T H E DECOMPOSITIO?; O F OXALIC ACID BY ARSEXIC ACID* BY G. BROOKS KISG AND JAMES H. WALTOX
The study of the decomposition of oxalic acid by arsenic acid was made in order to compare the dehydrating properties of arsenic acid with those of sulfuric and phosphoric acids. The decomposition of oxalic acid by sulfuric and phosphoric acids has been studied by Lichtyl and Kalton and TVeber' respectively, who have shown that the reactions follow the unimolecular reaction law, and are strongly inhibited by water. I t was of interest to determine whether or not arsenic acid would decompose oxalic acid similarly. RIenzies and Pottel.3 have shown from a phase rule study of the two component system, As20j and H 2 0 , that AsZO5.4H2Oand 3XszO~.jHzO are the only hydrates of As?Oj which can be isolated. K h e n heated, As~Oj.4H20 fuses a t 36.1' and rapidly changes to 3As20j.5H20. No indications of As205. 3 H z 0 (H&04) were obtained. In the present work solutions of ASZOSin xvater were used for the decomposition of oxalic acid.
Experimental Part Preparation and Purification of Materials-The arsenic acid used was the C.P. quality of General Chemical Co. Solutions of various concentrations were made by dissolving the solid in water, and the acid content determined by the titration method developed by Menzies and P ~ t t e r .This ~ method was checked against a gravimetric analysis and found to be rery satisfactory. C.P. sodium oxalate was dried in an oven a t 140' for several days, and thereafter kept in a desiccator. Sodium oxalate was used rather than oxalic acid as it is easier to obtain in the anhydrous state and because a t the temperatures used, the oxalic acid would sublime to an appreciable extent. Apparatus and Method of Procedure-Preliminary experiments showed that CO and COZ were present in the reaction products from the decomposition of the sodium oxalate by arsenic acid. The rate of reaction was followed by measurement of the volume of gas evolved at various time intervals during a decomposition, the apparatus employed and the method of procedure being essentially that used by ITiigj4Khitford,j and Dittmar6 in their studies on the decomposition of organic acids by sulfuric acid. I O cc. of arsenic acid solution were taken for each decomposition, and at the time of adding the acid to the reaction flask, a sample was taken in a weighing * Contrihution from the Laboratory of General Chemistry of the University of Wisconsin.
* Lichty: J. Phys.
Chem., 11, 2 2 5 (190;). Walton and Weber: Unpublished Results. 3 Menzies and Pot,ter: J. Am. Chem. Sac., 34, 1462 (1912). Wiig: J. Am. Chem. Soc., 5 2 , 4729-4751 (1930). 5 Whitford: J. Am. Chem. Sac., 47, 9 j 3 (1925). 8 Dittmar: J. Am. Chern. SOC.,52, 2j46 (1930); J. Phys. Chem., 33, j 3 3 (1929).
G. BROOKS KING AND JAMES H. WALTON
2378
pipette for subsequent analysis for arsenic content. A qualitative test on the decomposition products showed that no reduction had taken place in the reaction.
Results Completeness and Order of Reaction-The decompositions were carried out a t temperatures of 140' and 150". The concentrations of arsenic acid which could be used were limited; with solutions containing less than 7 0 % As206, water condensed in the upper part of the flask and in the capillary tubes, thereby concentrating the acid solution and hindering the flow of gas, thus making the results unreliable. At concentrations above 78% A s z O ~a, white solid separated on shaking, undoubtedly some of the arsenic acid precipitating from a supersaturated state. The results given in Table I show that the reaction is a quantitative one.
TABLE I Temperature of run Theoretical volume of gas Volume of gas found
140' 34.0 33.5
1500
35.7 35.6
1500 34.4 34.3
The decomposition of oxalic acid by arsenic acid follows the unimolecular reaction law, as shown by the following velocity constants obtained when the data for any one decomposition were substituted in the equation for a reaction of the first order: 31.0; 33.6; 3 2 2 ; 30.4; 32.0; 31.3; 30.8; 31.1. That the true speed of decomposition and not the speed of solution was measured was shown by the fact that the velocity constants were not affected by the volume of arsenic acid solution. j , 7, and I O cc of the acid solution were used and the deviation from a constant value for k was within the limits of experimental error. The value of k x 103 for any particular run was taken as the average value between 30 and 70 per cent decomposition in from two to four determinations. The Eflect of Water on the Reaction Rate-The data obtained a t the temperatures 140' and 150' are summarized in Table 11, and represented graphically in the curves, Fig. I .
TABLE I1 Effect of Water on Reaction Velocity
70Water 9.54 6.64 5.58 j.16 4.90*
At 140' Mol of H1O
5.85
3.95 3.28 3.02 2.86
k X 103 30.2 63.2
* Some separation of solid.
71.0
85.2 84. I
C&
Water
10.34
A t 150' Mol of Hz0
10.22
6,39 6.33
8.60 7.75 4.03 1.2*
4.67 3.64 0.67
5.22
k X 103 116 123 I44 I74 220
320
c
DECOMPOSITION O F OXALIC ACID BY ARSENIC ACID
2379
From the curves in Fig. I , the temperature coefficients of the reaction have been calculated and are given in Table 111.
TABLE I11 Molality of water Temperature coefficient
2.0
3.0
4.0
5.0
2.7
3.06
3.48
3.90
FIG.I Effect of m t e r on reaction rate a t I . 140' Subtract 1.00 from log K X 2. 150'
103scale
Although the effect of the addition of very small amounts of water can not be ascertained from these data, it is evident that mater has a marked retarding effect on the velocity of reaction. Similar results were obtained by TTiig,4Whitford,j Dittmar? and Lichty' in their studies on the decomposition of organic acids by sulfuric acid, also by Walton and Weber2 in the study of the decomposition of sodium oxalate with phosphoric acid. The temperature coefficients vary considerably with change in molality of water, howerer the change is regular and in accord with results obtained
2380
G. BROOKS KING A S D JAMES H. WALTON
by other investigators in similar decompositions, showing an increase as the concentration of water increases which is in accord with Dhar’s7 statement that a negative catalyst increases the temperature coefficient of a reaction. Solubzlity o j Oxalic Acid in Solutions o j Arsenzc Aczd-It has been shown by Dittmar,KWiig4 and others that the decomposition of organic acids with sulfuric acid may take place as a result of a preliminary reaction between the two components to form an oxonium salt. If the decomposition of oxalic acid with arsenic acid is similar t o that with sulfuric acid, then indications of
FIG.2 Solubility of HE104 in solutions of HsAsOd at 25’.
compound formation between arsenic acid and oxalic acid might be obtainable. Knox and Richards* have shown that solubility determinations furnish a convenient method for such a study. Their method consists in determining the solubility of one acid a t a particular temperature in solutions of increasing concentration of the second acid. Indications of compound formation are obtained by a minimum in the solubility curve. I n the present study, several solutions of arsenic acid of various concentrations were prepared. An excess of oxalic acid mas added to each solution which was then kept in a thermostat for several days t o reach equilibrium. Samples were then withdrawn, the oxalic acid titrated with standard potassium permanganate, and the total acidity determined by titration with standard alkali. The difference between the total acidity and that due to oxalic acid gave the acidity due to arsenic acid. The results obtained are summarized in Table IT’ and the curve for these data is plotted in Fig. 2 . Dhar: J. Phys. Chern., 28, 951 (1923). *Knox and Richards: J. Cheni. SOC.,115, j08 (1919).
7
DECOVPOSITION O F OXALIC ACID BY ARSENIC ACID
TABLE Iv Solubility of Oxalic Acid in Arsenic Acid Sorrnality of solvent
Normality of solute
0 0
2.412
2.2
2.20
5.41
1.92
j.Oj
1.77
13.6 17.7 23.8
0.8j8
? j .0
0 .j 1 6
36.3 39 i
0.96j
1.43 I
.08
1.72
The data are expressed in normalities of solute (H2C204)and solvent (H3As04). A distinct minimum appears in the curve, indicating compound formation between the arsenic and oxalic acids.
Discussion Illechanism of flw Deconqmition and Inhibition--The
decomposition of osalic acid by arsenic acid is probably best explained by assuming the formation of an unstable oxalic acid-arsenic acid complex, the concentration of which will govern the speed of reaction. The inhibiting effect of mater on the reaction may be explained on the basis of Taylor's9theory of ncgotive catalysis in which it is postulated that the added substance decreases the concentration or the active mass of one or more of the reactants through compound formation. I n the present case, water would be assumed to form rnolecular complexes with either arsenic acid or O X R I ~ Cacid or with both. The hydrates of these acids are, of course, ne11 known.
Coinparison of Arsenic, Sii(fitric a n d Phosphoric dcirls-By estrapolation to o concentration of water in the osalic acid decomposition by arsenic acid a t I j o o , a value of approximately 500 is obtained for k >( 10~.If-alton and If-eber' obtained 30; as a value for k X 103 using pure ortho phosphoric acid much higher value than either of the above is obtained for pure a t I jo'. sulfuric acid at, 2 5 ' . The acid strength of the acids as measured by their dissociation constants follow: Oxalic 3.8 X I O - ? ; Phosphoric 1.1 X IO-?, .Irsenic 0.j X IO-?. If the tendency for molecular compound formation governs the speed of reaction, then on the basis of I
