little SnOz or Sbz03 is ignited in a micro-crucible. After cooling, the residue is mixed with about 15 times its bulk of NHJ. The mixture is heated until the evolution of iodine vapors ceases. If the oxide bas not volatilized completely, an additional quantity of ammonium iodide is added and the ienition is repeated. 31. An Acid Sets a Base Free from a Salt. Weak acids are liberated from their salts bv stronp acids: likewise strong bases displace weak bases from their salts. Consequently, the notion that a base could be liberated from its salt by the action of an acid seems paradoxical. Nevertheless, this strange result can be actually achieved. If boric acid is added to a solution of an alkali fluoride, they react a t once. The products are the alkali salt of fluoboric acid and an alkali hydroxide : CALEYAND BURPORD, Ind. Eng. Chenz., Anal. Ed., 8 , li4 (1938).
Hence this singular instance shows that a weak acid can liberate a strong base from a neutral salt. The key to the paradox lies in the formation of a stable complex compound. Procedure: The reaction is carried out on a spot plate. Potassium fluoride solution (1 per cent) is treated with one drop of 1 N boric acid, to which several drops of phenolphthalein have been added. The mixture turns red almost a t once, indicating the production of a free base. 32. Permanganute Is Not Capable of Oxidizing Oxalic Acid. One of the most familiar redox reactions is:
Consequently, to state that permanganate is not capable of oxidizing oxalic acid seems to present a paradox that is contrary to extensive experience. Nonetheless, permanganate does not react at all, or with extreme slowness, if the oxalic acid solution contains molybdate in addition to mineral acid. The molybdate is completely inactive toward permanganate. The paradox is only apparent; this strange effect can be explained. The familiar redox equation holds only for free, i. e., ionogenically dissolved oxalic acid. The addition of molybdate produces no visible change, but in reality, the water-soluble complex molybdic-oxalic acid, H2MoO3C2O4,is formed at once. This heteropoly acid yields no, or only a few, GO4- ions, producing instead MoO&Oiions. The analytical properties of these two ionic species are entirely different. The most striking difference, in the present connection, is that the complex ions do not react with permanganate. In other words, the oxalate ions are masked, or sequestered in the complex.^ Procedure: Single drops of 0.1 N oxalic acid are placed in adjacent depressions of a spot plate. Four drops of 6 per cent ammonium molybdate solution are added to the first drop of oxalic acid, four drops of water to the other. Both mixtures are then acidified with sulfuric acid (1:4), and one drop of 0.02 iV permanganate is added to each. After the pure oxalic acid has discharged the color of the permanganate, three drops of permanganate are added to each of the samples. When stirred, the pure oxalic acid again discharges the permanganate color, whereas the specimen containing molybdate remains unaltered. 33. Ammonium Polysu@de Brings about an h i & tion. Alkali sulfides, free sulfur, and alkali polysulfides are quite common reducing agents. The action is due to the readiness with which free sulfur and sulfidebound sulfur are oxidized. Alkali sulfides are particularly used as reductants in the preparation of organic compounds. Consequently, it appears paradoxical that a metallic compound can be oxidized through the action of an alkali polysulfide. This exceptional case is constantly used in the classical scheme of qual-
-
(Continued on Page 353)
TEIGL, 2. anal. Chcm., 74, 391 (1928).
SPOT REACTION EXPERIMENTS (Continuedfrom $age 349)
itative inorganic analysis. After the heavy metals are precipitated from acid solution by hydrogen sulfide, the precipitate is digested with ammonium sulfide. This separates the " a c i d sulfides (As, Sb, Sn) from the "basic" sulfides, which do not dissolve in the ammonium sulfide. Yellow ammonium sulfide, which contains polysulfides, is always used in this production of the soluble sulfa salts of As, Sb, and Sn. The reason is that any divalent tin in the original solution will be precipitated by HzS as black SuS, which, in contrast t o SnS*, is not soluble in colorless ammonium sulfide. However, stannous sulfide dissolves in yellow ammonium sulfide, because ammonium sulfo-stannate is formed: SnS
-
+ (NHSSSI
(NHI)&~SI
When the solution is acidified, yellow SnSa is de. . posted:
+
(NH4)nSnS3 2HCI-t 2NH4CI
+ HIS + SnS,
The conversion of Sn(2)+Sn(4) is manifestly an oxidation. Procedure: A small pinch of finely powdered sulfur is placed in a depression of a spot plate and moistened with two drops of alcohol. This serves to lower the surface tension of the sulfur and thus facilitates the wetting with ammonium sulfide, and then the polysulfide forms more readily. Five drops of colorless ammonium sulfide are stirred in; the yellow color signifies the production of polysulfide. Five drops of colorless ammonium sulfide are placed in an adjacent depression. One drop of freshly prepared 1 per cent solution of stannous chloride is added to each of the liquids and stirred. A black precipitate, SnS, forms in- the colorless ammonium sulfide. In the other depression, the precipitate appears and then dissolves, leaving a clear solution. If dilute hydrochloric acid is added to both sDecimens, the stannous sulfide remains unchanged, whereas yellow stannic sulfide precipitates from the sulfo-stannate solution.
