Polarographic Behavior of Monohaloacetones - Analytical Chemistry

---

1908

ANALYTICAL CHEMISTRY

This method of analysis is being used in this laboratory for the determination of acrylonitrile in the presence of a known concentration of potassium persulfate, along with the various oxidation products of acrylonitrile mentioned earlier. The precision of the method compares favorably with the usual precision of polarographic analyses. For 31 sets of duplicate observations of diffusion-current values in the range of 80 to 130 (expressed as millimeters wave height), the variance of measurement was 0.2337. The method should be applicable to polymerization systems containing acrylonitrile and potassium persulfate initiator, It should also be suitable for solutions containing methacrylonitrile instead of acrylonitrile, since the half-wave potentials of these two compounds differ only by about 50 mv. (11). As sodium, rubidium, and cesium ions exhibit half-rrave potentials very close to potassium ion ( 7 ) it should be possible to determine acrylonitrile or methacrylonitrile in the presence of any one of these ions when the concentration of the alkali metal ion is known. ACKNOWLEDGMEYT

The authors wish to thank E. Vernon Lewis, of the Department of Mathematics, for his very valuable advice and assistance with the statistical calculations. The support of this investiga-

tion by a grant from the National Science Foundation is gratefully acknowledged. LITERATURE CITED

m'.

(1) Beesing, D. W., Tyler, P., Kurts, D. lf., and Harrison, S. A., AN-~L.CHEM.,21, 1073 (1949). (2) Bird, IT.L., and Hale, C. H., Ibid., 24, 586 (1952). (3) Dal Nogare, Stephen, Perkins, I,. I t , , and Hale. -1.H.. Ihid., 24, 512 (1952). (4) Gaudry. Roger, Org. Syntheses, 27, 41 (1947). (5) Goulden, C . H., "Methods of Statistical Analysis," 2nd ed., chap. 8, Wiley, New York, 1952. (6) Janz, G. J., and Duncan, N. E.. -4s.~~. CHEY.,25, 1410 (1953). (7) Kolthoff, I. hl., and Limgane. .J. J.. "Polaroaraghy," 2nd ed., vol. 2, p. 425, Interscience, S e w York, 1952. (8) Ibid., p. 720. (9) Petersen, G . IT., and Radke, 11. H., ISD. Esc. C m x . . rlx.4~. ED.,16, 63 (1944). (IO) Smeltz, K. C., and Dyer, E., J . .4m. Chem. SOC..74, 6 3 (1952). (11) Spillane, L. J., AXAL.CHEM.,24, 587 (1952). (12) Strause, S. F., 11,s.thesis, University of Deiaware. Sewark, Del., 1953. (13) Terent'ev, A. P., BuDskus, P. F., aud Yashunskii. V. G., Zhur. Anal. Khim., 9, 162 (1954). (14) Van Rysselberghe, P., Alkire, G. J., and McGee, J . 11..J . Am. Chem. Soc., 68, 2050 (1946). RECEIVED for review October 18, 1954, Accepted August 5 , 1'95:.

Polarographic Behavior of Monohaloacetones Analysis of Mixtures PHILIP J. ELVING' and ROBERT E. VAN ATTA'

The Pennsylvania State University, University Park, Pa. Iodoacetone, bromoacetone, and chloroacetone each give one polarographic wave in the pH range 1.5 to 10. Rapid hydrolysis of the compounds prevents measurement at higher pH. The reductions are pHindependent, irreversible, and diffusion-controlled; a two-electron transfer is indicated for the reactions resulting in each wair, with the final product being acetone, which is nonreducible in the potential range investigated under the conditions used. The polarographic behavior of the three compounds is summarized and compared w-ith respect to the factors of pH, ionic strength, buffer component nature and concentration, ketone concentration, temperatui e, and drop time. The factors involved in the simultaneous analysis by polarographic measurement of mixtures of the three haloacetones with satisfactory precision are discussed. Sodium acetate-acetic acid buffer (pH 4.6, ionic strength 0.5) is reconimended as the preferable medium for such an electrollsis. Within the allowable limits of concentration, mixtures of all three ketones can be analyzed with an error of less than 2 relative %.

T

HE polarographic behavior of the three monohalogenated acetones-iodoacetone, bromoacetone, and chloroacetonehas been investigated in order to determine the relationship of half-v-ave potential, E l / %and , diffusion current, i d r to pH, ionic strength and buffer concentration effects, the probable nature of the reduction, and the extent of applicability of polarography as a method of analysis for such compounds, individually and in mixture. 1 2

Present address, University of Michigan, Ann Arbor, Mich. Present address, Southern Illinois University. Carbondale. Ill.

The ketones have been studied polarographically in unbuffered solution by Winkel and Proske (16), who gave reduction potentials and claimed that the introduction of halogen made possible reduction of the carbonyl group, which in acetone itself is not reducible polarographically. Pasternak and Halban ( l a ) , without giving experimental conditions or data, stated that benzyl bromide and chloride give waves of pH-independent El/, resulting from two-electron processes. On the basis of this pH independence, they surmised that the halogen is being reduced in the haloacetones and not the carbonyl group. The present investigation serves to substantiate the latter opinion in regard t o the fundamental feature of the haloacetone reduction. EXPERIMENTAL

Apparatus. Sargent Model XI1 and X X I polarographe and a Leeds and Northrup Type E Electrochemograph ere wed. Potential measurements with the Sargent instruments were checked with a potentiometer. Beckman Model G and H-2 pH meters were used for p H measurement; a Type E electrode was used above p H 10. Capillaries used for the dropping mercury electrodes were prepared from Corning marine barometer tubing; m values (open circuit in distilled water) were: a t 25" C., 1.02 mg. per second a t 60 em. of mercury and 1.58 mg. per second (90 em.); a t 25' C., 1.39 nig. per second (45 em.); a t 0" C., 1.09 mg. per second (55 em.). Jacketed H-type polarographic cella ( 7 ) containing reference saturated calomel electrodes were used; water a t 25.0" i 0.1" or 0.0" =I=0.1' C. was circulated through the jacket, Resistance of the cell solution system was measured with a General Radio Co. Type 650-A impedance bridge. -411 potentials are reported us. the saturated calomel electrode and are corrected for internal resistance drops. Reagents. Iodoacetone was prepared from Eastmsn Kodak practical grade chloroacetone by an exchange reaction with potassium iodide in 95% ethyl alcohol solution. Vacuum distillation of the reaction mixture after removal of the bulk of the solvent yielded a product with the following physical properties: boiling point (4 mm.) 45" C., d i 5 2.0324 [literature values (2,

V O L U M E 27, NO, 1 2 , D E C E M B E R 1 9 5 5 Table I. Sulnher

Composition of Buffer Solutions" PH 1.3

4 . tj 4 lj 7.0

8 2 8.2

8.' 8.8

9 . .i 10.5 11 5 1 2 ,.i 12 6

Components 0 4521 0 50.M 0 10.11 0 27X 0 18111 0 50.11

0 10.11 0 5o.w

0 50.11 0 1GM 0 141.1

0

low

0 4511

a Nonhuflpred background iolutions mere prepared using XHICI. LiC1, a n d KC1 as electrolyte. Their p H was between 4.6 and 5.0 when chloroacetone w.ns prwent, between 3 . 3 and 3.6 for bromoacetone, and between 3.5 and 3.8 for iodoacetone.

1 3 ) : boiling point (11 mm.) 518.1~ C., dl6 2.1iI. Bromoacetone (Jasonols Chemical C o ~ p and . the Delta Chemical Works) was vacuum distilled: boiling point (13 nim.) 38" C., d i 5 1.6132 [literature (3, 8): boiling point (8 mm.) 31.5" C., d Z 3 1.6341. Chloroacetone was prepared from pure acetyl chloride by reaction with diazoniethane and dry hydrogen chloride in ether solution ( I S ) ; distillation of tlie reaction mixture yielded a product of boiling point (736 nim.) 118-9" C., d i 5 1.1260 [litcrature ( 1 ) : boiling point 118-20" C., d:' 1.1231. Stock aqueous solutions of the ketones (npproxiiuately 5mM) were relatively stable to hydrolysis over periods of 3 or 1 hoiirs when prepared by the initial dissolution of the ketone in a minimal quantity of pure acetone before dilution with water, and when stored out of contact with light; the acetone concentration in each stock solution was not greater than 1Om.U. Acetone in cell solution cmicentrations up to 5OmJf \\-as polarographically nonreducible under all conditions used in the investigation. Buffer solutions (Table I ) were prepared by initial dilution to within a few milliliters of thr desired final volume of calculated amounts of analytical reagent grade chemicals, plus sufficient potassium chloride, where necessary, to give the desired ionic strength. The solution p H was then adjusted to the desired value by careful addition of huffer components, followed by dilution t o t,he final volume. Ionic strength values wcre calculated in the usual manner, using the ionization constants listed in the literature. Sitrogen used for deoxygenating was purified by bubbling it through concentrated sulfuric acid, alkaline pyrogallol, distilled water and, finally, a portion of the test solution being investigated. Procedure for Polarographic Examination. The test solution was prepared by mixing an accurately measured volume of stock stmdard ketone solution wit'h the buffered or nonbuffered background solution. The test solution in the H-cell was purged for 5 minutes with nitrogen and then electrolyzed; a nitrogen atmosphere was maintained above the solution throughout the electrolysis. Values of t , the drop life, determined for the limiting current portions of each polarogram, were used to correct i,i \-dues for the effect of the electrocapillary curve where necessary.

1909 1.1 X 10-5 for chioro-, bromo-, and iodoacetoiie, respectivelj.; the corresponding n values, based on i d a t various experimental conditions and calculated from the IlkoviE equation, are 1.5 & 0.2, 1.7 & 0.2, and 1.7 i~ 0.2. (These values are, of course, open to considerable error because of the method and assumptions made in estimating D values.) Delayed maxima were observed in the neighborhood of -0.5 volt a t higher concentrations of iodoacetone-Le., in all cases the maximum in question appeared approximately 0.2 to 0.3 volt' beyond the point where the limiting current was first reached. No attempt was made to suppress these maxima, since in no case did their presence affect the wave itself (this was carefully checked) or int.erfere with the measurement of the limiting current; they did not appear in polarograms of mixtures. The maxima resembled a small water wave, but did not occur in the right location for a lvater wave and were obviously related to the concentration of iodoacetone, being most sharply defined with the higher concentrations of iodoacetone [see (10) for a discussion of similar maxima], As t,he concentration was reduced from 0.5 to 0.3mM, the height of the maximum diminished until, a t the latter concentration, it v a s just barely in evidence. In general, the maxima Tvere great'est in the acid region, although present throughout the p H range. Similar maxima have been encountered with iodoacetic acid ( 5 ) . At concentrat'ions of bromoacetone greater than about l m 3 1 maxima develope 1, which were greatest in magnitude in the acid region. Current-Controlling Processes. Ratios of the currents at different drop times and at different temperatures were of the magnitudes expected for diffuciori-controlled processes. The current ratios at two drop times-i.e., heads of niercur>--in ench of buffers 1 t o i were 1.24 i 0.0-4, 1.24 5 0.02, and 1.28 i 0.02 compared to calculated values of 1.00 for a kinetic-controlled process, 1.23 for diffuFion control: and 1.52 for adsorption control; these are based on h values of ti0 and 90 em. where the hackpressure corrections were 1.8 and 1.7 cm., respectively. The percentage temperature coefficients for the current change between 0" and 25' C. in buffers 2 and 5 n-ere 1.5 m d 1.5$2.2 and 2.0, and 2.5 and 2.2 for the chloro-, bromo-, and iodoketone3, respectively; temperature coefficients were calculated by the usual compound interest formula. For all three ketones, El/* values a t 0" C. were slightly more negative tliaii those at 25" C.e.g., 0.02, 0.0Eil and 0.03 volt for the chloro-, hromo-, and iodoketones, respectively.

0.101

'

I

POLAROGRAPHIC REHAVIOR

1~::u.liketone gave one well defined cathodic wave in the p H range ;.overed, which was 4.6 to 10.5 for chloroacetone and 1.5to 9.5 for bromo- and iodoacetone. S o wave was found a t higher pH Irecause of hydrolysiy of the compounds, probably to hydroxyaceto:ie which is not reducible in the potential range investigated, although ITinkel and Proske ( 1 7 ) reported it to be polarogasphicslly reduced a t approximately - 1.8 volts. 1\Ieasurements in the vicinity of t,his potential could not be made because of the concurrent or prior discharge of hydrogen ion or the cation of the background electrolyte. S o chloroacetone wave was found a t p H 1.5 because of the confluence of t'he wave with a nat'er wave ( 1 1 ) anti subsequent hydrogen ion discharge. The n value (number of electrons t'ransferred per molecule of haloketone reacted) approximates 2 over the entire p H range for all t'hree ketones. Values of D were estimated by comparison of the molecular volumes of the haloacetones with those of the corresponding haloacetic acids as being 1.2 X 10-6, 1.2 X 10-5, and

3

1

I

5

7

I

9

11

PH

Figure 1. EI/,-pH relationship for iodoacetone at 25' C. and ionic strength 0.5

Variation of Ell2 with pH. The E1,,-pH relation for each ketone is a straight line, indicating pI-1 independence of the reduction process over the p H range inve*tigated. Elil is -0.14 volt for iodo-, -0.35 for bromo-, and -1.15 for chloroacetone. There are, hoirever, two significant exceptions. At p H 8.8 and 9.5, Eli2 for iodoacetone (Figure 1) becomes more negative by approximately 0.03 volt per pH unit; theqe values have been checked repeatedly. Such behavior may indicate interaction between the electroactive species and ammonia, as the ammonia concentration in the buffer increases rapidly from p H 8.2 to 9.5 and a t the latter p H is approximately 0.8J.f. If the interaction product were more difficult to reduce than the noninteracted ketone, the change in El/zmight reflect the extent of interaction. Imine formation which has been postulated for ketones in ammonia solution Tvould probably result in a species possessing two

ANALYTICAL CHEMISTRY

1910

Table 11. Polarographic Analysis of Mixtures of Three Monohaloacetones in Acetate Buffer of pH 4.6 and Ionic Strength 0.5 Mixture

NO. 1 2 3

A A’b

Taken, mmole

0.89 0.92 0.181 0.237 0.237 0.92 0.89 0.92 0.89 0.096 0.181 0.181 Average

Chloroacetone id Found’, Ha. mmole 3.90 0.88 4.03 0.91 0.78 0.18 1.03 0.23 0.23 1.03 0.90 3.95 0.89 3.92 0.93 4.12 0.91 4.01 0.411 0.09 0.79 0.18 0.78 0.18

Error,

Taken, mmole

Bromoacetone id Founds, pa. mmole

1.1 1.0 0.6 3.3 3.3 2.0 0.0 1.0 2.2 6.7

0.58 0.59 0.121 0.59 0.58 0.121 0.121 0.59 0.58 0.121 0.121 0.121

2.96 2.99 0.62 2.99 3.03 0.61 0.59 2.92 2.85 0.61 0.61 0.59

%

Error,

Taken, mmole

0.0 0.0 0.9 0.0 1.7 0.9 0.9 3.4 3.4

0.65 0.62 0.123 0.62 0.65 0.62 0.65 0.122 0.131 0,250 0.131 0,122

%

0.58 0.59 0 12 0.59 0.59 0.12 0.12 0.57 0.56 0.12 0.12 0.12

Iodoacetone id Found“. pa. mmole

3.24 3.07 0.60 3.09 3.19 2.93 3.06 0.62 0.63 1.27 0.64 0.60

0.65 0.62 0.12 0.82 0.64 0.59 0.61 0.12 0.13 0.26 0.13 0.12

a

10

Error,

%

0.0

0.0 0.6

0.0 1.5 4.8 6.1 C 1.6 C’ 0.8 D 0.9 4.0 D’ 0.6 0.9 0.8 E 0.6 0.9 1.6 1.4 1.2 1.8 Values based on diffusion current constants found in pure solutions of individual ketones a t 0.5m.M concentration level: 2.94 for chloro-, 3.29 for bromo-, and 3.26 for iodoacetone. b Polarograms designated b y primes shown in Figure 5. B B’

reducible groups, although the imine reduction wave may be masked by the background electrolyte discharge. Unfortunately, the rapid hydrolysis of iodoacetone in this p H region prevents investigation a t higher p H values to determine whether the change in ,TI/*with pH is a true pH dependency or the result of an interaction with the buffer components. Perhaps a more probable explanation for the apparent p H dependency is that El/z appears to shift, because of geometrical interference of the crest of the chloride wave with the foot of the iodoacetone wave, which immediately follows this chloride wave. An analogous case is that of the effect of the iodoacetone wave on the bromoacetone wave in mixtures of the three ketones; for mixture 1 of Table 11, Eli2 for bromoacetone was approximately 0.05 volt more negative than when measured in a solution containing bromoacetone only, whereas El,2 for iodoacetone in the same mixture becomes more positive by 0.02 volt. El/zvalues for bromoacetone determined in phosphate buffer (pH 7.0 and 8.2) were approximately 0.08 volt more positive than Eliz determined in other buffers of corresponding pH. This effect is associated with the presence of phosphate. -4s the ionic strength of the medium is increased by adding potassium chloride, this “phosphate effect” is diminished until, a t an ionic strength of 2.0, E112 values in the phosphate buffers agree with those determined in other buffers a t the same ionic strength. Such behavior was not observed with chloroacetone, but is evident to a slight degree with iodoacetone. In the latter cabe, E112 values in phosphate buffer a t pH 7 and 8 are slightly more negative than the corresponding values in other buffers. A possible explanation, based on interaction of phosphate with ketone to form a more easily reduced species, is not supported by the experimental data, inasmuch as the effect appears to be independent of phosphate concentration. I n addition, such an interacted species would necessarily have to show behavior with increased ionic strength, which is practically identical with that of the noninteracted form, in order to produce the Ells values obtained a t ionic strength 2.0. Variation of Diffusion Current with pH. The diffusion current constant ( I = i d / C m 2 ’ 3 t l ’ * ) for chloroacetone decreases by about 6% from p H 4.6 to 8.2, above which it decreases rapidly to 0 a t p H 11.5 as the rate of hydrolysis increases with increasing alkalinity of the solution (Figure 2). Ammonia buffers give I values about 10% greater than do other buffers with a similar tendency to decrease as pH increases. Bromoacetone I values decrease about 15% as pH increases up to 8.8, and then decrease rapidly to 0 a t pH 10.5 because of increased hydrolysis rate. Buffer 5 (pH 8.2) produced a greater value for I than all other buffers investigated except buffer 1. I for iodoacetone increases slightly from pH 1.5 to 4.6 and sharply (about 10% up to pH 8.2, after which it drops rapidly t o 0 a t p H 10.5 because of hydrolysis. Values in nonbuffered

30

-

14.0) 2.0 (3.0)2.5

I

4.0 -

(2.0)

3.0

0

2

6

4

PH

Figure 2. Effect of pH on diffusion current constant values for monohaloacetones at 25’ C. and ionic strength 0.5 A . Iodoacetone B . Bromoacetone ( I scale in parentheses) C.

0 0

Chloroacetone Nonbuffered solutions Solutions buffered with ammonia buffers

solutions are approximately 5% greater than those in buffered solutions in the same pH region. Effect of Ketone Concentration. The variation of I with concentration is rather unexpected (Figure 3). I for iodoacetone is relatively constant for each buffer over the concentration range investigated. I n the case of bromoacetone, I is constant in buffers 1 and 5, but decreases appreciably with decreasing bromoacetone concentration in the other buffers. With chloroacetone, I decreases appreciably with decreasing concentration for buffers 2 and 5 ; with all other buffers a slight decrease is observed from 0.50 to 0.25mM chloroacetone, followed by a relatively sharp increase when the concentration is further reduced to 0.lOmM. With all the buffers used, increasing the ketone concentration causes El/* to become slightly more negative, approximately 0.02 volt. This effect is probably a result of the irreversible character of the reduction process and the accompanying unsymmetrical nature of the waves obtained. Such behavior did not occur in the analysis of mixtures of the three compounds,

V O L U M E 2 7 , NO. 1 2 , D E C E M B E R 1 9 5 5

1911

effect on Eli2 of iodoacetone, especially in the region of ionic strength 0.5 to 2.0. The tendency a t lower ionic strength is for the reduction to become less difficult with increasing ionic strength. I increases slightly as ionic strength is increased except in nonbuffered solution, where the reverse is true. The 2 effect of decreased buffer component concentration is noticeChlnroaceione able; Eli2 is about 0.02 volt more negative in buffer 2a than in 4 buffer 2, whereas the average I values are about 20% greater over the ionic strength range investigated. Such behavior may indicate interaction between active species and buffer component. 3 Specifically, in buffer systems with low buffer component concentrations, the postulated interaction might remove enough of t,he reacting buffer component (presumably the acid or base) 3 so that the background electrolyte becomes essentially the salt component. Under such circumstances, the resulting behavior I might well resemble that of a nonbuffered system, especially in view of the relatively small amount of ammonia or acetic acid 2 needed to give the required p H for a 0.1M buffer. The fact that the Ell2 values in the lower buffer component concentration Bromoccetone runs were somewhat more negative than in nonbuffered run8 at 4 approximately the same p H lends some credence to this theory. -5 Smaller buffer Component concentration and increasing potas4T sium chloride concentration enhance the possibility of behavior 3 of the type indicated. Further support for the postulated behavior is afforded by the 01 values for the wave in buffer 2a (1.25 i 0.12), which agree almost exactly with those for non3 buffered solution (1.19 i 0.08), whereas corresponding values in L 7 0 . lodoacetone buffers other than 2a are considerably smaller (0.88 ==! 0.13). 01 0 2 0 3 0 4 0 5 E112 values in buffer 5a are about 0.03 volt more negative than C. rnM those in buffer 5, whereas I values are of the same order of magnitude. Figure 3. Relation of diffusion current constant values to concentration of monohaloacetones at 25" C. and ionic strength 0.5 3

GSi-"

I

Numbers refer t o buffer solutions as designated in Table I.

I

h

Chloroacetonr I ia-

where unsymmetrical wave shapes could be more easily accounted for-Le., the limiting current portion of one wave becomes the introductory portion of the following wave, etc. (Table 11). Effect of Ionic Strength. Increasing the ionic strength of the solution renders chloroacetone more easily reducible throughout the p H range investigated (Figure 4), while the I values, in general, tend to increase slightly; similar behavior is observed in nonbuffered solutions where the ionic strength was varied by addition of the background electrolyte. These effects are the reverse of the behavior of bromoacetone under similar conditions and are more pronounced than in the case of iodoacetone. The effect of lower buffer component concentration is slight but definite. R7ith buffers 2a and Sa, E112 for chloroacetone is approximately 0.01 volt more negative throughout the range of ionic strengths studied than with buffers 2 and 5. I values for buffer 2a are slightly greater than those for buffer 2, whereas they are of about the same magnitude for buffers 5 and 5a. There is no significant change with increasing ionic strength in LY, determined as being 0.48 =t0.04 from the slope of the wave (Eli4 - Ea/, = 0.056/, a t 25" (3.). Increase in ionic strength in both buffered and nonbuffered solution causes El/*for bromoacetone to become more negative, whereas I generally decreases. In buffers 2a, 3, and 4, I increases slightly with increasing ionic strength. Changes in buffer component concentration have little effect with buffers 2 and 5, indicating little if any interaction between reducible species and buffer components. There is no significant change in 01 (0.70 =!= 0.05) with increasing ionic strength except for phosphate buffers 3 and 4, where LY a t ionic strength 0.5 is appreciably smaller (0.56) than with any other buffer a t the same ionic strength. Such behavior may account, in part, for the discrepancy in observed with these buffers a t low ionic strength. Increase in ionic strength of the medium has no appreciable

'I

1.14-

-E

1.10

-

0.40

I-

0.36

-

1%

Bromoacrtonr

-

0.32

0.28

U

0.20 ' lodoacetone W

0.1

I .o

0.5

I .5

2.0

P Figure 4. Effect of ionic strength and buffer component concentration of half-wave potential of monohaloacetones at 25" C. Numbers refer t o buffer solutions as designated in Table I; N refers t o nonbuffered ammonium chloride solution; a n d half-shaded points are common t o two or more curves.

1912

ANALYTICAL CHEMISTRY

Reduction Mechanism. On the basis of the observed polarographic behavior, the reduction of a monohaloacetone a t the dropping mercury electrode is an irreversible, pH-independent, diff usion-controlled process involving the fission of the carbonhalogen bond with acetone as the reduction product. This conclusion is substant,iated b y the similarity to the polarographic reductmionof other halogen-containing carbon compounds ( 4 ) , as well a.s by the electrolytic reduction of chloroacetone to acetone in hydrochloric a.cid solution with graphite or lead electrodes ( 1 4 ) and the chemical reduction of chloroacetone to acetone by zinc and hydrochloric acid (9). The reduct,ion process probably involves fission of the carbon-halogen bond to form a carbanion:

0

0

I/

+ 2e-

CHs-C-CH2-X

jl

-+

CH3-C-CHs-

+ X-

(1)

followed by the rapid combination of this species with hydrogen ion from the solvent:

0

0 CHh-C-CH2-

+ HT

I1

--t

CH3-C-CHs

i2 )

This mechnnisin accounts for the p H independence of the reduction processes, since Reaction 1 is the slow, pqteatial-determining step with the electrons acting as a displacement reageiit. (It is immaterial whether the elect,rons add to the carbon atom simultaneously or in rapid succession and IThether Equation 2 involves H + or HzO.) The two principal energy steps are the rearward approach of the electrons to the alpha carbon and the simultaneous fis~ionof the carbon-halogen bond. The formation of the carbon-hydrogen bond (Equation 2) does not figure in the controlling electrode reaction-Le., the reduction is p H independent. Increasing the electron densit'y on the alpha carbon increases the energy required for the first phase of the reaction. The electron density at the alpha carbon is determined by the substitue!it halogen; therefore, the relative difficulty of reduction should 1)e: chloroacetone>hromoacetone>iodoacetone; the espcrimental results verify this conclusion.

sired wave; buffers 5 , 6, and 7 are eliminated because of geometrical interference of the chloride wave of the buffers with the iodoacetone wave; and buffers 3 and 4 cannot be used since the more positive wave for bromoacetone in these media results in inadequate separation of the iodo- and bromoacetone waves. Thus, acetate buffer 2 is the only logical choice for use in siniultaneous analysis of mixtures of the three compounds. The acetate buffer used ITas prepared of such ionic strength as to produce a test solution of ionic strength 0.5JIwhen mixed with the sample. A number of mixtures of the three haloacetones were analyzed n-ith the results shown in Table I1 and Figure 5 . The concentratione of the three ketones used in mixture 1 are approximately the highest concentrations of the three ketones which do not result in maxima effects in the waves. The concentrations shown for mixtures 3 and E are approximately the minimum concentrations measurable without sacrifice in accuracy through intcrference of bromo- with iodo- or iodo- with bromoacetone waves. The I values for each of the three m v e s in the mixtures remain relatively constant over the range of conceiitrations investigated; the average deviation of these values from the I value in 0.5m3f sol itions of the individual ketones is of the order of 2% or less; thiiq. the most important criterion for analytical allplication of the method is satisfied. -1lthough a relatively large shift due to the geometry of the wave3 is observed for the E1.2 of bromoacetone, no adverse effect is introduced with regard t o the current, measurementq. I

AA-ALYTICAL APPLICATIOYS

Determination of Individual Haloacetone. The data cited on the E,.? and I behavior of the haloacetones serve to indicate the general feasibility of determining any of the monohaloacetones when present in admixture with other substances, some of which may be also polarographically reducible. Such data also indicate the critical conditions involved in developing a specific analytical procedure. I n general, any of the buffers studied are satisfactory for analytical work involving the determination of iodoacetone, ivith buffers 2, 4, and 6 being preferred on the bmis of niininium I variation with concentration changes and ionic strength effects. On a similar ha+ buffers 1 and 5 are best suited for polarographic analytical n-ork involving bromoacetone. Sone of the buffers investigated are ideal for analytical n-ork with chloroacetone; on comparison of ioriic strength effects and constancy of I values n-ith ketone concentration change, buffer8 3 and 4 ivould he most suitable, Analysis of Mixtures of Haloacetones. As suggested by Winkel and Proslre (16 j, it is possible to analyze mixtures of iodoacetone, bromoacetone, and choroacetone by the polarographic method. The major problem in the analysis of such a multicomponent mixture is the selection of a suitable background electrolyte for the electrolysis. It is desirable t o choose a system in which the waves are sufficiently well defined for accurate measurement and in which the diffusion currents obtained are, if possible, directly proportional t o concentrations of the electroactive species. I n the case of the haloacetones, it was not possible to choose a mediiim in n-hich both of these conditions would be perfectly met. The chloroacetone wave cannot be obtained in buffer 1 because of the hJ-drogen discharge preceding the de-

Figure 5 . Polarograms obtained for various m i x t u r e s of monohaloacetones Letters refer t o mixture as designated in Table 11.

Possible Interferences. Interference due to the possible presence of polarographically reducible compounds, other than the three ketones, in any solution for which haloketone analysis is desired, may be readily predicted from the El,*values for such compounds. For example, iodoacetic acid and ethyl monobromoacetate would interfere whereas the bromo- and chloroacids, chloroestere, and chloroacetaldehyde n-odd not. Acetaldehyde, the only polarographically reducible parent of any of the halogenated compounds listed, would not interfere since i t does not give a polarographic wave in the acid p H region (6); in any case, the wave for acetaldehyde occurs a t a much more negative potential ( - 1.7 volts) than the chloroacetone wave and the purging of the solution with nitrogen before electrolysis would probably remove the bulk of the aldehyde. AAALYTICAL PROCEDURE

Measure out a sample containing not more than 10 mg. and not less than 2 mg. of each of the three ketones (iodo-, bromo-, and chloroacetone). Transfer the sample to a 100-ml. calibrated volumetric flask, add 70 ml. of 0.7M sodium acetate-acetic acid buffer (pH 4.6), and dilute t o the mark with distilled water.

V O L U M E 2 7 , NO. 1 2 , D E C E M B E R 1 9 5 5 Rinse the cell and electrode several timefi with the solution to be analyzed. After purging for 5 minutes with nitrogen, electrolyze the solution over the potential range of 0 to -1.5 volts us. the saturated calomel electrode, maintaining the nitrogen atmosphere above the solution. If the electrocapillary curve for the base solution is not known, note t, the drop time, a t potentials of -0.2, -0.5, and -1.5 us. S.C.E. Run a Pimilar curve on the base solution and, if necessary, correct the sample curve for the latter curve. Using the intercept method, determine the diffusion current for each of the three waves. Calculate the amount of each ketone present from its diffusion current constant, C = id/Im2W6, where C is the ketone concentration in millimoles per liter, i d the measured diffusion current in microamperes, Z the diffusion current constant for the ketone (Table 11),m the mass of mercury in milligrams per second flowing from the capillary, and t the drop time in seconds measured at -0.2, -0.5, or - 1.5 volts. The weights or percentages of the three ketones present may then be calculated from the millimolar concentrations by conversion to grams and correcting for the dilutions involved. .4CKNON LEDG\IE\-T

The authors n ish to thank the -1tomic Energy Commission which supported the xvork described. LITER4TURE CITED

(1) Buchman, E. R., and Sargent. H , J . Am. Chnn. Soc., 67, 401

(1945). (2) de Clermont, P., and Chautard. P., Bull. soc. chim Paris, 43,

614 (1885).

1913 (:?)

(4) (5) (6)

(7)

(8) (9) (10)

"Ilirtioiiary of Organic Compounds," (Heilhron, I.. editor), rol. 1, p. 264, Oxford Univ. Press, New York. 1946. Elving, P. J., Record Chem. P ~ O Q (Kresge-Hooker T. Sci. Lab.), 14,99 (1953). Elring, P. J., Rosenthal, I., a n d Kramcr, 11.K., J . Am. Chem. Soc., 73, 1717 (1951). Elring, P. J., and Rutner. E., ISD. ENG.C H E M . . - 1 s . 4 ~ .E n . . 18, 176 (1946). Komyathy, J. C., llalloy, F., and Elving. P. .J.. . ~ L L C. H F x , 24, 431 (1952). Levene, P. A., Org. Synchesea, 10, 12 (1930). Linnemann, E., Ann. Chim., Justus Liebigs. 134, 170 (1863). Aleites, L., "Polarographic Techniques," pp. 137-9. Interscience,

Xew York. 1955. (11) Orlemann, E., and Kolthoff, I. AI., J . Am. C k e m . 9 o c . . 64, 1070

(1942).

€I. yon, Hela. Chirn. A d a . 29, 190 (1946). (13) Scholl, R., and Matthaiopoulos, G., Ber. deut. chem. Ges.. 29,1557 (1896). (14) Saper, J., Bu2l. soc. chim., 51, 653 (1932). (15) Van .Itts, R. E..Zook, H. D., and Elving, P. J . . J . A i m , C h c m . SOC.,76, 1185 (1954). (16) Winkel, A., and Proske, G.. Ber. deut. chem. Ges , 69, 693 (1935). ( l i ) Ibid., p. 1917. (12) Pasternak, R., and Halban,

RECEIVED for review April 23, 1955. .Iccepted Septemher 18. 1935. .Ihstracted from a thesis presented t o T h e Pennsylvania State College in .lugtist 1952 h y Robert E. Van A t t a i n partial fulfillment of requirements for Ph.D. degree. Detailed tables of d a t a covering t h e polarographic hehavior of the three haloacetones are available from the senior author.

Polarographic Nitrate Determination RANDALL E. HAMM and C. DEAN WITHROW Department o f Chemistry, University o f Utah, Salt Lake City, Utah

A new- polarographic wave has been found for the induced reduction of nitrate ion when present in a chromium(II1)-glycine complex solution. The current measured from this waw is proportional to the nitrate ion concentration. Procedures have been developed for determining nitrate in the concentration range 1.0 X 10-6 to 2.5 X mole per liter, with as good or better precision than the older polarographic methods for nitrate. The effects of gelatin, pH, and interfering ions on this reduction wave have been studied. The nitrite ion gives a similar reduction ware. As a result of the study of nitrate and nitrite reduction a rough calculation has been made of the number of electrons involved in the reduction processes.

A

S C M B E R of polarographic methods (1, 4-73 for the determination of nitrate by means of reduction induced in the presence of polyvalent cations have been described. The original observation of Tokuoka ( 8 ) and Tokuoka and Ruzicka (9) showed that the reduction potential of nitrate TTas shifted to more positive potentials when certain polyvalent cations tvere present. Lanthanum (9),uranyl ( 5 ) : molybdate ( 1 , 4),zirconyl ( 6 ) , and cerium ( 7 , 9 )ions h a r e been studied. Collat and Lingane ( 2 ) have recently niade a more thorough study of the reduction products in the case of lanthanum, cerium, and uranyl induced reductions of nitrate and have reported that nitrate can be reduced directly from acid solutions. I n the coime of a study of the complex ions of chromium(II1) with glycine it n-as observed (3) t h a t a chromium(II1)-glycine complex, prepared by aging a 1 to 3 mixture, gave an induced nitrate reduction 1%-arewhich had some unique properties that should make it especially valuable for the determination of ni-

trate. This paper is a description of the polarographic nictliotl which has been developed as the result of this new wave. Further work is in progress in this laboratory in an attempt to esta!)lish the nature of the processes that are responFilIle for the unusual shape of the wave; ho\i-ever, the present n.ork has estahlishetl a reproducible and highly sensitive method for the tletPrniination of nitrate or nitrite. EXPERIMENTAL

Equipment. h Sargent, Model XXI, polnrograph n-as used throughout this investigation. A Beckman, Xlodel G. pH meter, calibrated with 0.05.V potassium acid phthalate solution, was used for making all pH measurements. Analyses were performed in an H-type cell which had a saturated calomel reference electrode in one side, the experimental solution in the other side, and a saturated potassium chloride-3% agar bridge in the cross member. This cell was suspended in constant temperature bath which Tvas maintained at 25.0" =k 0.1 C. for making all measurements. Dissolved osygen Tvas removed from the esperimentsl solut,ion hy bubbling purified kiydrogen through the cell for about 5 minutes before the recording of a polarogram. An external potentiometer circuit v a s used for determination of the exact potential applied to the cell, wherc measurements of potential %-eretaken. The dropping mercury electrode wed was n piece of marine barometer tubing with the characteristics, 1112 3!iii; = 1.G3 nig.2'3 sec. -l'z. Reagents. All of the reugents used except the following were reagent grade chemicals. Chromium perchlorate, hydrated Cr(C10,)7, 0.Frederick Smith Chemical Co. Glycine, Eastman Kodak (white label). Gelatin, Baker and Adamson. A stock solution which was 0.10036 in potassium nitrate xms prepared from a dried and weighed sample of analytical reagent grade potassium nitrate. More dilute solutions Tvere prepared from this bv normal volumetric t,echniques. A stock solution of sodium nitrite which was approximately 0.10M in nitrite was