July, 1935
OXIDATION STATEOF ARGENTIC SALTS
(21.2 N , 90% nitric acid, density 1.49) gave, however, the striking result that, as the acid becomes nearly anhydrous, scarcely any argentic silver is produced-after seventy-five hours not even enough to give more than a slight brown color to the solution. The table further shows that the percentage of argentic silver produced a t any given acid concentration decreases, though its absolute concentration increases, when the total silver concentration increases. It is seen that a reagent 0.2 formal in Ag(NOJ2 can be prepared in 10 or 15 N nitric acid.
Summary
It has been shown that a t 0 and 24' ozone slowly oxidizes argentous nitrate in nitric acid solution to a black soluble argentic nitrate; but that the reaction ceases before i t has become complete, with the establishment of a steady state. This steady state is shown to arise, not at all from the occurrence of the reverse reaction, but from the reduction of the argentic salt by the water with evolution of oxygen. (For the overall reactions see the equations a t the beginning of this paper.) The argentic salt is assumed to consist mainly of bipositive silver on the basis of the results of the next following paper. The initial rate of the silver oxidation at 0' was found to be proportional to the argentous and
[CONTRIBUTION FROM
1229
ozone concentrations. This is accounted for by the following mechanism which involves the direct production of tripositive silver in the form of an oxygenated ion
+
+
Ag+ 0 3 = Ago 0 2 (slow), and A g o + A g + 2" = 2Ag++ Ha0 (fast)
+
+
+
+
The initial rate is found to be only slightly affected by the acid concentration or by the ionic strength. The rate of the argentic decomposition by water (at 0 and 24') was found as a first approximation to be proportional to the square of the argentic concentration (except when the latter is high), and inversely proportional to the argentous concentration, and to be greatly diminished by increasing the acid concentration. The mechanism is discussed (see equations 13, 14, and 15); but complications in the reaction make the conclusions somewhat uncertain. The effect of temperature on the reaction rate was found to be very much larger than for most reactions. The steady state a t 0' was roughly determined a t various acid concentrations. The percentage of the silver converted into the argentic form varied from 18% in 1.8 Nnitric acid to about 90% in 12 N nitric acid when the total silver was about 0.04 formal, but in 21 N (90%) nitric acid scarcely any formation of argentic salt resulted. PASADENA, CALIF.
GATESCHEMICAL LABORATORY, CALIFORNIA
INSTITUTE OF
RECEIVED MAY8, 1935
TECHNOLOGY, NO. 477 I
Argentic Salts in Acid Solution. 11. The Oxidation State of Argentic Salts BY ARTHUR A. NOYES, KENNETH s.PITZERAND CLARENCE L. DUNN Purpose of this Investigation In the preceding article Noyes, Hoard and Pitzer' described the general nature and the rates of the reactions by which argentic salt is produced by the ozone oxidation of acid argentous nitrate, and by which it is reduced by water. Before proceeding further with the description of these researches, however, it seemed essential to establish the state of oxidation, whether bipositive or tripositive, in which argentic silver exists in nitric acid solution. To a description of experiments made with this purpose this article is devoted. It is remarkable that this question has re(1) Noyes, Hoard and Pitzer, THISJOURNAL,
s?, 1221 (1935).
mained unanswered so long in the case of a readily obtainable oxidation state of a common element-an oxidation state, moreover, that explains the striking catalytic effects of silver salts. Furthermore, the question is of much theoretical interest in relation to the valence states cornpatible with the periodic system, inasmuch as copper exists mainly unipositive and bipositive, and gold unipositive and tripositive. Previous Investigations The most definite indications previously existing as to the oxidation state of argentic silver in its solutions were furnished by the composition of the solid phases that separate in preparing or
1230
ARTHURA. NOYQS,KENNE~H 8. P~TWR AND CLARENCE L. DUNN
in diluting the sdutions or upon adding other subtandes to them, but these indication$ seem contradictory. First, it may be mentioned that Barbieri2 found upon mixing an aqueous solution of potassium peroxysulfate with a solutiori of silver nitrate in pyridine (Py) that orahge-colored rrqfstal6 separated of the composition
[email protected], and that this compound forms isonWphous crptals with the pk'eviously p r e p a M cupric mmpound Cu&eOs*4Py,both these facts Showkg the existence of a complex salt caritalning bipositive silver, as is more clearly expressed by tRe formula [Py4Ag]+*&OB'. There have siriee been prepared3 complex salts of bipositive d v e r with various other bases (having coordinaticn tldinbers four or six) and with other anions than peroxysulfate. On the other hknd, s e e r a l investigators ( R i g - ~ o nY , ~ ~ s t ,Barbieri,6 ~ Jirsa') habe shdwn that argentic solutions give prwipitates w th water of solid basic salts that contain tfipositive silver. This is proved by the fact tkat they yEeld on drying cornpourids such as AgON03. Agd&*4AgOor (AgO)2S04.6AgO. Such sa€ts are precipitated even froivi a solution prepared by dkmlving pure silver oxide in nitdc acid, and even though this di&olt.ing is attenaed by evolution of oxygen (Jirsa). The results obtained by Yost5 and by Carmans are especially significant, in that they show that all the silver is originally precipitated in the tripositive state. Thus the compounds prepared by them from neutral potassium peroxysulfate and silva- nitrate contain per atomic weight of silver 2.0 oxidation equivarents (measund with referace to reduction t o argentous salt). Carman showed, however, that if the preripitates stood 6fie to three hours before filtering, the ratio drops t o a fairly constaht value of 1.51-1.53, correspond&g to about equal quantities of 'bipositi!ve and tpipositive siher. The faCt that the rate of oxidation of &rgentous salt by peroxysulfate, well as by ozone, (2) Barbieri, R t t r accad Lincer, [51 21, 560-503 (1912). Gam ch+m rtal , 42, 7 (191h)
(3) Hieber and Muhlbauer, Ber , 61, 2149 (1928). Morgan and Burstall. J Chem Soc.. 2B94 (1930), 8Sab&i, Ber , 60, 24242427 (1927), and A f f raccad Lincei, [e] la, 44-47 (1932). (4) HtgBon. J. Chem. Soc., 119,2048(1921) (5) Yost, THISJ O U ~ N A L ,I S , 162-164 (1920) I%, 88Q-bd7 (1931) (6) BarGiefi, k'& accad L h O i , f8] ( 7 ) Jirsa, Z anorg Chem , 148, 130-144 (1925); 168, 33-80, 01-05 ('1920) (8) Cdrfnan, Trahs Faraday Soc , k,606-577
