the compound has not been isolated. This is not surprising because the lability and reactivity of the intermediate compounds of a catalyst are the very characteristics responsible for the catalytic effect. The compound IR--SI represents the addition of iodine to RS to form RS.1,. This type of addition compound is also involved as the intermediate in the conversion of metal sulfides into iodides, a reaction that will be discussed later. The production of such compounds is probably the reason why insoluble sulfides also catalyze the iodme-sodium azide reaction almost instantaneously. In connection with this catalysis by metal sulfides and other thio compounds, it should be noted that these materials are frequently oxidized by iodine. Consequently, the catalyst is consumed during its period of activity, that is, it is poisoned. However, the decisive factor is that the velocity of the primary and secondary reaction is greater than the rate a t which the catalyst is poisoned, and consequently the iodine-sodium azide catalysis can be used to detect traces of sul6des. Procedure (catalysis by soluble sulfides) : Two drops of 0.02 N iodine solution and three drops of water are placed in a depression of a spot plate. Two drops of a three per cent of sodium azide in 0.02 N iodine solution and three drops of water are placed in each of the three adjacent depressions. A drop of one per cent starch solution is added to each and mixed by blowing through a pipet. Each of the four depressions contains the same quantity of iodine and three contain, in addition, equal quantities of sodium azide. A drop of 0.2 per cent aqueous thiourea solution is added to the pure iodine solution and to the first iodine-azide mixture and mixed by blowing. The iodine solution remains practically unchanged (blue) while the neighboring iodine-azide mixture immediately becomes colorless. Two drops of the decolorized reaction mixture are transferred by means of a pipet to the iodieazide solution in the adjoining depression. After mixing, this solution also becomes colorless. Three drops of the latter are transferred to the blue solution in the fourth depression. Several seconds after mixmg, this solution also loses its color. The experiment shows that the catalyst has not disappeared when the iodine-azide reaction is completed. It is still active and despite dilution remains capable of accomplishing the acceleration a second and third time. Procedure (catalysis by insoluble sulfides): The iodine and iodine-sodium azide solutions used in the preceding experiment are diluted to one-eighth their original strength. Solid sodium azide (tip of knife blade) is added to the diluted sodium azide solution to make the following experiment more distinct. One drop of each of the solutions is placed on qualitative Elter paper containing starch. It is convenient to lay the filter paper on a broad rubber stopper. A blue spot appears in both cases. A piece of galena or pyrites, whose surface has been ground smooth, is
pressed against these iodine-starch spots. No change will be observed in the pure iodine-starch fleck, which accords with the known resistance of natural sulfdes . to the action of iodine. In contrast, a distinct white\ impression of the ore will be left on the iodine-starch spot that also contained sodium-azide. The experiment demonstrates that even insoluble sulfides in compact masses can initiate the iodine-azide reaction. 2. Autocatalysis of the Permangunate-Owllic Acid Reaction. The reaction of oxalic acid with potassium permanganate in the presence of sulfuric acid can be written: 5&C204
+ ZKMnO, + 3HsS04 = K&O, + 2MnS0, + 10C02+ 8H10
However, this reaction does not proceed directly because a drop of potassium permanganate solution is not decolorized immediately when it is added to a solution containing excess oxalic acid and ,some sulfuric acid. The pink color of the MnOd- ion persists; it disappears only after some time. As more drops of the permanganate are added the time required for the discharge of the color of the individual drops constantly decreases. Fmally, the color disappears at once, until all of the oxalic acid has been consumed. This peculiarity, which is not observed in other oxidations with permanganate, indicates in itself that the foregoing equation does not truly represent the actual direct course of the reaction between permanganate and oxalic acid. The fact that the reaction goes slowly at first and picks up speed as it proceeds shows that the correct conditions for the oxidation of oxalic acid are being established during the period when the reaction is going slowly. In fact, a study of the kinetics of this reaction by Skrabal' showed that while the equation is perfectly correct with respect to the relative weights of permanganate and oxalic acid, it is, however, only the net equation produced by combining 10 partial reactions, some of which proceed instantaneously, others with measurable velocities. While the color is fading slowly, that is, in the "incubation" period, the Mn (VII) of the permanganate undergoes a slower, reduction to the unstable Mn (111) as shown in (1); this Mn (111) salt reacts instantaneously with oxalic acid and is reduced to Mn (11) as in (2). Incubation Mn (VII) period Mn (111)
{
+ o d i c acid + oltalic acid
--
Mn (111)
+
COB (1) slow Mn (11) COB (2) instantaneous
+
The Mn (11) formed in (2) can now react with RMnO, as indicated in (3) Mn (11)
+ Mn (VII)
-
Mn (111)
(3) rapid
The Mn (111) thus produced undergoes immediate reaction with oxalic acid according to (2). Consequently, (1) and (3) compete with respect to producing the reactive Mn (111). Since, however, (3) goes faster than (1) it is obvious that (3) will predominate, provided sufficient Mn (11) is present. This means that
2. anorg. Chem.. 42, 1 (1904).
the Mn (11) formed by the sequence (I), (2) assumes the role of a catalyst. I t reacts with one participant (KMnO,) of the slow oxalic acid-permanganate reaction and forms as intermediate product Mn (111), which reacts rapidly with the second participant (oxalic acid). Consequently, this catalyst does not have to be introduced separately into the system, but is formed during the incubation period. The supply is not only maintained but is constantly increased. This is a clear case of autocatalysis. These details make it obvious that the incubation period of the permanganate-oxalicacid reaction can be eliminated if catalytically active Mn (11) is added before starting the reaction. This fact also establishes the autocatalytic character of this reaction. Procedure: Two drops of six per cent oxalic acid solution and two drops of dilute sulfuric acid (1: 2) are placed in each of two adjacent depressions of a spot plate. A drop of 0.02 per cent manganous sulfate solution is added to one of the mixtures; a drop of water is added to the other. Both solutions are then mixed by blowing into them through a pipet. Three drops of 0.03 per cent potassium permanganate solution are added to each of the two solutions and the mixing is repeated. The solution containing the Mn (11) salt will lose its color almost instantly, whereas the pink will persist for about four minutes in the other solution. Then it begins to fade gradually. The experiments demonstrate both the existence of an incubation period of the permanganate-oxalic acid reaction and the shortening of this period by the catalytic action of manganous ions. 3. Oxidation Reactions Catalyzed by Nitric Acid. Nitric acid and nitrates in acid solution are used more frequently than any other oxidizing agents. A general oxidation equation can be written: 2HNO8
+ 3R
-
3R0
+ 2 N 0 + H20
(R = dinlent metal or other reducing agent)
2HN01+ R
-
RO
+ 2 N 0 + HzO
(1)
This equation also indicates the true course of the oxidation when nitric acid is used. because the NO formed in (1) reacts with nitric acid to produce nitrons acid : 2N0
+ HNOI
-
HzO
+ HNOI
(2)
Conseqnently, if nitric acid originally contains nitrous acid, the latter reacts first as shown in (1). Nitric oxide (NO) is produced, and by reaction with nitric acid (2) again forms nitrous acid, which continues the oxidation. Accordingly, nitric acid does not function at all as the direct oxidant, but acts simply as a source of nitrous acid. Consequently, if nitric acid is to act as an oxidizing agent, nitrous acid mnst either be present from the start or it mnst he formed by a subsequent reaction. However, even traces of nitrous acid suffice,because, as shown by (1) and (Z), it is furnished in increasing amounts by the nitric acid. Even the purest nitric acid acts as an oxidant, although very slowly at first, and only after an incubation period. Its action can be explained by the fact that nitrous acid is formed in this first stage. I t then continues the oxidation as represented by (1) and (2). These two equations, therefore, picture the mechanism of the oxidizing action of nitric acid. The net oxidation equation of nitric acid, given earlier, is obtained by combining them (with adjustment to involve the same number of molecules of HN02). The dominant role of nitrons acid can be shown experimentally by the fact that many familiar and common oxidizing actions of nitric acid either do not occur or are greatly weakened if the nitrous acid, which is present or produced, is destroyed. The following experiments demonstrate both the primary formation of nitrous acid, and also the retardation of the oxidizing action of nitric .acid if the nitrous acid is removed. 3a. Detection of Nitrous Acid During the Reduction of Nitrates in Acid Solution by Nascent Hydrogen. The formation of nitrous acid in the 6rst stape of oxidation by nitric acid is easily demonstrated through its reaction with alkali iodide. Free iodine is produced and is detected by the familiar starch test. Nascent hydrogen, generated by the action of zinc on nitric acid, is used as the reducing agent in the following experiments. Consequently, if zinc is added to nitric acid (or an acidified-nitrate solution) containing potassium iodide, the following reactions must be taken into account: ...-
Thanks to this oxidizing action, nitric acid, in contradistinction to other acids, is capable of dissolving even such metals as copper, etc., that are below hydrogen in the electromotive series. The metal is converted first into an oxide (CuO, for instance) and this, in turn, into the soluble nitrate through the action of the hydrogen ions. More than a century ago it was observed that pure nitric acid is much less active in the oxidative solution of such metals than impure acid, or acid which had been used previously and which consequently contained nitrozen oxides or nitrons acid. ~ a &it was found that the solvent power of 2H+ Zno --c Zn++ 2H formation of naxent and pure nitric acid is accelerated tremendously by the H + H ,H, } molecular hydrogen addition of nitrite. Therefore nitrite can be regarded reduction of nitrate to nitrite as a catalyst. Detailed studies by Abel of the kinetics (molecular hydrogen is in+ 2H HnO + of oxidation by nitric acid have shown that the actual active herd active agent is not the nitric acid but the nitrous acid, ox~datmnby nitrite of iodide 4H+ which, as will be seen presently, can he formed from 2N01- 21;o free iodine 2H10 2NO In nitric acid in the course of the oxidation. I t should be noted that there has been no mention It is well known that nitrous acid is a stronp oxidume here of another reaction that can occur, namely, the agent. Its action can be represented:
-
-
~
+
+
-
+
-
-
++
-+
1
,
+
reduction of free iodine by nascent hydrogen: IS colorized solution and, if necessary, a fresh particle of 2H -+ 2HI. However, this reaction proceeds so very nitrate. The experiment shows the acceleration by slowly that i t is of no significance with respect to the nitrate of the reduction of potassium permanganate by speedy oxidation of iodide ion by nitrite. Conse- nascent hydrogen. quently i t does not interfere. 36. Prevention of the Ozidation of "Molybdenum Procedure: One drop each of 1 N potassium nitrate, blue" by Concentrated Nitric Acid in the Presence of of one per cent potassium iodide solution (acidified Urea. Colorless, acidified solutions of alkali molybwith dilute hydrochloric acid immediately before the date are turned deep hlue by metallic zinc (and other experiment), and of starch are placed in the depression reducing agents). The hlue is due to the formation of of a spot plate. Since pure, dilute nitric acid does varying amounts of lower oxides of molybdenum not liberate iodine from potassium iodide the solution (MozOo, MmOs, etc.) or their soluble salts. Such remains colorless, or turns a barely visible blue. A systems are commonly referred to as "molybdenum granule of zinc is then added and a deep blue appears blue solutions." Molybdenum blue is gradually oxiimmediately. The experiment demonstrates the forma- dized by atmospheric oxygen, and very rapidly by tion of nitrous acid during the reduction of an acidi- oxidizing agents of all kinds, including, of course, nitric fied nitrate solution by nascent hydrogen. acid. The solution is decolorized by a reaction that 3b. The Reversible Reaction NOs- g NOz- as a can he represented : Catalyzing Reaction. Potassium permanganate is re3M020a 2HNOs 6MoOa 2 N 0 H1O duced onlv slowlv . hv . nascent hvdroeen: . It has already been shown that nitric acid owes its 2Mn0410H 6HC 2Mnf+ + 8H2O oxidizing ability solely to its content of nitrous acid, or Accordingly, the color is discharged only gradually if to the formation of the latter. However, nitrous acid zinc is added to an acidified solution of potassium per- is completely and almost instantly destroyed by urea. manganate. This reaction can be accelerated catalytically by nitrate, because the following partial reactions occur through the valence changes of the nitro- Consequently, molybdenum blue will resist the action gen. First, the nitrate is reduced to nitrite by nascent of concentrated nitric acid if urea is present. The hydrogen; that is, by one participant of the slow re- following experiment substantiates this statement. A action as has been shown in Experiment 3a (primary solution of molybdenum blue is prepared by warming reaction 1, below). The nitrite thus formed is re- a five per cent solution of ammonium molyhdate with oxidized to nitrate by potassium permanganate, the zinc and dilute sulfuric acid. A drop of the deep blue second participant of the slow reaction (secondary re- solution is placed in each of two adjacent depressions action 2, below). These two partial reactions proceed of a spot plate. Two drops of a saturated water soluvery rapidly and the sequence of reduction and oxida- tion of urea are added to one portion, two drops of tion, that is, the reversible change N (V) i2 (111) goes water to the other. Two drops of concentrated nitric on until all the permanganate is consumed and the acid are then added to both solutions and mixed by color bas vanished entirely. This series of reactions blowing through a pipet. The solution containing urea is that of a typical intermediate reaction catalysis. will remain blue for hours, while the other becomes The summation of the equations representing the two colorless within a few seconds because the lower oxides rapid, partial reactions (1) and (2) gives the equation of molybdenum are oxidized. (3) for the reduction of permanganate by nascent A white precipitate fonns in the hlue solution conhydrogen, which, of itself, proceeds slowly, hut is taining urea; it consists of the difficultly soluble urea rapid if nitrate is present. nitrate. This compound can be brought into solution by adding five drops of water. If this solution is com5NOa- + 10H 5N01- + 5H10 (1) primary reaction pared subsequently with a blank test prepared by 5N02- + 2MnOai + 6Ht 5NOa2Mn++ 3H.O (2) secondary reaction addincr nine drops of water to one drop of molybdenum blue solution, the colors will be identi'cal. 2Mn0,- + 10H + 6H+ 2Mn++ + 8Hz0 (3) = + (2)
-
+
+
-+
+
-
+
Procedure: A granule of zinc and two drops of 1 N sulfuric acid are nlaced in each of two adiacent depressions of a spot plate. As soon as a distinct evolution of hydrogen is under way, a drop of 0.1 N potassium permanganate solution is added to each and mixed by blowing through a pipet. A tiny particle of potassium nitrate is dropped into one mixture. The color disappears completely within a few seconds, while the other mixture, that contains no nitrate and hence is uncatalyzed, does not begin to lose its color until several minutes have elapsed. The experiment can be repeated by again adding permanganate to the de-
-
+
-
+
+
-
3d. Prevention of the Oxdution of Potassium Thiocyanate by Nitric Acid in the Presence of Sodium Azide.
The foreeoine ex~erimentsshowed that molvhdenum blue is not oxidized by nitric acid in the presence of urea because the real oxidizing agent, nitrous acid, is destroyed completely. The action of concentrated nitric acid on potassium thiocyanate in the absence or presence of a compound that destroys nitrous acid will be compared in the present experiment. Potassium thiocyanate is oxidized to cyanate and sulfate by nitric acid: -
3KCNS
-
A
+ 8HNOI +
3HI0
-
aKCNO + 3Hd04 + 8N0 + 4H.O
Consequently, if this oxidation occurs in the presence of barium chloride, precipitation of barium sulfate ensues. Urea cannot be used to remove nitrous acid in this instance because a precipitate of slightly soluble urea nitrate will be formed. Sodium azide is used: HNOn
+ 3HNs
-
2Hn0
+ 5N1
Therefore, if the oxidizing action of nitric acid is due solely to nitrous acid, a solution of potassium thiocyanate containing barium chloride should remain unchanged on the addition of nitric acid, provided sodium azide is present. On the other hand, a precipitate of barium sulfate will follow the addition of nitric acid if no azide is present. Procedure: Ten grams of potassium thiocyanate are
dissolved in 10 ml. of five per cent barium chloride solution. A black spot plate is kept a t 120' for five minutes in a drying oven, or the plate is warmed on an asbestos mat heated by a gas flame. A drop of the KCNS-BaCla solution is placed in each of two adjacent depressions of the warm spot plate. A pinch (pen-knife tip) of sodium azide is added to one of the drops and stirred in with a thin glass rod. Two drops of concentrated nitric acid are added to each solution. After a few seconds, a strong evolution of gas occurs in the azide-free solution, while only a few bubbles of nitrogen appear in the other. When no more gas is given off, ten drops of water are added to each solution and mixed bv blowing through a ninet. A nrerinitate ~, ~-~ ~ - --- - - - ~ r - r - - - -- rof barium sulfate will be seen in the azide-free solution. whereas the other is perfectly clear. ~
~
~~
~ - 0
~~
