638
GEORG CRONHEIM
THE CAT.4LYTIC ACTION OF NATURAL MINERAL WATERS. I1 GEORG CRONHEIM1 T h e N e w Yo?k State Rerearch Instztute, Saratoga S p a , Saratoga S p r i n g s , N e w York Received July $6, 1949
In a previous paper (I), the catalytic decomposition of hydrogen peroxide by natural mineral waters has been described and the mechanism of the process has been explained by means of a chain reaction. In order to come to a better understanding of the physiological effect of natural mineral waters, the investition of their catalytic activity was extended to so-called “peroxidase” reactions. It is known that ferrous salts act as catalysts in the oxidation of many organic compounds by hydrogen peroxide (Wieland), This effect resembles qualitatively the action of the enzyme peroxidase and, in the case of natural mineral maters, is sometimes referred to as their “peroxidatic” activity. In the present studies different types of natural mineral waters from Saratoga Springs, i’iew York, were used. These waters are mainly solutions of the halides and bicarbonates of alkalies and alkaline-earth alkalies, supersaturated with carbon dioxide. In addition they contain small amounts (4-13 mg. per liter) of divalent iron and traces of a number of other metals and acids.2 The “peroxidatic” activity of the mineral waters toward the oxidation of formic acid was investigated. This acid belongs to a group of biochemically important compounds, is quite stable, and can easily be determined. Furthermore, its oxidation products do not introduce any new compound into the whole reaction system. A series of typical experiments is presented in table 1 . It shows that the free formic acid is completely oxidized if both mineral water and hydrogen peroxide are present. Without mineral water, the formic acid is oxidized only a t a negligible rate. Nevertheless, it is pot a simple bimolecular reaction, HCOOH
+ HzO2 = 2H2O
because the experimental results give no constant value for suggested, therefore, the following chain reaction:
+ H20z= Fe+++ + OH- + OH + HCOOH = COOH + Hz0
Fe++ OH
+ C02 + H20 OH + HCOOH = OH + C02 + H2O COOH
+
H202 =
OH
. . . . . .. . . . . .
(1) K,,in,ol.
Weiss ( 5 ) (2)
(3) (4)
(5)
Present address: The G. F. Harvey Company, Saratoga Springs, New York a complete analysis of these waters, Bee 0. Baudisch and D Davidson (Arch. Internal. Med. 40,507 (1927)) and the publications of the Saratoga Spa. 1
* For
689
CATALYTIC ACTION OF NATURAL MINERAL WATERS. I1
The chain is broken by
COOH
+ OH = HzO + COz
(6) According to Weiss, the length of the chain depends upon the average acid and ferrous-ion concentrations. Therefore, the oxidation of formic acid should stop very soon, since all ferrous ions are quickly oxidized by the excess of hydrogen peroxide. The experiments, however, prove that even after 72 hr. the oxidation of formic acid still proceeds at a measurable rate. Furthermore, if new formic acid and TABLE 1 Decomposition of formic acid and hydrogen peroxide b y different natural mineral waters 200 ml. of mineral water 20 ml. of normal formic acid 60 ml. of hydrogen peroxide (1.5 per cent) or distilled water; temperature, 37°C.; 10 ml. each of the reaction mixture titrated with N/20 sodium hydroxide or N/5 potassium permanganate
+
TIME
haws
1
+
COESA WATER
OEYSEB WATEE
Ha01 added
ml.
I Ho0 added
HI& added
ml.
ml.
I HnO
HATEORN WATEP
added
ml.
ml.
4.
d.
15.25 15.05 14.90 14.85 14.80 14.80
0 24 48 72 96 120
5.50 3.00 0.60 0.10 0.05
7.65 7.55 7.50 7.50 7.45
5.90 3.60 0.75 0.30 0.20
8.00 7.90 7.85 7.80 7.75
1.75 1.30 0.80 0.60 0.40
4.85 4.85 4.80 4.80 4.75
0
10.35 7.60 5.20 2.50 2.15 1.80
0
10.35 7.65 5.40 2.90 2.45 2.10
0
10.35 4.80 4.05 3.45 2.90 2.40
0
24 48 72 96 120
D1sT'L*D WATEP
10.35 10.30 10.20 10.10 10.06 10.05
hydrogen peroxide are added, the reaction again continues until the second and even a third portion of the acid is oxidized. In order to explain this discrepancy the above chain mechanism has to be completed by a group of reactions in which ferric ion takes part. From observations by Kuhn and Wasserman (3) and Simon and coworkers (4), it is known that trivalent iron can be reduced quantitatively to divalent iron by hydrogen peroxide in acid solution.
HzOz ;rt H+ 3- IIOpFew HOI = Fe* HO2
+
+
(7) (8)
640
GEORG CRONBEIM
Thus, ferrous ions necessary to start a chain reaction according to reaction 2 are regained. Furthermore, the radical HOz formed in reaction 6 is able to start another chain according to HOe
+ HCOOH = COz + OH + H,O
(9)
The basic reaction in th$ new extended mechanism which regulates the oxidation of formic acid, is apparently the equilibrium 6 and it is determined by two factors: the concentration of hydrogen peroxide and its dissociation as indicated by the pH of the solution. TABLE 2 Influence of the concentration of hydrogen peroxide upon the catalytic oxidation of formic acid 200 ml. of Hathorn water 30 ml. of normal formic acid 60 ml. of hydrogen peroxide of different concentrations; temperature, 37°C.; 10 mi. each of the reaction mixture titrated with N / 2 0 sodium hydroxide or N/5 potassium permanganate
+
TME
-48
CONCENTRATION OF ADWED HSWROGEN PEROXIDE
I
1
+
0.3%
1
1.5%
1
3.0%
1
7.5%
mi.
mi.
12.05 10.90 10.30
10.90
120
9.45
1.25 0.75
0 24 48 72 96 120
1.95 1.30 0.75 0.40 0.15 0.10
9.60 7.50 5.70 2.50 0.45 0.30
I
mi.
mi.
9.85 5.20 0.40
5.90 2.25
ii ~
~
1 ~
I
21.05 16.25 11.20 5.05 0.20
49.90 40.25 29.00 0.10
The experimental proof for the first part of this statement is summarized in table 2. It shows clearly that the oxidation of formic acid depends directly upon the concentration of hydrogen peroxide. The importance of the dissociation of hydrogen peroxide can be shown only indirectly by the following consideration: In the course of the gradual oxidation of free formic acid the pH of the solution increases constantly, thereby shifting the dissociation equilibrium 7 more and more to the right side. This in turn increases the rate of reaction 8 and thereby of the whole chain mechanism. In other words, one has to expect that the oxidation of formic acid proceeds at an ever-increasing rate. This has been verified by the experimental results, as shown in table 1.
641
CATALYTIC ACTION OF NATURAL MINERAL WA1IERS. I1
The fact that changes of the pH even in a rather acid medium really deet the dissociation of hydrogen peroxide to a measurable degree is proven by a series of experiments where formic acid was replaced by various amounts of hydrochloric acid. The reaction studied in this case is the catalytic decomposition of hydrogen peroxide where, according to Cronheim (l), the dissociation equilibrium 7 is of the same importance as in the oxidation of formic acid. Since hydrochloric acid is not affected by either mineral water or hydrogen peroxide, the pH remains constant throughout each experiment. The results summarized in table 3 prove the correctness of these statements. The decomposition of hydrogen peroxide is dependent upon the pH of the mixture, but for a given pH it proceeds at an almost constant rate. The foregoing considerations and experiments leave no doubt that even in “peroxidatic” reactions not the absolute amount of ferrous or ferric ions is of TABLE 3 Decomposition of hydrogen peroxide at different pH
+
+
200 ml. of mineral water (HathornNo. 2) 60 ml. of hydrogen peroxide (3 per cent) 15,21, 26, and 30 ml. of normal hydrochloric acid; temperature, 37’C.; 5 ml. of the reaction mixture titrated with N/10 potassium permanganate; K is calculated from the equation for monomolecular reactions PH OF
1
5.78
MIXTURE.
I
5.40
2.88
2.48
Zd$, -ml. K X iW
TIME
hours
MI.
0 24 48 72 96 120 144
20.40 18.60 15.95 13.50 11.35 9.40 7.85
ml.
1.63 2.75 3.02 3.13 3.37 3.35
20.40 17.95 15.40 13.05 10.95 9.00 7.35
ml.
2.21 2.75 3.00 3.13 3.53 3.67
20.40 17.40 14.25 11.50 9.20 7.25 5.70
2.84 3.60 3.86 4.05 4.31 4.31
20.40 16.10 12.20 9.20 7.05 5.30 4.05
4.30 5.04 5.10 4.82 5.15 5.24
predominant importance for the catalytic activity, but the conditions which regulate their respective reactions with dissociated or undissociated hydrogen peroxide. Simultaneously with the oxidation of formic acid a catalytic decomposition of hydrogen peroxide occurs in the system according to the mechanism outlined in the previous paper. The rate of this decomposition as determined by titration with potassium permanganate (tables 1 to 3) is rather slow, owing to the acidity of the mixture. Compared with the above-mentioned mixtures with hydrochloric acid, it is, however, somewhat greater in systems containing formic acid, owing to an additional effect of the catalytic and peroxidatic reactions. A quantitative study of this effect, which in itself is insignificant, leads to some quite interesting results. Table 2 contains figures for the simultaneous decompositions of formic acid and hydrogen peroxide for different initial con-
642
GEORG CRONHEIM
concentrations of the latter component. The ratio (decomposed hydrogen peroxide)/(decomposed formic acid) is approximately 1, if both constituents are present in equimolar concentrations. If the concentration of hydrogen peroxide is increased, this ratio increases too, but at a lower rate. An explanation of this fact is given by the folloiving consideration: The first step for both chain mechanisms is reactlion 2 betmen ferrous ion and hydrogen peroxide with the formation of the radical OH. This radical can continue the chain with either hydrogen peroxide or formic acid, the probability depending upon their relative concentration. If the concentration of hydrogen peroxide is increased, its chances for the reaction
OH
+ HzOz = HOz + HzO
(10) increase accordingly. I n order that the radical HOz can continue the chain, it must first dissociate:
H0z e Hf
+ 02-
(11)
Since this equilibrium in the present acid mixture is shifted most,ly to the lefthand side, there i-iill be a relative accumulation of HO? radicals in the system, resulting in a greater probability for one of the folloiring reactions:
HOz
+ OH
+ COOH HOz + H0z
HOz
= = =
HzO
+
0 2
+ COz HzOz + 02 HzOz
(12)
(13) (14)
I n each case, the chain will be broken. The fact that the radical €IO2 is formed only by the decomposition of hydrogen peroxide explains satisfactorily the above-mentioned observation that the rate of decomposition does not increase proportionally to the concentration of hydrogen peroxide. The importance of the dissociation equilibrium 7, as outlined in the previous paragraphs, is due to the fact that it affects not only the length of the chain by its direct participation in several of the chain reactions but that it also determines the number of chains through reaction 8. The,only other way by which an alteration of the whole reaction mechanism seems possible is through factors influencing the transition of the iron ions between the di- and tri-valent states. In order to study this possibility, small amounts of various metal salts irere added to the reaction mixture. All the metals selected exist in several valence states and therefore might be able to influence the oxidation-reduction equilibrium of the iron salts. They included copper(I1) chloride, nickel(I1) chloride, cobalt(I1) chloride, manganese(I1) acetate, lead(I1) acetate, and palladium(I1) chloride. The experiments were carried out in the described manner with bottled water from the spring Hathorn No. 2, which contains 12 mg. of divalent iron per liter. The results of some typical experiments are shown in table 4. It can be seen that Cu++, Mn++, and to a small extent Pd++ ions have an
643
CATALYTIC ACTION OF NATURAL MINERAL WATERS. I1 \
TABLE 4
Influence of diflerent metal salts upon the decomposition offormic acid and hydrogen perozide b y natural mzneral water 200 ml. of mineral water (Hathorn No. 2) 20 ml. of normal formic acid 60 ml. of hydrogen peroxide (3 per cent) 10 ml. of M / W salt solution; temperature, 37°C.; 10 ml. each of the reaction misture titrated with N/20sodium hydroxide or N/5 potassium permanganate ADDED SAL1
.
1
CUPRIC CHLORIDE
+
+
+-
I zzEI I COBALI ACETATE
MANGANESE ACEIATE
I
LEAD ACEPAIE
I
PALLADIUM CHLORIDE
1
NONE
TIME
hours
24 48 72
0 24 48 72
1 1 I ml.
6.20 0.05
19.65 11.45 1.80
1 ~
Titrated with N/20 NaOH ml.
ml.
ml.
ml.
7.40 1.45 0.05
7.10 1.85 0.20
5.35 0.40
7.25 1.45 0.05
~
19.65 13.70
19.65 13.45
;.g ;:,ll
~
I 1 19.65 12.00
19.65 13.65
5.70 0.05
1
ml.
Id.
6.10 0.65
7.40 1.50 0.05
19.65 12.85 5.65 0.55
1
19.65 13.65 6.70 2.55
TABLE 5
In$zience of cupric salts upon lhe deconiposilion of formic acid and hydrogen peroxide by natural mineral water 200 ml. of mineral watcr (Hathorn No. 2) 20 ml. of normal formic acid 60 ml. of hydrogen peroxide (3 per cent) cuprlc chloride; temperature, 37°C.; 10 ml. each of the reaction mixture titrated with N/20 sodinni hydroxide or N/5 potItssiuni permanganate
+
+
+
- 0 24 48 72
21 10 2 10
21 IO 13 50 3.80
21.10 15.30 6.95 0 85
21 10 16.30 11.25 5.10
644
GEORG CRONHEI?d
Since small amounts of copper are supposed to be necessary for the normal oxidative function of iron in hemoglobin, a more quantitative investigation of its influence was made (table 5 ) . The figures shorn an increase in the catalytic activity with even such a small amount as 10 micrograms of cupric ions added. With 1.0 mg. of cupric ions the rate of decomposition is increased about four times. Since cupric ions alone are without effect, the result is due to the combined action of iron and copper. In the previous paper (1) it was stated that the absolute rate of the catalytic activity of natural mineral waters is greatly affected by secondary processes between intermediate reaction products and components of the mineral waters. Further proof for this statement is given in the following experiments. TABLE 6 Decomposition of f o r m i c acid and hydrogen peroxide b y deposit f r o m natural mineral waters Deposit from 10&400 ml. of mineral water (Hathorn No. 2) 15 ml. of normal formic acid 60 ml. of hydrogen peroxide (3 per cent) 200 ml. of distilled water; temperature, 37°C.; 10 ml. each of the reaction mixture titrated with N/20 sodium hydroxide or N/5 potassium permanganate RESIDUE FBOYI..
+
+
+
1 I
. . .. . ..
TIME
200mi.
I
+ 0.1
400 ml.
200 ml.
mg. CuCln
Titrated with N/20 NaOH hours
0 5 24
I
ml.
11.10 ~
9.80 1.20
1 ~
ml.
11.10 9.50 0.60
1 ~
ml.
11.05 7.15 0.10
I ~
mi.
11.10 7.40 0.25
Titrated with N/5 KMnOI ~~
0 5 24 48
21.50 20.25 10 50 0.90
21.50 19.50 8.40 5.65
21.50 17.15 6.80 4.35
21.50 18.40 2.85
SOLUBILITY OF AMMONIUM PERSULFATE
*
645,
results are another proof that although iron salts are absolutely necessary for the studied reactions, their concentration is only of secondary importance. SUMMARY
The catalytic oxidation of formic acid by natural mineral waters and hydrogen peroxide was studied. The oxidation is started by the ferrous ions present in the mineral water and is continued by a chain mechanism. The main factors influencing this chain mechanism are the concentration of hydrogen peroxide and its dissociation as indicated by the pH of the solution. The addition of cupric salts greatly increases the rate of the reaction. REFERENCES (1) (2) (3) (4) (5)
CRONHEIM, G . : J. Phys. Chem. 46, 328 (1941). KRAUSE,A., AND KANIOWSKA, D.: Ber. 6% 1982 (1936). A.: Ann. 609, 203 (1933). KUHNR , . , AND WASSERMANN, SIMON, A., AND HAUFE,W.: 2.anorg. Chem. 230,148,160 (1936). WEISR,J.: J. Phys. Chem. 41, 1107 (1937).
SOLUBILITY OF AMMONIUM PERSULFATE IN WATER AND IN SOLUTIONS OF SULFURIC ACID AND AMMONIUM SULFATE J . F . GALL, G. L. CHURCH, AND RALPH L. BROWN Research Division, Pennsylvania Salt Manufacturing Company, Philadelphia, Pennsylvania Received July 66, 194.9 INTRODUCTION
In spite of the growing importance of per compounds, much of the physical data relative to even the inorganic members is lacking. Thus we have found only two references to the solubility of ammonium persulfate (1, 2), and these report results of single measurements only. We have had occasion to measure the solubility of this salt in water and in solutions of sulfuric acid and of ammonium sulfate; the results obtained are presented in this paper. EXPERIMENTAL
Solutions of sulfuric acid and ammonium sulfate were prepared from C.P. reagents. Ammonium persulfate of the same quality was finely pulverized and added in considerable excess to the solutions in brown-glass rubber-stoppered bottles. Continuous agitation at constant temperature was obtained by tumbling in a water thermostat held within &0.5OC. of the desired temperaturs. In most cases two saturation bottles were used at each composition, and samples were withdrawn a t intervals of 1 hr. or 4 hr. Analysis was made for persulfate,
