Catalytic and Induced Reactions in Microchemistrv J
J
I. M. KOLTHOFF AND R. S. LIVINGSTON, University of Minnesota, Minneapolis, Minn.
LTHOUGH catalytic and oxidation of thiosulfate by hyThe general theory of catalysis and inducinduced r e a c t i o n s are tion in homogeneoussystems is discussed. drogen peroxide. On the other widely used in the detechand, we could predict that the Induced and catalyzed tion and quantitative estimaprecipitations are chlorine-chloride pair would not mentioned briefly, tion 3f c e r t a i n ions or comby a general be an efficient catalyst for this discussion of analytical applications. Fireaction, since the rate of oxidapounds, not much theoretical nally, the catalytic and inducing effects of tion of chloride ion by hydrogen work h a s been done on the new tests are peroxide is known to be slow. mechanism of most of these resilverand mercury and described, Unfortunately, such information actions. It is rather deplorable that most of our knowledge in about the rates of the reaction this field is empirical. If the steps idvolved is commonly not theor:y underlying catalytic and induced reactions were known available and a consideration of the free energies of the probmore explicitly, it might be expected that a systematic applicaable reaction steps is the only systematic guide which we can tion of the theoretical fundamentals would result in the disuse in selecting catalysts. While this method does not enable covery of many useful reactions for the qualitative detection one to predict that any possible catalytic pair will be an efficient and quantitative estimation of microquantities of substances. catalyst for a given reaction, it does greatly simplify the probHomogeneous catalysis of oxidation-reduction reactions lem by excluding many substances which might otherwise b e occurs when the catalyst can exist in an oxidized and a reconsidered as possible catalysts. duced form and when these forms of the catalyst are capable As an example of this method, let us consider the oxidation! of reacting rapidly with the reducing and oxidizing agents, of arsenious acid by ceric ion in an acid solution. Since t h e free energy of a reaction is the algebraic sum of the free enerrespectively. A simple and well-known example of this type of reaction is the catalysis by the iodine-iodide couple of the gies of its constituent half-cell reactions, the free energy d a t a oxidation of thiosulfate ion by hydrogen peroxide in slightly may be conveniently summarized by plotting the half-cell acid solution. The kinetics and stoichiometry of this reacpotentials (4) of reactants and possible catalysts against pH tion may be summarized by the following chemical equations: (Figure 1). It should be emphasized that this plot is not, intended to i m d v that any electrochemical mechanism of (1, rate determining) catalysis exists,*it is merely a convenient method of repreHzOz I-+ IO- HzO LO-$-H+Ft.HIO (2, rapid, reversible) senting free energy or equilibrium data. I n preparing t h e HIO-t-I-+H+~Iz+HzO (3, rapid, plot no attempt has been made to allow for secondary effects Iz+I-*Is(4,rapid, reversible) (5,rapid, probablycomplex) of pH-i. e., the electromotive force of the Ce++++,C e + + + r8- + 2s20a--++ s40s= + 31HzOz -t 25203= 2H + = S4Oe- 2Hz0 (6, stoichiometric) I
-+
+
+
+
Equation 6, which represents the stoichiometric reaction, is the sum of the five preceding equations which represent the reaction steps. Although the specific reaction rate of reaction step 1 is very much smaller than that of reaction step 5 , the steady-state concentration of 1 8 - is under ordinary conditions s'o small that the absolute rates of reactions 1 and j are equal. Any one of a number of similar processes would serve equally well to illustrate this type of catalysis. Typical examples are the decomposition of hydrogen peroxide catalyzed by the iodine-iodide ( I ) , the bromine-bromide (b), or the chlorine-chloride (22) couples, and the oxidation of chromic to dichromate ion or of ammonium ion to nitrogen (24) by peroxysulfate ion catalyzed by the trivalent-monovalent silver couple. Acid-base catalysis has not been mentioned in the foregoing discussion, only because it mas felt that a discussion of this phenomenon would not contribute anything useful to the study of catalysis in analytical chemistry. Of course (generalized) acid-base catalysis as well as kinetic salt and solvent effects (7) play important roles in determining the rate of oxidation-reduction reactions. Efficient catalysts for a given reaction can be selected (mithout direct trial) whenever the rates of the reaction steps are known. For example, in the case just cited it is well known that the rate of oxidation of thiosulfate ion by iodine is extremely rapid and that hydrogen peroxide oxidizes iodide ion more rapidly than it oxidizes thiosulfate ion. On the basis of this evidence, we could predict with certainty that the iodineiodide couple would act as an efficient catalyst for the 20 9
half-cell a t pH = 1 is not the e. m. f. which would be obtained in a half-cell prepared with equal amounts of cerous and ceric salts dissolved in 0.1 N acid, but is the potential which the half-cell xould have if the activities of the ceric and ow rous ions were equal.
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210
When the half-cell potential of the catalytic pair lies between the half-cell potentials of the compounds taking part in the main (stoichiometric) reaction, the free energy of each of the two compensating catalytic reaction steps-steps 2 and 5 in the first case considered-will be negative, and the pair in question is a possible catalyst. All oxidation-reduction couples which do not satisfy this condition cannot act as catalysts for the given reactions. For the ceric arsenite reaction a number of catalytic couples satisfy this condition. Of these, only compounds of iodine (17) and bromine are found to catalyze the reaction measurably, and the effect of bromine is too small to be of any practical value. The diagram indicates that the iodine-iodide pair will probably not be effective, but that either the hypoiodous-iodide or the iodate-iodine couple should be capable of acting as an efficient catalyst. It is very probable that the hypoiodous-iodide pair is the active catalyst, since it is well established that hypoiodous acid is intermediate in the oxidation of arsenite by triiodide ion. While the diagram does include as possibilities a number of compounds which are not efficient catalysts, it definitely excludes such, otherwise probable, couples as the manganic-manganous and cobaltic-cobaltous. The method is of greater practical value when the free energy of the main reaction has a smaller (negative) value-i. e., when the potentials of the corresponding half-cells are not so widely separated.
Induced Reactions In some cases the rate of an oxidation-reduction reaction is greatly increased when one of the reactants is simultaneously oxidizing or reducing some other substance. Such reactions are called induced or coupled reactions. The simultaneous oxidation of manganous ion and arsenious acid by chromate ion,
+ H8AS08 + Mn++ + 6H+ = HaAsO4 + Mn++++ Cr++++ 3H20
CrOl*
’
is an example of a simple induced reaction. In the absence of arsenious acid the oxidation of manganous ion by chromate ion is extremely slow. But in the presence of arsenious oxide, and under favorable conditions, one mole (1 equivalent) of manganous ion is oxidized for each mole (2 equivalents) of arsenious acid oxidized. This corresponds to a maximum induction factor-i. e., the ratio of the equivalents of the induced oxidation to the equivalents of the primary oxidationof one-half, A possible, although by no means established mechanism of this process may be represented by the following equations:
+ + +
+
(7,relatively rapid) Crop“ HaAs08-+HaAsO4 CrOs2CrOsHaAsOs 10H+ -c -+. HsAs04 2Cr+++ 5H20 (8, slow, probably complex) CrOsMn++ 6 H + + + Mn+++ Cr+++ 3Hz0 (9, rapid, probably complex)
+ +
+
+
+
+
The sum of Equation 8 and twice Equation 7 gives the stoichiometric equation for the primary reaction. 2Cr04‘
+ 3HsAs08 + 10H+
-
3HsAs04
+ 2Cr+++ + 5H20 (10, stoichiometric)
When the reaction step represented by Equation 9 occurs very much more rapidly than that represented by Equation 8, an induction factor of 0.50 is realized and the total stoichiometric reaction is represented by Equation 11.
- ++
++
++
GO4 H&Oa Mn++ 6H+ = Hs&04 M n + + + O f + + 3Hz0
(11, stoichiometric)
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Tn many of the practically important examples of this type
of reaction, the oxidation of a reducing agent by atmospheric oxygen is greatly accelerated by its simultaneous oxidation by some oxidizing reagents. The oxidation of stannous chloride by dichromate in the presence of air is a well-known example of this type of induced reaction. The value of the induction factor increases as the ratio of the concentration of dissolved oxygen to the concentration of the primary oxidant is increased. In practice this ratio depends on such factors as order of mixing, rate of stirring, etc. I n some cases the induction factor is also a function of the concentration of hydrogen ion and even of the concentration of the reducing agent. If a proper choice of conditions results in a value of the induction factor much greater than unity, the induced reaction is said to be an example of “induced catalysis” (6). If on the other hand the induction factor cannot be made to exceed unity (or some other small number), the reaction is considered to be an example of “simple coupling.” For all practical purposes, reactions of the ‘(induced catalysis” type are of greater importance. A possible mechanism for reactions of this type is catalysis (of the reduction of oxygen) due to a catalytic couple, one member of which is an intermediate in the primary reaction. When the oxidation by oxygen is a chain reaction, “induced catalysis” may be the result of the starting of reaction chains by an intermediate of the primary reaction. This latter mechanism has been clearly demonstrated for such reactions as the atmospheric oxidation of sulfite, induced by its primary oxidation with hydrogen peroxide. For this special class, the induction factor will be greatly decreased by the presence of a suitable inhibitor (3). I n some cases even microquantities of an inhibitor produce a marked decrease in the reaction rate. This suggests that the inhibiting effect of a substance could be used to determine its concentration michrochemically. Unfortunately, the inhibiting effect is seldom very specific, which greatly reduces the general utility of the method. However, it might be used to determine the concentration of an inhibitor in a solution which was known to be free of all substances which could act either as an inhibitor or as an inducing agent for the reaction.
Induced and Catalyzed Precipitations The theory underlying induced and catalyzed precipitations is different from induced and catalyzed reactions in homogeneous systems. Two different cases may be distinguished: 1. A constituent forms a slightly soluble crystalline compound upon addition of a suitable reagent. If the constituent is present in extremely small quantities (microconstituent) no precipitate is noticed, either because the amount is too small to be visible or because of super- or undersaturation. If in such a case a larger amount of another constituent (macroconstituent) is added, forming a slightly soluble compound with the same reagent, the microconstituent may be co-precipitated. Especially if the micro- and macrocomponents form mixed crystals, the enrichment of the microcomponent in the precipitate may be very pronounced. Thus a very dilute solution of lead does not yield a precipitate upon addition of sulfuric acid. If barium is added to the lead solution, mixed crystals of barium and lead sulfate are precipitated, the latter containing all the lead originally present in the solution. Direct application of such “induced” precipitations t o microdetection can be made when the microcomponent gives a colored precipitate with the reagent whereas the macrocomponent which is added yields a colorless precipitate. For example, nickel mercuric thiocyanate, which has a greenyellowish color, is precipitated by an alkali mercuric thiocyanate solution from relatively concentrated solutions
JULY 15, 1933
ANALYTICAL EDITIOIL'
only. However, when zinc is added to the solution the white crystals of zinc mercuric thiocyanate are colored green if nickel is present. Korenman (18)makes use of the same principle for the detection of cobalt. In the presence of small amounts of cobalt the zinc mercuric thiocyanate has a blue color. Montequi (25) found that copper mercuric thiocyanate, which normally has a green color, yields a violet precipitate when precipitated together with zinc. The reaction is applied as a sensitive method for the detection of copper. 2. HETEROGENEOUS CATALYSIS.It is beyond the scope of this paper to discuss the theory of heterogeneous catalysis. This has many analytical applications but the exact mechanism involved in most of these reactions has not been investigated yet. Upon treatment of various metal solutions with suitable reducing agents no metal is separated, although the reducing action of the agent is great enough to cause reduction t o the metallic state. An alkaline stannite solution reduces the bismuth salts to metallic bismuth (27). Use of this reaction is made in the sensitive detection of bismuth. Lead salts are slowly reduced by alkaline stannite, this reduction being fkrongly accelerated by traces of bismuth. By adding some lead salt to the solution the reaction for bismuth with alkaline stannite solution can be made 250 times more sensitive ( I S ) , Various examples of induced reductions involving the formation of precipitates are given by Feigl (8). The catalytic effect of mercury in the reduction of arsenic by stannous chloride or hypophosphite reagent is mentioned in She practical part of this paper.
Analytical Applications A substance is transformed into a reaction product which is readily detected. If the speed of transformation is slow, a suitable catalyst is added. Thus manganese is oxidized to permanganate by potassium persulfate in the presence of silver as a catalyst. More general use is made of the catalytic and inducing properties of substances, for the qualitative detection and quantitative estimation of the catalyst. Care should be observed in the interpretation of the results. Neither a positive nor a negative result as a rule is entirely conclusive regarding the presence or absence of a certain constituent (catalyst), for the following reasons: 1. 'The specificity of the catalytic or inducing action of the substame to be determined or detected must be considered. From the considerations presented in the theoretical part of this paper it may be inferred that in oxidation-reduction reactions various oxidizing or reducing substances, whose halfcell potential is intermediate between those of the two slowly reacting systems, may exert a catalytic effect. In many cases these changes in free energy are not known exactly. Therefore, after it has been found that a particular substance exerts a catalytic effect, a great number of substances (which might possibly serve as catalysts) should be tested also, in order to find how specific the catalytic effect is. Even if it has been found in an empirical way that the catalytic effect is quite specific for a certain reaction, it is not permissible to conclude the presence of this particular catalyst if a positive catalytic effect has been observed. The necessity for cautious interpretation may be inferred from the following example: It is mentioned by Feigl that traces of silver ion catalyze the reduction of tri- and tetravalent manganese in a hydrochloric acid solution. In testing the specific character of this test for silver the authors found that palladium exerts a similar effect, the reaction being even more sensitive for palladium than for silver. 2. The generality of the catalytic or inducing action must be considered. After it has been found that a substance exerts a catalytic effect, the question arises as to whether
211
general use of this effect can be made for the detection and determination of the catalyst. It should not be overlooked that other substances present may make the catalyst ineffective. For example, silver and manganese can be detected by making use of the catalytic effect of silver upon the oxidation of manganese to permanganate by potassium persulfate in strongly acid medium. Halide ions interfere, because they either precipitate the silver or reduce the permanganate formed. In a study of the catalytic effect of silver the authors found that in the presence of an excess of solid magnesium oxide manganese is oxidized to permanganate on boiling with persulfate in the presence of silver as a catalyst. This effect is quite specific for silver, and traces of this element can be detected in this way (see below). However, in testing the limitations of this test for silver it was found that cobalt and chloride make the silver ineffective. In certain cases where two or more substances exert a similar catalytic effect it should be possible to make advantageous use of the masking effect, in order to make the reaction more specific. For example, it was found that traces of iodine (or iodide, or iodate) and of osmium exert a tremendous catalytic effect upon the reaction between ceric sulfate and arsenic trioxide in acid medium. I n this case the iodide can be made ineffective by the addition of a small amount of mercuric mercury. In the quantitative estimation of a catalyst use is made of the relation between the concentration of the catalyst and the rate of the catalyzed reaction. It should be realized that quite generally indifferent electrolytes affect the rate of a reaction. Therefore, this relation between concentration of catalyst and reaction rate should be known for the particular mixture in which the catalyst is to be determined. The difficulty as a rule can be eliminated in a practical way. The rate of reaction is determined first by mixing a measured volume of the solution in which the catalyst is t o be determined with a measured volume of the reacting components. After the approximate amount of the catalyst is determined, about an equal but exactly known amount is added to the same volume of the solution originally taken, without changing the volume appreciably (additions with microburet or micropipet) and the measurement is repeated. If the rate of reaction is proportional to the concentration of the catalyst, the amount of the catalyst is readily calculated from the two sets of measurements. Various modifications of this procedure are possible. The simplest procedures for measuring the rate of reaction are found in those cases in which one of the reacting components is colored. The time required for obtaining a colorless solution is measured. A reversal of this procedure is possible if a color develops after one of the reacting components has been quantitatively oxidized or reduced. I n other cases the amount of reaction product formed is measured quantitatively. The catalytic effect of copper in the oxidation of cysteine to cystine and the subsequent measurement of the oxygen evolved in a given time was used by Warburg (28) for the determination of microquantities of copper in blood and was applied to milk by Zondek and Bondmann (29). Great care, however, should be observed, since relatively large and variable readings are obtained in blanks ( 2 ) .
Practical Applications Since the catalytic or inducing effect is often exerted by extremely small traces of a substance, it is possible t o make use of the catalysts for the microdetection and estimation of the particular substance. It is beyond the scope of this paper to give an exhaustive treatment of all catalytic and induced reactions which find application in qualitative and quantitative microanalysis. As examples, the catalytic and inducing effects of silver and mercuric mercury are discussed below and some new reactions are described.
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INDUSTRIAL AND ENGINEERING CHEMISTRY
Catalysis by Silver 1. The catalytic effect of silver ions in oxidation reactions is t o be attributed to the intermediate formation of a higher oxidation state, probably of trivalent silver ions. The latter have a high oxidation potential and are not stable as such in aqueous medium. The catalytic effect of silver ions in the oxidation of manganese to permanganate, of trivalent chromium to chromate, and of trivalent cerium to ceric cerium in strongly acid medium by persulfate has been known for a long time. Use of it was made by Marshall (23) in the detection of traces of manganese. T o the authors' knowledge the reaction has never been used for the detection of silver, although it was found t o be highly sensitive. PROCEDURE. One milliliter of a manganous sulfate solution containing 0.1 mg. of manganese is diluted to 5 ml. with 4 N sulfuric acid and 1 ml. of the solution to be tested for silver is added. About 0.2 to 0.4 gram of potassium persulfate is then added and the tube is placed in a boiling water bath. In the absence of silver the solution remains colorless, even after 5 minutes of heating. In the presence of 0.05 y (= 0.00005 mg.) of silver a distinct violet color appeared after 2 minutes of heating. The test is therefore sensitive to one part of silver in 20,000,000 parts of solution when 1 ml. of the unknown is tested in the above way. The reaction is specific for silver. Osmium tetroxide, auric gold, palladium, cobalt, copper, mercury, and various other cations did not catalyze the reaction. Halide ions interfere. Even traces of chloride in the solution prevent the formation of permanganate or make the reaction for silver much less sensitive.
It was found in this work that silver not only catalyzes oxidations with persulfate in strongly acid medium, but also in weakly acid and in alkaline medium. Thus manganese is oxidized to permanganate in sodium hydroxide, or sodium carbonate medium or in the presence of excess magnesium oxide, Silver also catalyzes the oxidation of lead t o lead peroxide by persulfate in neutral or alkaline medium. Five milliliters of 0.1 per cent lead nitrate solution were treated with 5 ml. of 1 N sodium acetate and divided into two parts. One drop of 0.1 per cent silver nitrate was added to one part and 0.5 gram of potassium persulfate to both. The tubes were placed in a boiling water bath. A reddish brown colloidal precipitate of lead peroxide was formed in the presence of silver, whereas the blank remained clear, except for a slight white precipitate on cooling, This test for lead is not very sensitive and of little analytical significance, since many ions interfere with the silver catalysis. On the other hand the oxidation of manganese to permanganate in alkaline medium in the presence of silver as a catalyst is of analytical significance. DETECTION OF MANGANESE IN WEAKLY ALKALINEMEDIUM. Fifty milligrams of magnesium oxide (manganese-free), 1 drop of 0.1 per cent silver nitrate, and 0.2 to 0.3 gram of potassium persulfate are added to 10 ml. of the solution to be tested for manganese. The tube is placed in a boiling water bath for 3 to 4 minutes, then removed and allowed to settle. A red-violet coloration in the supernatant liquid shows the presence of manganese, Sensitivity: 0.3 mg. manganese per liter gave a distinct pink color. It should be mentioned that the magnesium oxide should be tested in the same way for manganese. It was found that c. P. products of magnesium oxide contained traces of manganese. With large amounts of manganese (100 mg. per liter) a heavy precipitate of manganese dioxide is formed which masks the silver catalysis. The solution to be tested should not contain more than 10 mg. of manganese per liter. The reaction is specific for manganese, although cobalt and nickel interfere, giving black precipitates. The presence of 0.05 mg. of nickel did not affect the sens tivity of the reaction, but the same amount of cobalt decreased the sensitivity markedly. Chlorides interfere. By working in strongly alkaline medium the interference by chlorides is eliminated. DETECTION OF MANGANESE IN STRONGLY ALKALINE MEDIUM w PRESENCE OF CHLORIDE.To 10 ml. of solution are added 1 drop of 0.1 per cent silver nitrate, 0.5 to 1 ml. of 4 N sodium hyaroxide, and 0.2 to 0.4 gram of potassium persulfate. The
VOL. 7, NO. 4
mixture is heated to boiling for 30 seconds to 1 minute and allowed to settle. A pink coloration of the supernatant liquid indicates the presence of manganese. The test is sensitive to 5 y of manganese, corresponding to 0.5 mg. of manganese per liter. Cobalt interferes, but small amounts of nickel and copper may be present. DETECTIONOF SILVER IN WEAKLYALKALINE MEDIUM. To 10 ml. of the solution to be tested are added 1 ml. of a manganese chloride solution containing 100 mg. of manganese per liter, 50 mg. of magnesium oxide, and 0.2 to 0.3 gram of potassium persulfate. The tube is heated in a boiling water bath for 3 to 4 minutes, removed, and allowed to settle. A red-violet color indicates the presence of silver. The test is sensitive to 3 y of silver in 10 ml., corresponding- to a concentration of 0.3 mg. - of silver per liter. The reaction is specific for silver. Nickel, cobalt, and chloride interfere. In strongly alkaline medium chloride does not interfere, but the reactionlbkcomes slightly less sensitive. DETECTION OF SILVERIN STRONGLY ALKALINEMEDIUMIN PRESENCE OF CHLORIDE.To 10 ml. of the solution are added 0.2 ml. of manganous sulfate solution containing 0.1 gram of manganese per liter, 1 ml. of 4 N sodium hydroxide, and 0.2 to 0.4 gram of potassium persulfate. The mixture is boiled for 30 seconds to 1 minute and allowed to cool. A blank without silver is used for comparison. A pink coloration shows the presence of silver. The test is sensitive to 5 y of silver, corresponding to 0.5 mg. of silver per liter. Copper catalyzes the oxidation, but the reaction is not sensitive for copper. Palladium inhibits the silver catalysis; gold has a similar effect but not as pronounbed as palladium. Cobalt interferes. but small amounts of hickel may be iresent. 2. Silver chloride is more or less soluble in stronger chloride solutions with the formation of a complex AgC12- ion. It was found by Lang (20) that brown solutions of manganic manganese (MnI'I and Mn") are fairly stable in 2.5 N hydrochloric acid. Addition of a trace of silver catalyzes the reaction between the higher-valent manganese and chloride and the solution becomes colorless within a short time. Feigl and Frankel ( I d ) found that silver also catalyzes the reaction between ceric cerium and hydrochloric acid. I n using a, suitable manganic chloride solution in 2.5 N hydrochloric acid as a reagent they could detect on a spot plate 0.4 y of silver _(1drop of a solution containing 8 mg. of silver per liter). The authors found a sensitivity of 1 y silver, but in testing the specific character of the reaction noticed that palladium has a stronger catalytic effect than silver. In a spot plate 0.2 y of palladium can be detected. With a suitable mixture of ceric nitrate and hydrochloric acid Feigl and FrSinkel could detect 0.05 y of silver. However, the authors found that the sensitivity was less than 1 y. I n addition it was found t h a t palladium behaves like silver; the test, therefore, is not specific. Thallous thallium interfered badly with the test, a yellow precipitate and a rapid decoloration being obtained. 3. Silver salts exert an inducing effect upon the reduction of mercuric chloride by phenylhydrazine (10). I n the absence of silver the mercuric chloride is reduced t o calomel and then slowly to metallic mercury. In the presence of silver the black mercury separates instantaneously. Hahn (14) found that traces of silver accelerate reduction of mercuric chloride to calomel by hypophosphite in a well-buffered solution and makes use of this fact in a sensitive detection of silver.
Catalysis by Mercury 1. King and Brown (16) found that mercuric mercury strongly catalyzes the reduction of arsenite or arsenate in hydrochloric acid medium to brown arsenic by stannous chloride. The addition of enough mercuric chloride to make its concentration 0.00001 M before the addition of stannous chloride, hastens the appearance of the coloration, increases the sensitivity of Bettendorf's test 10- t o 100-fold and enables the test to be made in a lower concentration of hydrochloric acid. The authors were able t o confirm the catalytic effect of mercury upon the speed of reduction of the arsenic, but
JULY 1.5, 1935
ANALYTICAL EDITION
did not find the extreme sensitivity mentioned by King and Brown. Thus, 0.2 ml. of N arsenic trioxide was mixed with 9 ml. of concentrated hydrochloric acid and 1 ml. of 1 M stannous chloride (in 6 N hydrochloric acid). Without addition of mercuric chloride, a slight coloration was visible after 15 minutes of M mercuric chloride was added before standing. If 1 ml. of the stannous chloride, a faint color was perceptible after 2 minutes of standing. With 0.03 ml. of 10-8 N arsenite acid, stannous chloride] and mercuric chloride a doubtful coloration was perceptible after 15 minutes of standing (limit of sensitivity); without mercuric chloride no change was noticeable after the same time. King and Brown mention that use of the catalytic effect ,of the mercury can be made in its quantitative approximation. This was confirmed by the authors of this paper. As an illustration the following experiments are given : One-half milliliter N arsenic trioxide] and x ml. of iM mercuric chloride were diluted to 9 ml. with concentrated hydro-. chloric acid, and 1 ml. M stannous chloride was added. As compared with a solution containing no mercuric chloride (x = 0), a distinct difference was noted after 2 minutes with 0.2 ml., 3 minutes with 0.1 ml., 5 minutes with 0.03 ml., 7 to 8 minutes with 0.1 ml. Silver and copper do not interfere. Palladium interferes badly with the arsenic and mercury test, since it gives a dark turbidity upon addition of stannous chloride (palladium metal). Gold also interferes. The authors found that mercuric mercury also catalyzes t h e reduction of arsenic by hypophosphite. Reagent 1, 10 per cent calcium hypophosphite in 25 per cent . . . hydroclhloric acid. Reagent 2, reagent 1freshly saturated with mercurous chloride. PROC’EDURE. Three milliliters of the solution to be tested are mixed with 2 ml. of concentrated hydrochloric acid and 3 ml. of reagent, and placed in a boiling water bath. TA13LE
I. REDUCTION O F ARSENICBY HYPOPHOSPHITE
Arsenic Present in 3 M1. Y
300 100 30 10 5
Time Required to Give Brown Turbidity Reagent 1 Reagent 2 Win. Mzn.
1.5 2 4 30 (doubtful)
..
