I X D C C E D REACTIOKS BY TESLEY G. V A N S O Y
Mellor’ and Friend2 both include in their text-books good discussions of the early work on induced reactions. I n general, F. ;\10hr3 has received credit as being the first to notice such a phenomenon. He noted, over a century ago, that an aqueous solution of sodium arsenite did not undergo any perceptible change when shaken with air, while a solution of sodium sulphite was oxidized rapidly under the same conditions. Further, he found that when a mixture of sodium sulphite and sodium arsenite was treated in this manner both salts %ere oxidized. Yarious names have been used to designate this type of reaction, such as chemical induction, auto-oxidation, sympathetic, coupled, and induced reactions. Kessler4 appears to have been the first to study the subject systematically and he introduced the term, induced reactions, and the other nomenclature which is used most commonly at the present time. He called the fastest reaction the primary reaction; the induced reaction caused by or helped along by the change the secondary reaction; the substance which takes part in both reactions the actor; the substance which takes part in the primary reaction the inductor; and the substance which takes part in the secondary reaction the acceptor. He found many examples of induced reactions but made no attempt to explain them. Since his time, numerous theories have been suggested to explain the nature of these reactions. The most important of these will be mentioned briefly in this article. Manchot‘s Peroxide Theory5:-In every process of oxidation there is formed a primary oxide which has, in general, the characteristics of a peroxide. The peroxide is formed directly from the reagents, and intermediate states are passed over. aFeO n o y = Fe20,, + * = Fez O3 (211 - 1)0
+
+
He believed that the formula could be deduced from a determination of the induction factor, which is the proportion in which the actor divides itself between the inductor and the acceptor. Amount of acceptor oxidized Induction Factor = Amount of inductor oxidized Luther and Schilow’s Classification6:-They believe that information on the nature of the intermediate compound formed can be obtained by suithlellor: “Chemical Statics and Dynamics,” 333. Friend: “Textbook of Inorganic Chemistry,” 7 I, g j . 3 Mohr: “Lehrbuch der chemischen-analytischen Titriermethoden,” 2 7 1 (1855). Pogg. Ann., 95, 216 (I8sj); 119, a18 (1863). 5 hlanchot: Ann., 325, 93, ~ o (1902). j 6 Luther and Schilow: Z. physik. Chem., 46, 7 7 7 (1903). 2
I594
WESLEY G. YANNOY
able variation of the reagents. “Observations of the occurrence or nonoccurrence of induction when the reagents are varicd systematically enables one, therefore, to a certain extent to decide which of the substa,nces play a specific rBle in the reaction.” The nature of the classification is as follows: Class A, the inductor is ‘specific’and the actor and acceptor are ‘non-specific.’ Classes B and C, the actor and the acceptor take the place respectively of the inductor in Class -1. Class D, both the acceptor and the inductor are ‘specific,’ and the intermediate body must be regarded as a complex derived from both of them. Classes E and F, the inductor and the actor, or the acceptor and the actor are ‘specific‘ respectively. hliller’s Classification’.-He prepared a classification based on kinetic measurements. According to this, the induced reactions are divided into three classes. Class I . Cases of catalysis combined with the destruction of the catalyzer. Class 2 . Cases in which the reaction between the actor and the inductor is the same whether the acceptor be present or not. The peroxide theory would come under this heading. Class 3. Cases in which the reaction between the actor and the inductor changes when the acceptor is present. This classification of course involves kinetic measurements and the comparing of the rates a t which the actor reacts with the inductor and acceptor separately and combined. In addition to these three classifications, a number of hypotheses have been advanced to explain the particular reaction or group of reactions studied. Schiinbein? in 1858 attempted to explain the formation of hydrogen peroxide or of ozone, when many different substances mere exposed to atmospheric oxidation on Brodie’s assuniption3 that the oxygen molecule consists of a positive atom united to a negative atom. Under these conditions he argued that the metal would unite with the negative atom and the water with the positive atom. However, he could riot subrtantiate this view with any experimental evidence and, as a consequence, it has not been generally accepted. Hoppe-Seyler4 suggested the formation of naicent oxygen during the oxidation of a substance, which then united with water to give hydrogen peroxide. Traubej would not accept Hoppe-Seyler’s explanation and expressed the view that, since water was essential for the oxidation, the oxygen used by the substance undergoing oxidation came from the water. The hydrogen thus liberated would react with molecular oxygen giving hydrogen peroxide. This theory was modified somewhat by Bach6 and by Engler and Wild,7 who held that the oxygen molecule could combine as a whole wit’h subl l i l l e r : J. Phya. Chem., 11, 91 IIg07’8. 75, 99 (18j8). 3Brodie: J. Chem. Soc., 4, 194 ( 1 8 j z ) ; 7, 3 0 j ( 1 8 j 5 ) . Hoppe-Seyler: Z. physiol. Chem., 2, 2 2 2 (1881); Her., 16, IIoppe-SeyelPr: Z. physiol. Chem., 2 , 2 2 2 (1881); Ber., 16, 6Bach: Compt. rend., 124, 9j1 r189j). Engler and lf-ilcl: Ber., 30, 1669 :1897).
* Schonbein: J. prakt. Chem..
11;
(1883).
11; (1883).
ISDUCED REACTIONS
I595
stances other than hydrogen, giving rise to higher peroxides of the inductor which then react with the acceptor, Schilow’ made use of stable and instable intermediate compounds to explain the mechanism of the reactions which he studied. Miss Benson2 developed what she called a “Ferro-iod-ion Theory,” assuming the formation of an intermediate compound to explain the mechanism of the reaction, chromic acid, ferrous sulphate, and potassium iodide. Winther3 suggested short-wave radiation as an explanation of reactions of this kind. Each of the above hypotheses may apply t o a limited number of induced reactions; but it is rather difficult, to be sure just where any particular one may apply. PIIanchot’s peroxide theory apparently does not hold for all induced reactions, according to results as shown by PIliss Benson, and further, it seems that the peroxide which he postulates must vary in composition, without, any reason being given, in order to account for the reactions which he studied. The classifications by Miller, as well as those by Luther and Schilow, are general and could perhaps be applied to all induced reactions. I n many cases, however, the application would be somewhat tedious and, even after the classification, the true mechanism of the reaction would still be in doubt. With these facts in mind, Professor Bancroft4 has developed and introduced the following theory of induced reactions. Let A be an oxidizing agent (actor) which will not react or which reacts very slowly with a reducing agent C (acceptor); but which reacts with a reducing agent B (inductor) and with C in presence of B. The ratio of C oxidized by X to B oxidized by A is called the reduction factor. Case I . B may catalyze the reaction between A and C. If so, the induction factor will increase indefinitely with the relative increase of C in case one adds h to a mixture of I3 and C. The induction factor will approximate zero in case one adds B to a mixture of A and C so slowly that B is used up practically instantaneously. B will be specific and we cannot substitute another reducing agent for it. Electrolysis will not necessarily accelerate the reaction between arid C. Case 2 . C map react with a lower stage of A. Thus HBrCh --t HBrO? which reacts with C. Slow addition of B to a mixture of A and C should give approximately two for an induction factor with bromic acid as A . dddition of bromic acid to a mixture of B and C mill give an induction factor of approximately zero when B is in large excess, and a value of two when C is in large excess. K i t h different B’s the induction factor will vary for any given concentration of B and C. -4is specific; but B may be any suitable reducing agent. Case 3a. C may react with a stable oxidation product of B. So long as there is an excess of C, B xi11 remain apparently unchanged and the induction Schilow: Z.physik. Chern., 42, 641 (1903). Benson: J. Phys. Chem., 7, 356 ( 1 9 0 3 ) . Winther: 2. physik. Chem., 100, 566 ( 1 9 2 2 : . Bancroft: J. Phys. Chem., 33, 1184 f1929).
1596
WESLEY G . VASNOY
factor will be infinite. With an excess of A the induction factor will be determined by the actual ratio of C to A. .A case of this kind is h = 02,B = H I , and C = 802. This is what has been called a consecutive reaction. Case 3b. C may react with a higher instable stage of B. Thus Fe" + Fe" which reacts with C. If B is added slowly to a mixture of X and C, the limiting induction factor should be obtained. If B is added to h alone, it should be possible to get more reduction of A than corresponds, for instance with the formation of stable Fe"' (evolution of oxygen). B is specific; h may be any suitable oxidizing agent. One might perhaps substitute but a suitable anode for A. Case 4. B and C form a complex (hydroquinone or aldehyde with sulphite) which reacts with A. If we add X to a mixture of B and C containing so much C that there will be no free B, we shall get' a limiting induction factor provided the rate of reaction of h with C is negligible. Addition of B to a mixture of h and C will give approximately zero for the induction factor if the rate of reaction of A with B is high relatively to the rate of formation of the complex of B and C. B and C are specific; but A may be any suitable oxidizing agent. X suitable anode may be substituted for A. Case j . Nos. 1-4 may occur simultaneously in any combination. This will usually show itself by positive tests in at least two of the preceding cases. Each problem will then have to be considered on its own merits. The order of the reaction velocity equation for A and B may be normal for Xos. I and 4, whereas it may be abnormal for Nos. 2 and 3, being necessarily abnormal for A in S o . 2 and for B in S o . 3 . With Sos. I , 2 and 3, the order of the reaction velocity of A and B will be independent of C ; but this will not be true of KO.4. The effect of temperature should be studied. The relative stabilities of the various intermediate compounds may vary considerably with changing temperature. One difficulty here will be that at the higher temperatures .A may react with C a t any annoying rate even in the absence of B.
Experimental Part Chromic Acid, Arsenious Acid and Tartaric Acid Among other reactions, Schilowl studied the induced oxidation of tartaric acid by chromic acid in the presence of arsenious acid. The reaction was previously observed by Kessler who had verified this induction in the case? , SnO, and SOsacted as inductors. The where -AS203,FeS04,K 4 F e ( C S ) ~Sb203, results of Schilow's investigation are given in Table I, and Figs. Ia and Ib. He obtained a maximum induction factor of 2 . 8 with increasing sodium tartrate concentration and hence with an increasing ratio of tartrate to arsenite concentrations. He further showed that the induction factor was independent of the concentration of the arsenious acid provided the ratio of the concentrations of tartrate and arsenite were kept constant. When he LOC.cit.
I597
INDUCED REACTIONS
I
1
c
/---
E9 A/ i:
p
I -f
J GONG O F A C C E P T O R
0
FIG.r b
TABLEI Reaction Mixture: I O cc N / I O OK2Cr20i; I O cc S ~ IHzSOa; O 5 cc Y!IOO AsnOB;5 cc ofdcceptor. Titrated with K;ZOONa2S203and r\;!zoo I? Solutions. Conc. Ace. Conc. hcc. Molar Ind. Factor Conc. lnd.
Conc. Acc. Slolar
Ind. Factor
.70
Conc. Acr. Conc. lnd.
I/ZOOO
0.25
1(=0.20)
11200
I
1/1600
0.30
I .25
I!150
2.02
13 . 2
I!l400
0.34 0.40 0.46
1.43 I ,65
I/IOO
2.
0.58 0.75
3.3
I .07
5
1/10
2.32 2.69 2.79 2.80 2.80 I .$9
20.
djo 1/30 1/20 1/15
1.39
6.6
I/IZOO
I/IOOO I
!goo
1/600 I 1400
1/300
2 . j
IO.
40
66. 100.
132. 200.
1598
WESLEY G. \‘ANNOY
increased the concentration of the chromic acid the determinations became difficult to carry out due to spontaneous or voluntary oxidation of the tartaric acid. Considering these results in the light of the above theory it would seem that the nature of t,his reaction was described by Case 2 . Xamely, that the tartaric acid (acceptor) must react with a lower stage of chromic acid (actor) since the arsenious acid (inductor) was not specific and could be replaced by various other reducing agents. Under these conditions however with the maximum induction factor of 2 . 8 as determined by Schilow, the formula for the lower stage of chromic acid would be complex and improbable. (CrjOn, etc., depending on just how one postulated the reaction to occu1.)
TABLE I1 Solutions: X,’IOOK2CrPO;;K#’IOHP801;N I I O O .1s203j N ‘zoo K a 2 S 2 0 s ; K,’200 IP and 3Iolar KaKC4H406(From which the necessary concentrations for the acceptor were made by dilution). Time given for reaction to take place 1-24 hrs. 11-96 hrs. Temp. zz°C (Room). Control Reactions: I. I O cc K2Cr207= 2 0 cc Sa2S203 2. I O cc ICPCr?O;= j cc As203 = I O .I cc Ka2S203 3 . I O cc S a ? S 2 0 s= I O cc I P 4. I O cc K P C r 2 0 7 j cc 1 / 1 0 KaKCiH406 = 20 cc S a P S 2 0 3 I O cc K?Cr207 I O cc H?S04 j cc S a K C 4 H 1 0 B = 1-6.5 cc Na2S203 j. 11-0 cc KaP SP O3 Reaction Mixture: I O cc Ei? Cr207j I O cc H2S01; j cc As203and 5 cc hcc. as indicated.
++
+
Titre I1
Acc.
Titre I
I
31’1200
6.1
2
0.2
4
LI/400 0.2 ~II’IOO 3 . 9 3I/’20 4.5
j
hli’I0
4.2f
6
hIolar 2 ; cc Molar
so.
3
7
4,3+
5.9 3.8; 4.6
Ii2Crz01 Ii2Cr20i used by used by Ace. Ind. 4.1 IO IO,^ 9.8 14. 6.1 14.6 5 . 5
14.3
5.7
I.F.
I.F., (by dchilow)
0.41
0.40
1.0j
1.07
2.29
2.32
2.8
2.8 2.8
2.5f
Agent
Ka2Sz03 I*
,, ,I
,,
Fading end-point
I t was then thought that the reaction might be a combination of Cases I and z or possibly Case 1 in the above theory, the former however being the most probable. If this mere the solution then an increase in the ratio of the concentration of the acceptor to the concentration of the inductor should give rise to a considerable increase in the value of the induction factor. With this in view Schilow’s work was repeated and an attempt was made to work at higher concentrations n-ith the acceptor (tartaric acid). His methods were followed throughout. Karnely, definite quantities of potassium dichromate, arsenious acid, sodium tartrate and sulphuric acid were
INDUCED REACTIOSS
I599
mixed and allowed to react at room temperature. After a time sufficient for the reaction to go to completion either the excess arsenious acid was titrated with iodine in an excess sodium bicarbonate solution or the excess chromic acid by means of iodine liberated from potassium iodide in acid solution with sodium thiosulphate. This was done by adding an excess of potassium iodide and I O cc of 2N sulphuric acid, then after two or three minutes the solution was diluted and titrated as rapidly as possible. This method was shown to be satisfactory by Schilow. The method used to calculate the amounts of chromic acid used by the acceptor and the inductor was as follows: I . Excess Chromic Acid, M / 1 2 0 0 hcc. Used by Acc. 10.1- 6 = 4.1 pts. 19 ,,Ind. 2 0 . 1 - 1 0 . 1 = IO pts. 2 . Excess Arsenious Acid, M / 2 0 hcc. Used by Acc. 10.1 4.5 = 14.6 pts. ,, Ind. I O - 4.5 = 5 . 5 pts. The data obtained, as seen from Table 11, checked Schilow’s results in a very satisfactory manner for all concentrations of the acceptor up to the tenth-molar solution. In all cases however where the concentration of the acceptor was tenth-molar or above, checks could not be obtained. Accurate titrations could not be made because of fading end-point. On standing, potassium tartrate was precipitated from the solution. Further it was noted that K2Cr201oxidized NaKC4H406directly and to a considerable extent in an acid solution as is shown by the fifth control reaction in Table 11. These difficulties have been investigated and will be discussed in more detail.
+
1 )
Investigation of Tartrates The following tartrates were examined with the idea of trying to find one which would not react with chromic acid or one which would react only with the slowest possible rate. IO cc of a molar solution of the various tartrates were mixed with I O cc of N/IOOK2Cr20;and I O cc X!/IOH2S01. The time required for the complete reduction of the K2Cr20; was used as a means of comparison. so. Tartrate Time Recrystallized S a K C 4 H 4 0 6(washed with alcohol) I. 1 2 hrs. 2. Kahlbaum’s SaKC4H4o6 I2 ” ” pwd. 6 ” 3. >, 18-24 I’ 4. Baker‘s ” recrystallized ” (not washed) I8 ” 5. 6. Kahlbaum’s HBC~H~O~ I O minutes 19 hrs. K~C~H~OG 7 . B and A 1, 8 . 1Ierck’s 24 ” Na2C4H4o6 18 ” 9. B and A Llerck’s 19 IO. 23 ” Prepared 11. NaKC4H40s 20-23 ”
1600
WESLEY G. V A T S O Y
The rochelle salts were made as indicated by the folloiving equations using Merck's potassium tartrate and C.P. sodium carbonate. i?;a2('Oa for analysis). K,C',H,O, H~SO~--+-KHCIHIO~ K2SOI
+
zKRCrHaOs
+
+ Sa2C:Os--+zIiSaC,Hr0, + H?CO,
The acid tartrate crystals were washed thoroughly and dried. Calculated amounts were then dissolved in water with sodium carbonate. After the reaction was complete the rochelle salts were crystallized from the solution. From the tests it was concluded that chromic acid would oxidize slomly any pure tratrate. Further, that XIerck's sodium tartrate would serve bcst in the induced reaction under consideration since its rate of oxidation by chromic acid was lower than that of any other tartrate, with the exception of Merck's potassiuiii tartrate, and its solubility would be sufficient to prevent the precipitation of the acid tartrate in acid molecular solutions of the salt,
Investigation of the Fading End-point In the foot-notes of his article Schilow made mention of this fading endpoint and stated that it was probably due to the reduction of the iodine by the oxidation products of the tartrate. He however iiiade no allowance for this in his iodine titrations. Since good end-points w r e obtained in the cases of low tartrate concentrations two methods of obtaining good end-points in the case of high tartrate concentration seemed possible. First, precipitate the tartrate in the form of acid potassium tartrate, filter and titrate the filtrate. Second, after the reaction had come to equilibrium, dilute the desired amount and make the titration. Seither of these two methods hoivever proved to be satisfactory as fading end-points were obtained in both cases. Solutions of varied coniposition were then titrated with iodine in an attempt to locate the cause of the fading end-point. The eoncentration of the solutions used were, molar Sa2('4H406, for the rest the same as used in the previous experiments. so. I.
Solution
j
cc XsaOs
+
cc Ill'itrt.
+
IO.
z
End-point
Does not fade Fades \-ery slowly Does not fade
j cc SazC4K40c I O ec H2S0, I O ,I cc As203 j cc K2Cr20i I O ce H,PO, IO 3 4. 5 cc S a 2 C r H , 0 6 j cc K?('r?Oi IOCC H & 0 4 16.6 (in I hr.j Fadesrapidly (after chromic acid was reduced) I O cc H ~ S O I _Fades slowly j. I O cc Sa2C4H406 2.
3.
"
IO
+
+ +
+
+
From these titrations it maq evident that iodine waq slowly reduced by the acid sodium tartrate solution alone, ( S o j ) . Further, that this reduction took place much faster if the tartrate solution was first partially oxidized by a small amount of chromic
I S D U C E D RESCTIOXS
1601
acid as shoxn by the titration results of solution S o . 4. This showed that the solution of the difficulty must lie in the oxidation products of the tartrate or the reduction products of the chromate. This was tested by noting the time required for the reduction of a limited amount of iodine by solutions of sodium tartrate and sulphuric acid containing small amounts of the oxidation products of tartaric acid such as, oxalic acid, formic acid and glycollic acid. T o particular one of these could be said to cause the fading end-point altho they all served to make the tartrate a bit more easily oxidized by the iodine. The addition of chromic sulphate had little or no effect on the time required for the reduction of the iodine. There remains a possibility that glycollic aldehyde may be the material which caused the trouble ; however none could be obtained and this test has been necessarily omitted. Bunsen and Roscoe' in their article on "Photo-Chemical Researches" investigated the reaction between dilute solutions or bromine and tartaric acid. They noted the auto-catalytic nature of the reaction and concluded from their \york that the cause of increase in the speed of the reaction as it progressed did not lie in any peculiarity v+hich light possessed but rather to the mode of action itself. In other words the reaction was autocatalytic because of cheniical induction and not because of photo-induction. They however did not try to explain the cause of this chemical induction. It was possible however to do away with the fading end-point as far as concerned in this particular investigation by a slight change in procedure. Sarnely, by working always with an excess of chromic acid and thus making it unnecessary to titrate arsenious acid in the presence of partially oxidized tartaric acid. This procedure proved to be satisfactory and was used in all the remaining experiments. Direct Oxidation of Sodium Tartrate by Chromic Acid The oxidation of sodium tartrate was carried out at various temperatures, IOO', 7 2 " and 22' (room temperature) in order to determine the amount of oxygen taken up per mol of sodium tartrate. The reaction was allowed to run for a definite time after which the excess chromic acid was determined by adding potassium iodide and sulphuric acid, and titrating the iodine liberated with standard sodium thiosulphate. The oxidation at 100°C was carried out by boiling the reaction mixture under a reflux condenser for the required time after which it mas cooled in ice mater for five minutes, then titrated immediately. For the oxidation at 72°C the reaction mixture was brought t o 7 2 T in boiling water as quickly as possible after mixing, then placed in a drying oven and held at that temperature. *At the end of the required time it was cooled in ice water for five minutes and then the excess chromic acid was titrated immediately. Fig. za, b and c shows the variation in the amount of oxygen used per mol of sodium tartrate, per unit time at the various temperatures as given in Table 111. From this work it was apparent that the speed of the reaction Bunsen and Roscoe: Phil. Trans., 147, 399 (18j7).
1602
WESLEY G . VANNOY
TABLEI11 Reaction Mixture: 5 cc M/IOOKa2C4H4O6;IO cc M/IO H2S01 and 50 cc M/IOO K2CreOr Pt. I Oxidation at I O O O C . NO.
Time in hrs.
Excess KLhOl
I
I,/2
2
I
3 4
2
6
5
I2
6 7
24
43.31 32.75 21.6 15.75 13.70 12.46
48
IO .3
Pt. KO.
2
Oxidation at
I .72
2.8 3.48 3.63 3.75 3.96
72OC.
KpCr207
49.28 48.37 44.41 33.88 17.34 16.02
0.72
I/2
I
3 4
6
2
5
I2
6 7 8
24
48 62 86
P t . 3 Oxidation at
'5.31
14.5 13.99 22OC
Excess K~Cr207
I
2
2
I2
49.i9 49.28 48.47 43.6 39.14 35.18 28.90 19.57
24
4 5
6
48 60 77
7
IO0
8
432
used
I .63
5.59 16.12 32.66 33.98 34.69 35.5 36.01
Oxygen used per Mol. Tart.
0.07 0.16 0.j6 I .61 3.26 3.4 3.47 3.54 3.6
(Room Temp.)
Time in hrs.
3
0.67
28.4 34.25 36.3 37.54 39,7
Excess K?Cr2Or
I
No.
6.69
Oxygen used per Mol. Tart.
17.2j
Time in hrs.
2
9
K2CrLh used
K2Cr20, used
Oxygen used per Mol. Tart.
0.21
0.021
0.72
0.07
1.53
0.1j
6.4
0.63 0.96 I .48
10.86 14.82 21 .I
2.1
30.43
3.04
between chromic acid and tartaric acid was considerably decreased by a lowering of the temperature. Further, that under these conditions four atoms of oxygen was the maximum used per mol of sodium tartrate; or, at least, the number of atoms of oxygen used per mol of sodium tartrate approached four as a limit, which meant that the tartrate was oxidized toward oxalate. At room temperature, however, the reactions became distinctly autocatalytic and much slower. The induction period lasted for several
INDUCED REACTIONS
1603
hours, the speed of the reaction gradually increased until nearly two atoms of oxygen were used per mol of sodium tartrate and then began to decrease. I n other words glycollic acid was the turning point in the speed of the oxidation. The fact that the oxidation of tartaric acid by chromic acid was autocatalytic in nature was thought to be in some way connected with the fading end-point previously obtained with iodine titrations in partially oxidized tartrate solutions. Further, since the oxidation of tartaric acid by bromine was autocatalytic as shown by Bunsen and Roscoe,' it seemed that the explanation must lie in the oxidation products of the tartrate. The catalytic
Fig 2 a, b and c
effect of glycollic acid, oxalic acid, formic acid and chromic sulphate was tried on the oxidation of tartrate using chromic acid as the oxidizing agent instead of iodine as previously tried. The excess chromic acid was dcterK2Cr207was added to a mixture mined after five hours when 2 j cc of X'IOO of 5 cc 3 1 / 2 0 Sa2CrH10s, I O cc 4 9 H2S04and 0-10 cc W I Osupposed catalyst. In no case however, was the amount of chromic acid used changed by a variation in the amount of supposed catalyst, which of course eliminated them as possible autocatalysts for this reaction. I t was known from the researches of Harcourt and Esson2 and later by SkrabaF that in the oxidation of oxalic acid by potassium permanganate manganous sulphate was found to be the autocatalyst. Hence, could chromous sulphate, the corresponding salt in the above reaction, be the autocatalyst? This was investigated by reducing a M ' I O solution of chromic sulphate in a Jones reductor, then testing the catalytic effect of the chromous sulphate thus formed on the oxidation of tartrate by chromic acid.-It should perhaps be mentioned that chromous sulphate is unstable in air and is gradually oxidized back to chromic sulphate. It was therefore possible to obtain the effect of decreasing chromous sulphate concentration by making runs at different time intervals after the chromic sulphate had been reduced. LOC.cit. Harcourt and Esson: J. Chem. SOC.,20, 460 (1867). Skrabal: Z. anorg. Chem., 42, I (1p04).
I 604
WESLEY G. VAXXOY
rr.UiLE IT' Mixture put thru Jones reductor in order mentioned; I O O cc ?; H&04, so cc LLI:/Io Cr2(S06)3,zoo cc S H2Y0, and IOO cc H20. Total vol. 4 j o cc. Acidity of resulting chromous sulphate solution determined by titrating with standard sodium hydroside--.4 Sormal. Reaction RIist,ure: 50 cc S '100X2Cr207> 5 cc M;'IOOXanCcIl,Osand z cc of chromous sulphate solution taken at various time intervals after the reduction as indicated. Temp. z z 0 C (Room) 40 cc S,'IOO Ii2C'rnOi= 40 ec 1;I O O S a 2 S 2 0 3 Pt. I Time after reduction, 5 min. 2 cc C-'rS04 = 2 . 7 cc. B2Cr?Oi SO.
Time in hrs.
St12&03Titre
2.9 3.75 5.3
4 9
Pt. z so.
Ii2c1r207used
?2
IO.j
33
1.1.8
T m e after reduction, about Time in hrs.
12
hrs.
2
cc C r S 0 4 = .1 cc Ii2Cr.0;
> a 2 s 2 0 , Titre
Ii2Cri0, used
49 4
0 4
I
I
2
I2
3
23
7 . 2
4
36 45
12.7
5
Pt. 3 so. I
2 9
Ij.6
Time after reduction, about 48 hrs. Time in hrs. 4
8 24 33
48
2
cc CrSOc = Occ IGCr?Oi
Sa2S103Titre
50 49.6 47
I
