Spectrophotometric Analysis of Sulfuric Solutions of Hydrogen Peroxide, PeroxymonosuIfuric Acid and PeraxydisuIfuric Acid Maria H. Mariano Luboratdrio de Fisica e Engenharia Nucleares, Sacavbm, Portugal
A simple spectrophotometric method of analysis of mixtures of hydrogen peroxide and the peroxysulfuric acids is discussed. Ceric sulfate and ferrous sulfate solutions are used as titrants. The effects of concentration, acidity, dissolved oxygen, and cerous ions on the accuracy of the measurement are studied, as well as the variation with time of the composition of the samples for different sulfuric acid concentrations. Peroxymonosulfate interferes markedly with the cerimetric titration of hydrogen peroxide in diluted HzSO4 solutions. This interference is minimized in strong sulfuric acid media or in the presence of an excess of Ce(lll) ions. Dissolved oxygen can also induce large errors in the determination of peroxydisulfate with Fe(ll). An analytical procedure in which these different disturbing influences are avoided is outlined. IN OUR STUDIES on the 6oCo?-radiolysis and 5.3 MeV aradiolysis of aqueous solutions we need to determine the radiolytic behavior of mixtures of HzOz, H2S06, and H&0* in a sulfuric acid medium. A method for the measurement of these peroxides in the presence of each other had to be devised. Because of the particular conditions of our irradiation experiments (namely, the utilization of dissolved 210Po as a source of a-radiation) the measurement technique had to satisfy two main conditions: It had to be as simple and as rapid as possible and the behavior of the titrants under radiation had to be known. The method outlined by Boyle (1) attracted our interest because the titrants used are the most thoroughly studied chemical dosimeters : the Ce(IV)/Ce(III) and Fe(II)/Fe(III) systems. In this method, hydrogen peroxyde is measured by adding a known amount of Ce(1V) ions, the remaining ceric ions being determined spectrophotometrically at 320 mp; the total oxidizing power of the sample is determined by adding an excess of ferrous sulfate and measuring the resulting Fe(II1) ions at 304 mp. From these two measurements the sum of the concentrations of the peroxysulfuric acids is obtained. Peroxymonosulfate is then reduced by the use of arsenious acid followed by the determination of the peroxydisulfate concentration with Fe(I1). Since Boyle was not explicit over the magnitude of the errors, we decided to study this technique in more detail. In fact, the cerimetric method of [H202] determination in the presence of the peroxysulfuric acids is open to criticism (2, 3) and it has also been pointed out that oxygen interferes with the Fe(I1) method of peroxysulfate determination (4). The possibility of the composition of the sample changing during the manipulations because of hydrolysis of the peroxyacids also deserves some attention. EXPERIMENTAL The spectrophotometric measurements were made on a Zeiss PMQ I1 spectrophotometer equipped with a thermostated cell compartment. The temperature was measured with an iron-constantan thermocouple with an accuracy, obtained by a least-squares fit, of *O.l "C. The ceric ion molar absorptivity at 320 mp and 25 "C, in 0.4M sulfuric acid, was taken as E = 5580 1.M-1 . cm-l 1662
ANALYTICAL CHEMISTRY
(5). The Fe(II1) ion molar absorptivity at 304 mp, in the same acid medium and temperature, was determined by measuring the absorbances of ferric sulfate solutions of different concentrations. From these measurements the value obtained by a least-squares fit was E = 221 1 f 10 1.M-1. cm-l, in agreement with the determinations of other investigators (1, 6, 7). In sulfuric acid solutions of concentrations greater than 0.4M, the molar absorptivities were corrected for the variation with the acid concentration (7, 8). The temperature dependence of the molar absorptivity of ferric ion at 304 mp was also studied over the range from 14.5 to 30 "C. In this range, E was found to increase linearly with ISM-1 the temperature, a coefficient a = 7.05 . cm-l "C-l being obtained. This value is in good agreement with that obtained by Scharf and Lee in the temperature range from 18 to 30 "C(7). With the exception of the arsenious oxide (Carlo Erba, Italy) used in the reduction of the peroxymonoacid, all chemicals were Merck p. a. products. Hydrogen peroxide was obtained by y-radiolysis of very pure water or dilute sulfuric acid solutions. The peroxymonosulfuric acid was obtained by hydrolysis of a potassium peroxydisulfate solution in 5M sulfuric acid, at room temperature, or by heating a peroxydisulfuric acid solution in 1 M HzS04. In either case, by the time all the peroxydisulfate has been hydrolysed, a fraction of the peroxymonoacid formed has also been decomposed giving hydrogen peroxide. This was destroyed by adding a slight excess of a ceric sulfate solution. The stock solution of peroxymonosulfuric acid was then kept at a temperature of - 10 to - 5 "C to avoid further hydrolysis. In the preparation of the solutions, freshly purified water was used, with a conductivity of about 1 X 10-61t-1 cm-' at 25 "C. The analytical study was carried out on standard solutions of each of the peroxides alone, or in admixtures of known composition, in a 0.4M H2S04 medium at room temperature (between 23 and 26 "C). To avoid errors arising from the hydrolysis of the peroxyacids, the peroxydisulfate solutions were prepared immediately before use and the Hz02 content of the peroxymonosulfate solutions was determined.
RESULTS AND DISCUSSION The Cerimetric Measurement of Hydrogen Peroxide in Presence of Peroxysulfuric Acids. From data given in the literature, one concludes that ceric sulfate is a good selective oxidizing agent in the measurement of H202 since the reaction 2Ce4+ HzOz + 2Ce3+ 2H+ 0 2 is rapid. According to Baxendale (9),only a few seconds are needed for com-
+
+
+
J. W. Boyle, Radiation Res., 17,427 (1962). L. J. Csfinyi and F. Solymosi, Anal. Chim. Acta, 15, 501 (1956). B. D. Sully and P. L. Williams, Analyst, 87, 653 (1962). N. H. Furman, "Scott's Standard Methods of Analysis," Van Nostrand, New York, 1962, p 1015-D. (5) C. J. Hochanadel and J. A. Ghormley, J. Chem. Phys., 21, 880 (1953). (6) C. M. Henderson and N. Miller, Radiation Res., 13, 641 (1960). (7) K. Scharf and R. M. Lee, Radiation Res., 16, 115 (1962). (8) J. W. Boyle and H. A. Mahlman, Radiation Res., 16,416 (1962). (9) J. H. Baxendale, The Chemical Society, Spec. Publication No. 1 , 40 (1954).
(1) (2) (3) (4)
I
I
0
20
40
60
80
100
120
140
160
160
200
220
t (minutes) Figure 1. Compared effects of peroxymonosulfate and peroxydisulfate in the reduction rate of ceric ions
+ + + + +
+ + + + +
+ +
I-A-162 pM HzSOj 4.00 pM K2S208 27.9 p M H 2 0 2 160 pM Ce(1V) I I - ~ - 4 9 8 pM HzSOj 23.7 pM KzS208 27.0 pM H 2 0 z 148 pM Ce(1V) I I I - + 4 9 . l m M K2S208 28.6 pM HzOZ 119 pM Ce(1V) IV-o-0.1M K2S208 27.9 pM H 2 0 2 117 pM Ce(IV) pM H2S05 27.7 p M H2O2 119 pM Ce(1V) IV-0-681
pletion of the reaction even at micromolar concentrations of the two reactants. However, in carrying out the cerimetric titration of hydrogen peroxide in the presence of the peroxysulfuric acids we did verify that the [HzOz]values, as well as the measured concentrations of the peroxyacids, depend greatly on the time taken in the measurement and are sometimes very different from those corresponding to the initial composition of the sample. A typical example of this is shown in Curve I1 of Figure 1 which represents the reduction of a ceric sulfate solution when added to a sample containing 498 micromoles of peroxymonosulfate, 27 micromoles of hydrogen peroxide, and 23.7 micromoles of peroxydisulfate. In the first 2 to 3 minutes a very rapid Ce(1V) ion reduction takes place; the rate of reduction slows down as the reaction proceeds. About 70 minutes later, the analysis of the sample shows that the peroxymonosulfate concentration has decreased to 191.5 micromoles, the peroxydisulfate concentration increased to 37.5 micromoles, whereas the [H202] value, inferred from the number of ceric ions reduced, is equal to 58.3 micromoles. Variations of this kind cannot be accounted for by the hydrolysis of the peroxyacids (see Figures 4 to 6) and attracted our interest. The conclusions we arrived at, in our study of the behavior of mixtures of these peroxides in the presence of ceric and cerous ions, are summarized as follows: The magnitude of the errors observed in the H202 and peroxyacid determinations depends on the relative amounts of peroxyacid, hydrogen peroxide, and ceric sulfate as well as on the time of contact of the ceric sulfate with the sample under study. For a given molar ratio of peroxyacid over ceric ions, the errors decrease with increasing nydrogen peroxide concentration. For given amounts of ceric sulfate and hydrogen
peroxide, the rate of reduction of the Ce(1V) ions increases with the peroxyacid concentration (Figure 1) but, when the molar ratios [H2S05]/ [Ce(IV)I are about 10 or greater, the reaction between Ce(IV) and H202 is slowed down. Thus, the deviations observed in the measured [H~OZ] values may have positive or negative signs depending on the molar ratios of peroxymonosulfate over ceric ions. Peroxydisulfate has a minor effect when compared with that of the peroxymonosulfuric acid, in the cerimetric titration of hydrogen peroxide. In fact, the errors observed in a solution containing 0.1M K&08 are about the same as in a sample containing 6.8.10-4M peroxymonosulfate (Figure 1, Curves IV and V). In contrast to published information ( I , 2), the errors in the determination of the [H202] values were found to arise from the reduction of the Ce(1V) ions by both peroxyacids. This is inferred from results obtained when ceric sulfate is added to solutions of the pure peroxyacids (As seen in Figure 2, the pattern of the curves obtained in this case is similar to those of Figure l.), and from the fact that no reaction between the peroxyacids and hydrogen peroxide was observed to occur. It was found that the number of molecules of peroxyacid decomposed is greater than the number of Ce(IV) ions reduced, and that peroxydisulfate is a product of the reaction between peroxymonosulfuric acid and ceric sulfate. Figure 3 shows that concentrated sulfuric acid or cerous sulfate inhibit the interference of H2SOs in the reaction between Ce(IV) ions and hydrogen peroxide. From this study we concluded that in carrying out the cerimetric titration of hydrogen peroxide in mixtures of H20z and H2S05in a dilute sulfuric acid medium, reliable results will be obtained only if peroxymonosulfate is reduced or excess cerous ions are VOL. 40, NO. 1 1 , SEPTEMBER 1968
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0.450
0.400
0.350
0.300 0
6 0.250
2 2
0.200
0.l 50 20
60
40
80
100
120
140
t (Minutes)
os 00 0.05C
20
40 60 t (Minutes)
80
Figure 3. Comparison between the rates of reduction of ceric sulfate for different sulfuric acid concentrations and in the presence of excess cerous ions
100
0
%Ib
X
Figure 2. Effect of the pure peroxyacids on the reduction rate of ceric sulfate X
loi pM K&O*
A
680 pM HzSOs
o
+ 79.6 pM Ce(1V) + 72.0 pM Ce(1V)
added prior to the addition of the ceric sulfate. Errors arising from the presence of peroxydisulfuric acid can be neglected in those cases where this compound is not in great excess over Ce(1V) ions (Figures 1 and 2). When the reduction of peroxymonosulfate has to be performed and arsenite is used for this reduction, care has to be taken to avoid the simultaneous decomposition of the hydrogen peroxide by arsenite. Thus, with solutions having , H 2 0 2concentrations equal to, or greater than, 8 ~ 1 0 - ~ Muse can be made of a solution of arsenious oxide about lO-3M to reduce (in 5 to 10 minutes) a 4.1OU4Mperoxymonosulfate. For less concentrated H z 0 2 solutions, it is convenient to reduce the concentration of As20a with the corresponding increase in the time of reduction of H2S05 (about 25 min for 1 0 - 4 ~A S ~ O ~1 0 - 4 ~H~so~). The following set of reactions is believed to explain the results obtained on mixtures of peroxymonosulfuric acid, hydrogen peroxide, and ceric sulfate (9, 10):
+
+
+
+ +
+
+
43.0 pM H~SOF, 4.49pM Hz02 38.9 pM Ce(1V) +0.4k”zSO4 idem in 5MH901 43.0 pM HBSOL 4.49pM H202 38.7 pM Ce(1V) $400 pM Ce(II1) 0.4MH2S0i
acid and for the formation of peroxydisulfate. The competition represented by Reactions 5 and 6 explains the variations in the rate of reduction of the ceric ions as a function of the molar ratios [HzS05]/[Ce(IV)]and the inhibition of the reaction between Ce(1V) ions and HZOZfor high values of this molar ratio. The reverse Reaction 1 accounts for the phenomena,observed in the presence of excess of cerous ions and Reaction 1, followed by 2 and 7, explains the results in strong HzSO4 media where the peroxydisulfate formed in 2 is rapidly hydrolyzed regenerating the peroxymonosulfate. A similar mechanism would explain the reduction of Ce(1V) ions by peroxydisulfate, the Reactions 8, 9, and 10 being followed by Reactions 1 to 7.
-
SZOS-f Ce3+
(8)
+ H20 * HSOi $, HSOs-
(9)
SzOs2-f Ce4+ S208-
+
HSOs-
HSOs-
+ Ce4+S HSOs + Ce3+
+ HOz
Ce3+ $, OH
+
+ OH + Ce4++ OH-
HSOI-
0 2
(1)
(6)
(7)
This mechanism accounts for the fact that there is more HzSOsdecomDosed than Ce(1V) ions reduced by this peroxyKO) P. B. Sigler and B. J. Masters, J. Amer. Chem. SOC.,79,6353 (1957).
1664
0
ANALYTICAL CHEMISTRY
Determination of Peroxysulfuric Acids with Ferrous Sulfate. TIME OF MEASUREMENT. Our experiments show that, in 0.4M HzSOa,the total oxidizing power of the solution can be rapidly determined (time of reaction 3 min or less) when [Fe(II)]/[Sz082-]> 50. The rate of reaction increases with the acidity: In a 5M HzS04 medium, the oxidation of Fe(I1) ions is five times more rapid than in 0.4M H2S04. This suggests that the species which reacts with ferrous sulfate is the HS208ion rather than the S2082- ion or, at least, that the rate of reaction between H S 2 0 ~ -and Fe(I1) is greater than that corresponding to SZOS*- Fez+. EFFECT OF OXYGEN AND PH. Only when arsenite is present in the samples does dissolved oxygen interfere with the determination of peroxydisulfate with Fe(I1) (Table I). Errors sometimes as great as 5 5 % can be observed in some cases. The differences between the values obtained in the presence and in the absence of dissolved oxygen decrease as the hydrogen ion concentration increases.
+
,*
2-
e-
--/-c---A+---
The simplest way to explain this interference of arsenite and oxygen in the reaction between peroxydisulfate and ferrous sulfate is by some mechanism such as the following:
+ Fez+.-,HS04- + SOa- + Fe3+ SO4- + Fez+ SO4'- + Fe3SO4- + As3+ As4+ + S04'As4' + + H+ AS'+ + HO? HOz + Fez+ HOP- + Fe3+ HOz- + H+ H20z H 2 0 2+ 2Fe2+ 2 OH- + 2Fe3+
HS208-
--*
(1 3)
--*
+
-+
(14) (15
(16)
(17)
The overall Reaction 14 [the scheme involves the formation of an oxygenated complex of As(IV)] has been suggested to occur in the photolysis of arsenite solutions (11) and would account for the excess of Fe(I1) ions oxidized in the presence of As(II1) and O2 in dilute HzS04. However, such a mechanism does not explain the phenomena observed in strong acid solutions in which the complex formed by As(1V) seems to behave toward Fe(I1) in a way similar to that of the S o h radicals. This may be due to modification of the reactions of the As(1V) species because of protonation in strong acid medium. Hydrolysis of Persulfuric Acids. At room temperature, the aqueous solutions of peroxydisulfate are stable but, at low pH, a net decomposition in HzOzand HeSOs is observed. Figures 4 and 5 show this decomposition as a function of increasing acidity and time (12). It is inferred that the species which undergoes decomposition at room temperature is the HS208- ion:
S208'-
+ H30'
HS208-
HSOs-
+
Ft- HS208-
+ SO8 HS0.i- + HzOz
HSOs-
+ H?O
+
+ HnO
-
-
or, according to Kolthoff and Miller (13): HS208-
-+
HSO4-
+ SO1
(21)
followed by (1 2)
--*
.
- --
(11)
--*
0 2
c.- 0 -
(18) (19)
(20)
(11) M. Daniels, J. Phys. Chem., 66, 1473 (1962). (12) P. D. Bartlett and J. D. Cotman Jr., J. Amer. Chem. Soc., 71, 1419 (1949). (13) I. M. Kolthoff and I. K. Miller, ibid., 73, 3055 (1951).
so4 + Hz0 + HzSOj
(22)
The rate of hydrolysis of HpSOs is also catalyzed by the H30+ ion (Figures 4 to 6). From curves in Figures 4 to 6 it is also concluded that errors in the measurement of either of these peroxides, arising from variations in the composition of the sample during the determination, will be significant only for acid concentrations equal to or greater than 5M. Analytical Procedure. From the study outlined above we concluded that the three peroxides (HzOZ,HzSO~, and HzS208) can be rapidly and accurately measured in the presence of each other, with ceric and ferrous sulfates as titrants, under the conditions which follow: MIXTURES OF PEROXYMONOSULFATE AND HYDROGEN PEROXIDE. The solution is acidified with sulfuric acid until [H2S04]= 5M(avoid the heating of the solution). Hydrogen peroxide is then determined with an excess of ceric sulfate (at 320 mM, E = 6800 1.M-I cm-l) and the total oxidizing power of the sample is measured with ferrous sulfate (at 304 mh, E = 2874 1.M-I cm-l at 25 "C), without any further precautions. The reaction between H2S05 and Fe(1I) ions Table I. Effect of Oxygen and Acidity in Ferrous Sulfate Determination of Peroxydisulfate Solutions in Presence of Arsenite Analysis of Analysis of the the solutions aerated solutions, devoid of oxygen Solution composition [K2S208]X 10-4M [K2SaOs]X lO-4M 1 0 - 4 ~ ~ 10-3714 ~ ~ ~ 0 ~ 1.55 1.02 AS203 0.4M HzSO4 1.36 1 .oo 1.2 x 1 0 - 4 ~K ~ S ~ O ~ 1.65 1.17 1 0 - 3 ~ ~ ~ 2 0 ~ 1.65 1.15 0.4M HzSOa 1.70 1.16 8-44 x ~ O - ~ M K ~ S ~ O ~ 0.92 0.84 10-aMASz03 5M HzSO4 1.68 x 1 0 - 4 ~ ~ ~ ~ ~ 1.76 0 ~ 1.65
+
+
+
+
+
+
+
VOL. 40, NO. 1 1 , SEPTEMBER 1968
0
1665
1
12
t (Hours) Figure 5. Hydrolysis of peroxydisulfate in 5M sulfuric acid A 0
0
+
Total oxidizing power X = peroxydisulfate X = peroxymonosulfate X = hydrogen peroxide
:’i
10
I
0
I
50
I
150
100
I
200
1
250
t (Hours)
Figure 6. Hydrolysis of peroxymonosulfate in 5M sulfuric acid A 0
+
is rapid even at equimolecular concentrations of the two reactants. Use can be made of an excess of cerous sulfate, [Ce(III)]/ [H2S05]5 150, added prior to the cerimetric titration of H202, instead of acidifying the solution. MIXTURES OF PEROXYDISULFATE AND HYDROGEN PEROXIDE, If the adequate excess of ferrous sulfate over peroxydisulfate is used, the measurements will take only a few minutes so that errors arising from hydrolysis will be negligible even with 5M HB04 solutions (in a time interval of 20 min this 1666
ANALYTICAL CHEMISTRY
Total oxidizing power X = peroxymonosulfate X = hydrogen peroxide
error is less than 6 2 in a 10-4M K&Os solution). No particular precautions will thus be necessary. MIXTURES OF H202, H2SOs AND H&08. For solutions containing 8.10-6M to IO-*M of each of the peroxides, the following procedure can be used: Add to a 10-cc aliquot 1 cc of a saturated arsenious oxide solution. After about 10 min, add 1 cc of a 3.10-3M ceric sulfate solution and measure the remaining Ce(1V) ions at 320 mp, thus obtaining the [H202] value. Take 5 cc of the previous mixture, bubble NPfor about 10
min, and add 1 cc of a 10-lM FeS04 solution. After 5 min more of nitrogen bubbling, measure at 304 mp the Fe(II1) ions formed in the reaction of Fe(I1) with the remaining Ce(1V) ions and the S20e2-ions present in the sample. After correction for the Ce4+ Fe2f reaction, this measurement gives the peroxydisulfate concentration in the sample. Measure the total oxidizing power in another 5-cc aliquot by adding 1 cc of a 10-IM FeS04 solution. There is no need for N2 bubbling. The peroxymonosulfuric acid concentration is obtained by difference between the T.O.P. and the sum [H202] [H&05]. In the more acidic solutions, [HZS04] 5M, it is con-
+
+
venient to dilute the sample with water to minimize the hydrolysis of the peroxyacids (avoid an overheating of the solution). ACKNOWLEDGMENT
We are indebted to L. Santos for his advice concerning some valuable preliminary experiments and to M. Coimbra whose skillful technical assistance and understanding of the problems involved in the experimental side of this study have greatly contributed to the achievement of this work. RECEIVED for review March 4, 1968.
Accepted May 21, 1968
A Precision Photometer Using Milliwatt Light Sources and Photon Counting Edward H. Piepmeier, Donald E. Braun,' and Roxie R. Rhodes Department of Chemistry, Oregon State University, Corvallis, Ore. 97331
A precision photometer with digital readout has been built using milliwatt light sources and a pulsedetection system that counts single photoelectron events. The system can be used over more than two orders of magnitude of radiant power before pulse overlap becomes significant. Because the detection system primarily depends upon counting photoelectron pulses rather than upon pulse amplitudes or their integrals, neither the voltage applied to the photomultiplier nor the pulse amplifier gain need be highly stable to achieve precise results. Two colorimetric systems were studied and the results shown to agree with Beer's law.
THEADVANTAGES of a photometer that uses a pulse detection system with digital readout have been discussed by Ross (1). He originally introduced the pulse counting concept in precision photometric analysis by using a radioisotopic light source and a pulse height discrimination detection system. His particular technique takes advantage of the negative exponential distribution of pulse heights observed when the multiphoton pulses from a beta-activated scintillator are detected by a photomultiplier tube. An important characteristic of this system is the very stable light source. However, because the system depends upon distinguishing pulse heights, precise results require that the photomultiplier dynode multiplication system, amplifier, discriminators, and counting intervals be stable. Although common light sources are usually thought of as emitting energy continuously, they are sources of discrete photons. When light from these sources is measured with a detector such as a photomultiplier, the photons are converted to current pulses which may not be completely resolved in time from each other. Usually these pulses are smoothed to what appears to be a continuous or continuously varying signal and recorded by a readout device such as a meter or servo recorder. However, the number of photons per unit time reaching the detector can be decreased to a point where they are resolved from each other. This can be done by del Present address, Department of Chemistry, Pacific College, Fresno, Calif. 93702
(1)
H.H. Ross, ANAL.CHEM., 38,414 (1966).
creasing the intensity of the light source, by isolating a wavelength region of interest with a filter or monochromator, and by stopping down the optical aperture of the light beam. If the photon pulses are also resolvable from the noise pulses so that the noise can be subtracted from the total counting rate to give the photon counting rate, or if the noise is sufficiently low, then a practical detection system can result by counting pulses. This type of system, ideally depending only upon pulse counting rate, and being independent of pulse height, has a number of important advantages over the system used by Ross, or the more common techniques which provide a smoothed current or voltage proportional to the radiant power falling upon a detector. For instance, if individual photons of a chosen wavelength were counted, the counting rate rather than the integral of the pulse heights would be a direct measure of the radiant power reaching the detector, and would be relatively independent of variations that might occur in pulse heights. Therefore, a highly stable detection system would not require that the pulse amplifiers, pulse-height discriminators, and power supplies be highly stable. The stability of the detection system would be primarily dependent upon the stability of the counting time intervals. The stability of a photon counting detection system that used a photomultiplier as the detector would also be dependent upon the stability of the quantum efficiency of the photocathode (2) and the collection efficiency of the dynodes, which can be made to be only a moderate function of the applied voltage (3). Engstrom and Weaver (3) have discussed the ability of a photomultiplier to provide pulse counting rates that are relatively independent of the applied photomultiplier voltage. The stability of a single beam photometer that uses a photon pulse detection system would be primarily dependent upon the stability of the light source, the stability of the efficiency of the optical system, and the stability of the counting time intervals. The research presented here shows in detail how a low power tungsten bulb or hollow cathode tube can be used (2) J. R. Prescott and P. S. Takhar, IRE Trans. Nucl. Sei., NS-9, 36 (1962). (3) R. W. Engstrom and J. L. Weaver, Nucleonics, 17 (2), 70 (1959). VOL 40, NO. 1 1 , SEPTEMBER 1968
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