tate-acetic acid buffer and 5 ml. of the color forming reagent to both flasks. Dilute both flasks to the graduation mark. If more than one sample is to be analyzed, it is advisable to Fork with one sample a t a time. The sample(s) should stand for 15 minutes to allow the color to develop. The solution containing hydrogen peroxide is used in the reference cell and the other solution is used in the sample cell to determine the quantity AI - Aa, as given in Equation 5. The absorbance is measured at 510 mP. If higher concentrations of hydrogen peroxide are desired, the procedure can be modified slightly. After adding iron(I1) from a microburet to each of two 50-ml. volumetric flasks, add different volumes of the hydrogen peroxide solution-e.g., 1 and 1.5 m1.-to the flasks. Add 10 ml. of the buffer solution and the 1,lO-phenanthroline reagent and dilute to volume. The solution containing the least hydrogen peroxide is used in the sample cell. Checking Precision of 1,lO-Phenanthroline Method. Exactly 0.50 ml. of 0.194111 hydrogen peroxide solution was diluted t o 1 liter. This 9.70 X l O - 4 M solution was used in the subsequent investigation. The color was developed in the following manner. Add 2 ml. of the iron(I1) solution, 2 drops of concentrated sulfuric acid, 10 ml. of the hydrogen peroxide qolution, 10 ml. of the acetate buffer solution, and 10 ml. of the 1 , l O phenanthroline reagent to a 50-ml. volumetric flask, and dilute to volume. Allow this solution to stand 15 minutes for complete color development. This order of addition is recommended, and if more than one sample is to be analyzed, it is advisable to deveIop each color separately. The results obtained are shown in Table 11.
Effect of Solution Variables Using Bathophenanthroline. The solution variables are essentially the same as in the 1,lO-phenanthroline method. The iron(I1) concentration may not be in excess by more than 0.015 mg. per ml. when bathophenanthroline is used. Ethyl alcohol is used as a solvent because of the limited solubility of bathophenanthroline in aqueous solution. The molar absorptivity of the complex has been reported as 22,400 liters per mole-em. a t 533 mp ( I S ) . Diverse Ion Concentration. It was found that 100 p.p.m. of the following ions did not interfere: chloride, nitrate, acetate, sulfate, perchlorate, arsenate, nickel(II), zinc, cadmium, copper(II), mercury(II), titanium(IV), lead, bismuth, and tin(1V). Silver ions in excess of 30 p.p.m. causes a turbidity. Cobalt(I1) interferes because of formation of a yellow complex. PROCEDURE USING BATHOPHENANTHROLINE
This procedure is identical to the one using 1,lO-phenanthroline except that 95% ethyl alcohol is used for dilutions. The absorbance is measured a t 533 mM. Checking Precision of Method Using Bathophenanthroline. The color wasdeveloped in the follon-ing manner: Add 1 nil. of iron(I1) solution, 2 drops of concentrated sulfuric acid, 1 ml. of hydrogen peroxide solution, 20 ml. of the acetate buffer solution, and 5 ml. of the bathophenanthroline solution to a 50-ml. volumetric flask. Dilute t o volume with 95’% ethyl alcohol. Allow the solution to stand 20 minutes for
complete development of color. The results are shown in Table 111. SUMMARY
The proposed procedures for determining hydrogen peroxide are more sensitive than the peroxytitanic acid method because of the high molar absorptivities of the iron(I1) phenanthroline complexes. As indicated in Tables I1 and 111, to 10-6M solutions of hydrogen peroxide can be analyzed with satisfactory precision and accuracy. LITERATURE CITED
(1) Allen, Nelson, IND.ESG. CHEJI., ANAL.ED. 2, 55 (1930). (2) Bonet-Maury, P., Compt. rend. 218, 117 (1944). 1 3 ) Eisenberc. G. M.. ISD.ENG.CHEJI.. ANAL.ED.%, 327 (1943). (4) Erdmann, H., Seelich, F., Z. a n d Chem. 128, 303 (1948). (5) Hartmann, S., Glavind, S., d c t n Chem. Scand. 3, 954 (1949). (6) Isaacs, M. L. , J. Am. Chenz. Soc. 44, 1662 (1922). \
z
( 7 ) IlcCabe, S., Ind. Eng. Chem. 45, 111.1 (September 1953). (8) Miller, MacPherson, N.,Natl. Research Council Can. At. Energy Project Div. Rpsearch C R d 352 (N.R.C. KO. 1617). (9) Ownston. T. C. J., Rees, TI-. T., - --Analyst 75,’204 (1950). (10) Sandell, E. B., “Colorimetric Determination of Traces qf Metals,” 2nd ed., na.. 376-7, Interscience. Yew Tork. _ 1950. (11) Savage, J., Analyst 76, 224 (1951). (12) Schales, Otto, Ber. 71, 447 (1938). (13) Smith, G. F., Richter, F. P., “Phe\
,
nanthrolineand Substituted Phenanthroline Indicators,” G. Frederick Smith Chemical Co., Columbus, Ohio, 1944. RECEIVED for review May 1, 1958. Accepted August 4, 1958.
Polarographic Determination of Chlorate LOUIS MEITES and HENRY HOFSASS Polytechnic Institute o f Brooklyn, Brooklyn, N. Y. ,An investigation of the polarographic characteristics of chlorate in various mineral acid media has led to the development of a polarographic method for its determination, based on the rapid stoichiometric reaction between chlorate and excess ferrous iron in 3F hydrochloric acid. It is accurate Chloand precise to about & l % . rate solutions as dilute as 10-W can be analyzed.
D
the very positive value of the standard potential of the halfreaction ESPITE
C103-
+ 6 H + + 6e = + 3H20; E” = + 1.47 volt ( 4 )
C1-
chlorate ion is not reduced a t the dropping mercury electrode from alkaline, neutral, or moderately acidic solutions ( 2 , 12). This evidently reflects the existence of an extremely slow step in the reduction process. Consequently the polarographic determination of chlorate could only be accomplished by increasing the rate of the electroreduction (either by adding a catalyst or by using a strongly acidic supporting electrolyte) ,or by converting the chlorate into an equivalent amount of some other
substance which n-ould give rise t o a polarographic wave. The first of these approaches was employed by Koryta and Tanygl (S), who found that the catalytic current obtained in an acidic oxalate medium containing 10m,ll titanium(1V) was proportional to the chlorate concentration from 1 t o 15mX. This concentration of titanium(1T’) is so high that an accurate correction for its diffusion current would become extremely difficult in the analysis of chlorate solutions much more dilute than 1mM. This paper reports the results of investigations of both of these possible VOL. 31, NO. 1 , JANUARY 1959
119
types of methods, and describes an indirect procedure which should be useful for the determination of chlorate in many kinds of samples. EXPERIMENTAL
Polarograms were recorded with a calibrated pen-and-ink recording polarograph, especially constructed ( 7 ) . A conventional dropping mercury electrode assembly was used. The modified H-cell ( I O ) was provided with a silver-silver chloride-saturated potassium chloride electrode. Solutions were deaerated with prepurified nitrogen which was scrubbed with chromous chloride to remove traces of oxygen. All measurements were made a t 25' i 0.005" C. All chemicals were ordinary reagent grade and were used without purification. RESULTS A N D DISCUSSION
Osmium tetroxide is an efficient catalyst for certain processes involving the reduction of chlorate ( I ) . In 0.2 to 1F solutions of hydrochloric, perchloric, and sulfuric acids containing 0.001mM osmium(VIII), no catalysis of chlorate reduction could be detected over the range of potentials in which osmium is reduced to the +3 state (8). However, in the perchloric and sulfuric acid media a definite increase of current was observed just prior to the final current rise, indicating that the osmium(11)formed a t this more negative potential (8) does catalyze the chlorate reduction to some extent. Unfortunately, the catalytic chlorate wave does not begin until the potential becomes so negative that there is no room for the development of a plateau, and the phenomenon is analytically useless. Similar experiments using molybdenum(V1) in place of osmium(VII1) gave no catalytic current. Although the electroreduction of chlorate from these dilute acid media is very slow, a very well defined wave rising from zero applied potential was observed with fresh solutions of chlorate in more concentrated hydrochloric acid. The height of this wave decreased about 60% per minute of deaeration with a moderately rapid stream of nitrogen, and eventually fell to zero. Evidently the wave is due to a volatile product of a reaction between chlorate and the hydrochloric acid. According to the literature (6, I S ) , this might be either chlorine or chlorine dioxide; evidence favoring the first of these is presented below. Plotting the logarithm of the observed wave height against total deaeration time and extrapolating to zero time gave the initial wave heights which are shown as a function of the hydrochloric acid concentration in Figure 1. In 6 to 12F acid the initial wave height, 120
ANALYTICAL CHEMISTRY
by deaeration in the same way and a t the same rate as the wave observed in hydrochloric acid solutions. These v a i es are undoubtedly due to chlorate and chlorine dioxide, respectively, and the large difference between the halfviave potentials of the chlorine dioxide wave and the wave obtained in hydrochloric acid media must be due to the fact that the reducible substance responsible for the latter wave is actually chlorine. The polarographic determination of chlorate in sulfuric and perchloric acid media is clearly inadvisable because of 0 0 4 0 12 the effects of traces of chloride which might be present. HCI, F An analytically useful wave whose Figure 1. Initial diffusion currents for height is proportional to chlorate con0.9mM chlorate in various concentracentration is obtained in hydrochloric tions of hydrochloric acid acid solutions containing excess stannous or ferrous ion; these are oxidized Ed.8. = -0.3 Volt VS. S.C.E., m'/a = 2.5 by the chlorate t o yield equivalent amounts of stannic or ferric ion, respectively, and the wave heights are 21 i 3 pa., is in good agreement ivith unaffected by deaeration. Ferrous ion the figure calculated from the average is better for this purpose than stannous value of 1.5 for I / n (9) together with ion, because the ivave of the excess the experimental values of mz'3t1'6 stannous ion follows so closely upon the and C and the value of n for the over-all wave of the stannic ion produced (5, I f ) process that the height of the latter is difficult to measure accurately. C1036H+ 6e = C13H20 The formation of chlorine by the direct This indicates that the chemical reaction between chlorate and hydroreduction of chlorate by hydrochloric chloric acid is so rapid a t acid concenacid is rapid and quantitative when the trations above 6F that some loss of concentration of the latter is above 6F. chlorine during deaeration might be At lower acid concentrations, however, feared even from solutions containing this reaction becomes slow, as is imexcess ferrous ion. However, the rate plied by Figure 1 and shown explicitly of the reaction between ferrous ion and by Figure 2. chlorate is also dependent on hydroIn 3F hydrochloric acid, the reaction chloric acid concentration, and is imis only about 40% complete after 1 practicably slow when the acid concenhour. The extreme slonmess of this tration is as low as 1F. From these reaction is responsible for the observaconsiderations it appears that an acid tion that chlorate gives no wave in concentration close to 3F is most suitdilute hydrochloric acid solutions. able. The reaction between chlorate and 3F The effective diffusion current conhydrochloric acid is first-order with stant of chlorate in 38' hydrochloric acid respect to chlorate during nearly the containing 50mX ferrous ion is shown entire period covered by Figure 2, and in Table I. It is termed effective the first-order rate constant in the because it is calculated from the concenequation tration of chlorate which would have been present had no reduction by ferrous ion occurred. The error a t very low chlorate concentrations is largely is 0.00260 min.-l a t 25.00' C. due to the uncertainty in the residual Polarograms of chlorate in 5F sulcurrent correction caused by the airfuric or 8F perchloric acid containoxidation of the ferrous ion. This ing only traces of chloride, which would makes it advisable to prepare the ferfavor the formation of chlorine dioxide rous solution immediately before use rather than chlorine, according to the by weighing the appropriate amount of equations solid ferrous ammonium sulfate into air-free 3F hydrochloric acid in the C1035C16H' = 3C12 3Hz0 polarographic cell.
+
+
+
+ + + 5C103- + C1- + 6H+ = 6C101 + 3H10
showed two waves. The first rose from zero applied potential; its height decreased slowly on standing, and mas virtually unaffected by deaeration. The second had a half-wave potential of -0.8 volt us. S.C.E., slowly increased in height on standing, and was affected
Measured volumes of standard chlorate solutions were added to known volumes of freshly prepared air-free 3F hydrochloric acid-0.05F ferrous ammonium sulfate, and the limiting currents were measured a t -0.3 volt us. S.C.E. after about 5 minutes' deaeration, Corrections have been applied
I
I '
I.0-
.-E
' 0.5.-U
I
0 0
2
Moles 0
100
50 Time,
150
minutes
Figure 2. Current-time curve obtained with 3F hydrochloric acid initially containing 2mM chlorate Ed.a. =
-0.3 volt
VS.
S.C.E.,
m'/r f'/@ = 2.5
I
I
I
6 F e ' + / m o l e ClO; 4
8
I . IO
Figure 3. Effect of ferrous-chlorate ratio on measured diffusion currents in 3F hydrochloric acid im.
Current obtained with large excess of ferrous ion Data obtained with solutions containing 8 to 50mM ferrous ion and 1 to chlorate
9mM
stoichiometric was obtained from measurements of the diffusion current constant of ferric ion in 3F hydrochloric acid containing excess ferrous ion. With 50mM ferrous ion, the effective diffusion current constant of chlorate as defined above is 8.77 f 0.09, while the diffusion current constant of ferric ion is 1.46 f 0.014. With lOmM ferrous ion, the corresponding figures are 8.93 =k 0.09 and 1.48 f 0.024. In either medium Time, minutes
ICQ-
Figure 4. Rate of ferrous-chlorate reaction in solutions initially containing 3F hydrochloric acid, 0.8mM chlorate
=
R IF^++'
where R is the number of ferric ions resulting from the reduction of one chlorate ion. The values of R in the two media are 6.01 and 6.00, respectively. This proves that the stoichiometry of for the residual currents, which were the chlorate-ferrous reaction is corTable 1. Effective Diffusion Current measured just before the addition of rectly described by the equation given Constant of Chlorate in 3~ Hydrothe chlorate. Several different capillarabove. ies were used, and the value of m2/3t1/b chloric Acid containing 5 o m ~F ~ ~ was The presence of a layer of chloroform always measured immediately after rous Ion above the pool of mercury which accuthe limiting current measurement. mz/3t118, mulates at the bottom of the cell is necesThree milliliters of chloroform were pres~1g.*/3 ent to protect the mercury pool which [ClOs-I, sary to prevent the oxidation of the mermM id, pa. Sec.-1/2 I accumulated during the measurements. cury. The solid mercurous chloride 0.023 2.36 (9.9) 0.000980 which would thus be formed is not At low ferrous-chlorate ratios, the 0.100 2.34 8.88 0.00481 reducible, and the omission of the 0,206 2.44 8. 70a ferrous ion is quantitatively oxidized 0.00971 chloroform leads to low and erratic 2.44 8 . 70 0.00980 0,208 and the excess chlorate is reduced to diffusion currents. 2.34 8 . 89. 1.01 0.0485 chlorine which escapes into the gas 2.35 8.86 0.0490 1.02 Chlorate concentrations above about stream, so that the measured diffusion 2.39 8 .83a 0.0971 2.05 I m X give rise to a large maximum on current is too low. Experiments in 2.38 8 .75 2.04 0.0980 the polarogram, which may be sup2.39 8.92O 0.1905 4.06 which the ferrous-chlorate ratio was pressed by 0.002% Triton X-100. 2.48 8 .85 0.1923 4.22 varied around the stoichiometric value 2.37 8. 795 8.09 0.3883 The data shown in Table I for solutions gave the results shown in Figure 3. 2.40 8. 85 0,3922 8.33 containing smaller concentrations of from which it is apparent that this 2.37 8 .810 15.9 0.762 chlorate indicate that the presence of 2.41 8 . 79 0.769 16.3 ratio must be a t least 10 t o ensure satisthis concentration of the Triton does not 2.41 8. 67 0.9615 20.1 factory results. This corresponds to 2.41 8 .7 la 1.905 40.0 change the effective diffusion current an excess of about 50% over the quan2.41 8 .6ga 3.670 76.6 constant of chlorate by as much as tity of ferrous ion required if the reac2.41 8 .58" 5.60 115 i 0.5%. Hence the routine use of the 2.42 8 .75" 6.80 144 tion taking place is maximum suppressor in chlorate deter2.42 8 .55" 7.30 151 9.01 154 2.42 (7.06)= CIOs-+ 6H+ + 6 F e f + = minations seems advisable. Figure 4 shows typical data on the Mean (excluding values in parentheses) : C13H20 6 F e + + + 8.77 =k 0.09 rate of the chlorate-ferrous reaction. ,, Contained 0.002% Triton X-100. Evidence that the chlorate-ferrous From this and similar plots the followreaction shown by this equation is ing equation can be deduced for the Upper. 1OmM ferrous ion Lower. 50mM ferrrour ion
+
+
VOL. 31, NO. 1, JANUARY 1959
121
rate of this reaction in 3F hydrochloric acid
where the quasi-second-order rate constant, k, is 20.6 i. 0.8 1. per mole per minute a t 25.00° C. This information is of value in choosing a ferrous ion concentration for use in the determination of chlorate; with less than about 0.01;11ferrous ion the reaction is too slow for convenience. On the other hand, high concentrations of ferrous ion lead to undesirably large corrections for the ferric ion present in the reagent as a n impurity. A ferrous concentration of 50mU, which permits the accurate determination of chlorate concentrations from 0.005 to 7 m X (Table I), seems to represent a reasonable compromise between these factors.
Perchlorate and chloride will not interfere in the determination of chlorate by this method. Such strong oxidizing agents as hypochlorite, chlorine, ferric, pervanadyl, chromate, and bromate could be removed by a controlled-potential electrolysis of a weakly acidic solution prior to the determination of chlorate. Hence this information should make possible the polarographic determination of chlorate in a wide variety of materials. LITERATURE CITED
(1) Hofmann, K. A., Bel. 45, 3329 (1912). (2) Kolthoff, I. M., Lingane, J. J., “Polarography,” 2nd ed., p. 576, Interscience, hew l o r k . 1952. (3) Koryta, J., Tanygl, J., Chem. listy 48, 467 (1954). (4) Latimer, W. M., “Oxidation States of
the Elements and Their Potentials in Aqueous Solutions,” p. 56, 2nd ed., Prentice-Hall, New York, 1952.
( 5 ) Lingane, J. J., J . Ani. Chem. SOC.67, 919 (1943).
( 6 ) Luther, R., MacDougall, F. H., 2. phusik. Chenz. 5 5 , 477 (1906). ( 7 j &Teites, L., J . . Am. Chem. SOC.76, ’
5927 ( 1954).
(8) Meites, L., Ibid., 79, 4631 (195i). (9) .Meites,
L., “Polarographic Techniques,” p. 73, Interscience, NeTv Tork, 1955.
(10) Meites, L., illeitea, T., A N A L . CHEM. 23,1194(1951). (11) Meites, L., BIeites, T., unpublished experiments cited in (9). (12) Mjlner, G. W. C., “Principles and
Applications of Polarography and Other Electroanalytical Processes,” p. 289, Longmans, Green, Nen- York, 1957. (13) Taube, H., Dodgen, H., J . Bin. Chem. SOC.71, 3330 (1949).
RECEIVEDfor review June 19, 1958. Accepted August 8, 1958. Based in part on thesis submitted by H. Hofsass to the faculty of Polytechnic Institute of Brooklyn in partial fulfillment of requirements for the B.S. in chemistry degree, June 1958.
.
A Simple Sensitive Test for AI I phatic Ketones EUGENE SAWlCKl and THOMAS
W. STANLEY
Community Air Pollution Program, Robert A. Tuft Sanitary Engineering Center, U. S. Public Health Service, Department of Health, Education, and Welfare, Cincinnati 26, Ohio
F A simple, sensitive color test for acetone and other aliphatic ketones containing the structure, RCH2COCHzR’ , is based on a condensation of the ketones with 2-hydroxy-1 -naphthaldehyde under special conditions. Limits of iden’ification are about 1 to 5 y, the colors are stable, and a blue color, A,, 598 mp, is usually obtained. The procedure consists of the addition of 0.5 ml. of a 2-methoxyethanol test solution to approximately 75 mg. of the test powder, which is made of 1 part of 2-hydroxy- 1 -naphthaIdehyde and 3 parts of anhydrous aluminum chloride. The color is read in 3 minutes. The exothermic reaction of the reagent powder with 2-methoxyethanol furnishes the heat necessary for the development of the color. This test is superior in simplicity of operation, color differential between blank and test solution, color stability, and sensitivity to previous procedures.
N
for a simple, specific, and reliable test for aliphatic ketones has been shown by the numerous publications on this subject which have appeared during the last century (2, S, 7 , 8, 14, 17, 19, 20). There are four main colorimetric methods for the detection of ketones containing the struc122
EED
ANALYTICAL CHEMISTRY
ture RCH,COCH,R’. Many of these color tests involve the condensation of the ketone with an aromatic aldehyde in alkaline (1) or acid (23) media to give a colored diarylidene acetone (17). Some of the other aromatic aldehydes which have been used are furfural ( I @ , p-hydroxybenzaldehyde (WS), salicylaldehyde ( I I ) , and vanillin (18, 23). The disadvantages of the furfural test are its 24-hour duration (92) and the reportedly negative results for methyl ethyl ketone (27). The salicylaldehyde test requires 20 minutes of heating and is relatively insensitive for diethyl ketone (26). I n all these tests the blank is yellow, vihile the positive color is orange to violet. This is not a very sharp distinction because compounds such as acetophenone also react to give orange to red chalcone derivatives. A second color test is based on the formation of the blue dye, indigo, in the reaction between o-nitrobenaaldehyde and an aliphatic ketone. Acetyl derivatives, like acetophenone, acetaldehyde, and biacetyl, give positive results. The test is not very sensitivea n identification limit of 100 y being reported for acetone (9). A third color test uses sodium nitroferricyanide as the reagent (3, 9, 10, 24, 26, 28). The basis of the color reaction is stated to be the reaction of
sodium nitroprusside with the acetone anion to produce the unstable highly colored anion, [Fe(CN),(O?r’= CHCOCH3)]-4 (4). A satisfactory modification of this test has been used in clinical chemistry to test for the presence of acetone and acetoacetic acid in urine, spinal fluid, or blood serum ( I , I d , 14, 21). Ten micrograms of acetone can be detected by the method. This test is not specific because acetophenone ( 9 ) , pyruvic acid (9), aliphatic aldehydes (S), pyrrole (g), and mercaptans give positive reactions. I n addition, aliphatic ketones, such as methyl ethyl ketone, give a fairly insensitive light orange color with the reagent. A fourth method for the detection of aliphatic ketones is based on the formation of a red-violet complex between m-dinitrobenzene and the aliphatic ketone anion in alkaline solution (5, 15). This reaction has been successful in the detection and determination of ketonic steroids. It forms the basis of all reliable methods for the determination of 17-oxo steroids (16, 29). The main disadvantages are that it is not too sensitive for aliphatic ketones, the color is not stable, and compounds such as acetophenone and acetaldehyde give positive results ( 3 ) . The new color test for aliphatic ketones probably involves the reactions
