anodic shift. At the present time, one cannot say whether or not surface activity is a necessary requirement for the observed shift. Calcium ions, which are not surface-active, have been shown t o produce similar but larger shifts (8) (see also Figure 5 ) , while tetramethylammonium ions, which are less surfaceactive than DTAC, produce a smaller shift than calcium ( 4 ) . The large shift in the half wave for p-nitrophenol shown b y calcium ion in Figure 5 n-as almost entirely eliminated by the presence of DTAC. This confirms the suggestion (8) that prior interaction of calcium with nitrobenzene does not take place; instead, ion-pairing n-ith a reduction intermediate must occur. The effects of surface-active agents
on polarographic waves of nitro compounds can be useful for analytical purposes. For certain compounds, the shifting of waves to more cathodic potentials in acidic media or opposite shifts in basic media may yield better limiting currents. This should allow greater accuracy in determining wave heights. I n addition, reducible mixtures which would normally have overlapping waves may be resolved since the potential range required for the nitro reduction should be shortened. LITERATURE CITED
S. L., Aussen, B. S.,Rec. trur. cham. 73. 455 11954). (2) Dratovsky, hI., Ebert. M.>Chem. lzsty 48, 498 (1954). ( 3 ) Heyrovsky, J., Sorni, F., Foreijt, J., (1) Bonting,
Collectton Czechoslov. Chem. Comniuns. 12, 11 (1947).
(4) Hummelstedt, L. E. I., Rogers, L. B., J . Electrochem. SOC.106, 248 (1959). (5) Nolthoff, I. M., Lingane, J. J., "Polarography," 2nd Ed., Vol. 2 , Chap. 28, Interscience, New York, 1952. (6) Korshunov, I. -I.,Kirillova, A. S., J . Gen. Chem., I'.S,S.R. 18, 785 (1948). ( 7 ) Page, J. E., Smith, J. W., Waller, J. G., J . Phys. & Colloid Chem. 53, 545 11949). (8j Pietrzyk, D. J., Breese, R. F., Rogers, L. R., unpublished data. (9) Schniid, R. IT..Reilley, C. K., J . Am. Chem. Soc. 80, 2088 (1958). (10) Tomamushi, R., Yamanaka, T., Bull. Chem. SOC.J a p a n 2 8 , 673 (1955). RECEIVED for review October 23, 1961. Accepted April 25, 1962. Work supported in part by the U.S. -4tomic Energy Commission under Contract hT(30-1)-905.
Nitromethane as a Solvent for Solid Electrode Voltammetry J. D. VOORHIES and E. J. SCHURDAK Organic Chemicals Division, American Cyanamid Co., Bound Brook,
b Nitromethane has been used as a solvent for constant current voltammetry with the platinum electrode. Cathodic reductions and anodic oxidations of both organic and inorganic species are described. The performance of the platinum electrode in dry nitromethane solutions o f tetraalkyl ammonium salts and anhydrous M g (C104)2 has been studied. The maximum electromotive range i s obtained with 0.1M Mg(C104)?,the anodic and cathodic backgrounds occurring a t +2.4 and -2.4 volts vs. normal hydrogen electrode respectively, a t a current density o f 0.3 ma. cm.-* The effects o f water and oxidative pretreatments on the platinum electrode are also discussed. The anodic oxidation of iodide to iodine in nitromethane occurs in two discrete steps. The chronopotentiometric phenomena associated with these two steps are analogous to those reported in the literature for the polarography of the iodine, iodide redox system in acetonitrile. Some useful information on the reactivity and apparent reduction potentials of several positive chlorine compounds dissolved in nitromethane, 0.1 M in tributylethyl ammonium nitrate, i s presented.
I
s A ~ T U D I of the reactivity of organic compounds containing the nitrogen-cliloyine bond, nitromethane was a useful solvent for electrochemical
N. J .
measurements. The so-called positive chlorine compounds are not stable in many solvents because of general solvolysis, chlorination, or oxidation reactions(5,li). Common ionizing solvents such as water, alcohols, acetone, and acetonitrile react rapidly with some or all compounds containing the S-C1 bond with the result that meaningful electrochemical measurements cannot be made. -4lthough nitromethane is not completely inert to S-chloro compounds, several chloramides XT-hich react rapidly with u-at'er, acetone, and other solvents !\-ere dissolved in nit'roniethanesupporting electrolyte solutions and subjected to cathodic chronopotentionietry (4, 1 5 ) a t the platinum electrode with good results. Sitromethane vias then examined as a general solvent, for electrochemical osidation-reduction ytudies. The useful properties of this solvent for such studies are a wide polarization range coupled n i t h a low concentrat'ion of labi!e protons. Other properties contributing to its utility for elect'rochemistry are high dielectric constant [ea. 40 a t 20" C. (6'b)], a reasonnble lioiling point (101' C.)) and low viscosity [0.63 centipoise a t 25" C.(o'c)]. -1 niajor disadvantage of nitromethane as a solvent for voltammetry is t'he poor solubility of most' common supporting elect'rolytes. I n this study, t'he tetraalkyl ammonium salts, anhydrous magnesium and lithium perchlorates, and certain Lewis acid halides
such as ferric chloride and zinc chloride were the only salts soluble a t the 0.1111 level. Selson and I m m o t o ( 7 , 10) have cited the solubility of other metal perchlorates in a n electrochemical study of metal ion reductions in nitromethane. A more extensive search for soluble salts would be appropriate before further studies with this solvent are attempted. The general electroniotive range of the platinum electrode in nitromethanesalt solutions is limited on the cathodic side by the presence of small amounts of water TT hich may be asqociated n i t h the nitromethane solvent molecules (16) This limitation is very much dependent on the supporting electrolyte used, and a large number of organic compounds TT ith reducible functions can be reduced cathodically 11ell before the bachgr ound reduction. Oxidations of organic compounds such as hydroquinones anti phenols lire also possible. Inorganic redox ystenis studied TT ere fenic reduction and iodide oxidation. EXPERIMENTAL
2;
Reagents. Most organic compounds described were laboratory samples of high purity prepared a t t h e Bound Brook Laboratories, =Imerican Cyanamid Co.. Bound Brook,
K. J. All salts used were of reagent grade except for the tetraalkyl ainrnonium salts (Eastman). Tetrachloroglycouril was obtained from Diamond Alkali Co , VOL. 34, NO. 8, JULY 1962
939
and spectroquality acetonitrile was obtained from hhtheson, Coleman & Bell. Nitromethane (Matheson, Coleman & Bell), b.p. 1O0-Zo C., contained about 0.05% water. Apparatus and Procedure T h e voltammetric technique used throughout this investigation was voltammetry a t constant current or chronopotentiometry. A constant current in t h e range of 0.05 t o 1.0 ma. was drawn through two platinum electrodes, and the e.m.f. between one of t h e platinum electrodes a n d a reference electrode mas measured as a function of electrolysis time on a strip chart recorder (Sargent, RIulti Range). T h e constant current supply used was a commercial transistorized unit (Model 151B, Quan Tech Labs., Boonton, X. J.), The platinum indicating-working electrode was a platinum wire 2 mm. in diameter and 1.3 sq. em. in exposed surface area tightly fitted into Teflon. The reference electrode was a methanol-aqueous, normal calomel electrode containing about 20 volume per cent of methanol (RINCE) and having a reduction potential of +10 to +20 mv. us. aqueous saturated calomel electrode in dilute nitromethane -salt solutions. All subsequent e.m.f. values are reported with reference to this electrode. The electrolytic solution used throughout most of this study was nitromethane, 0.05 to 0.08V in tetramethyl ammonium chloride (ThIAC) . These solutions were essentially dry except for traces of water in the solvent, absorbed from the atmosphere and contained in the salt crystals used to prepare the solution (50.05% HzO). The water was determined b y Karl Fischer titration. The effect of water on the reduction of 1,4-benzoquinone is described in Results and Discussion.
I n the study of the N-chloro compounds, the chloride ion from TMAC reacted with the positive chlorine of the N-chloro compound to give chlorine. Thus, the reduction potentials for the first reduction step in a series of such compounds dissolved in nitromethaneTRIAC were almost identical and equal t o that of chlorine. Therefore, the electrolytic solution used for the chronopotentiometry of &'-chloro compounds was nitromethane, O . l J 1 in tributylethyl ammonium nitrate prepared by reacting a nitromethane solution of the iodide with a stoichiometric amount of silver nitrate. The silver iodide foimed was filtered off, and the resulting solution \vas partially dried with molecular sieve 51-1 (Fisher Scientific Co.) to less than 0.03y0water. Test solutions contained about 50 mg. of active chlorine in 100 ml. of solution based on a concurrent analysis of the active chlorine content of the compound in question. This was done so that relative potential measurements for a series of compounds would be made under the same conditions of concentration of reducible function and applied current density. The solution temperature throughout this study was 26" A 3" C. RESULTS AND DISCUSSION
Platinum Electrode Behavior. T h e useful cathodic range of the platinum electrode i n nitromethane-salt solutions is primarily dependent on the amount of water present (0.01 to 2%). The reduction of d r y nitromethane about 0.05M i n T M A C (0.03% HzO) at a precathodized platinum electrode occurs between -0.90 and -1.0 volt a t a current density of 0.3 ma. em.-* The cathodic
Table I. Chronopotentiometric Oxidations and Reductions in Nitromethane-TMAC i"T"2
Compound Ferric chloride Reduction Hydroquinone oxidation 1,4Beneoquinone ( 5 0 . 1 % "20
)
Reduction 1,4-Benzoquinone (2% H,O) Reduction 1,4Benzoquinone (4% (v.) X / l HCIadded) Reduction
s-s
+
I-
Reduction
940
0
ANALYTICAL CHEMISTRY
Concn. Current ( M X Density 103) (pa. Cm.-*) 5.06 373 5.06 298 0.25 38.4 5.28 720
C
-0.08
Amp. Sec.1'2 1. Moles Cm.2 0.333 0.343 (27' C.) 0.312
+0.74
0.54
T
(Sec.) 20.5 33.9 4.2 15.7
El/, +0.02 +0.03
-0.46 298 13.3 12.4 -0.37 531 531 8.45 -0.38 Poorly defined reduction step Interference of background reduction
0.19 0.19 0.29
5.17
377
15.9
-0.35
0.29
4.0
298
14.1
-0.35
z 0.28
5.74 9.84 532
FERRIC REDUCTION
\
r=LOSsec
\I EMF ( V o l t s
VI
MNCE)
5.352 sec IOGIDE OXIDATION
k, 7, =
I
7060
157 sec
0
Figure 1. Chronopotentiograms in nitromethane Top 5.06 X 10-3M FeC13, i,, = 373 po./rq. cm. CH~NO?, 0 . 0 8TMAC, ~ = 270 c . Bottom 5.62 X 10-3M TBEAI, io = 308 pa./rq. cm. CH,NO?, 0.1 M LiCIOI, T = 26' C.
r
depolarization shifts, in the anodic direction upon addition of water and approaches - 0.4 volt in nitromethane containing 2% water. The electrode process occurring a t the background depolarization in dry nitromethane solutions may be the reduction of a molecular complex betn-een water and nitromethane. Evidence to support this view is the fact that the cathodic background depolarization can be extended to about -2.6 volts a t 0.3 ma. cni.-2 in nitromethane, 0.1111 in lIg(C1OJ2. In this case, the magnesium ion may displace the nitromethane from its complex with water. -4 similar behavior n a s observed for 0.1111 LiC104 in nitromethane. Unfortunately, this effect n u q not discovered until the completion of this study, so very few experiments were carried out with 0.1V lIg(C10412 or LiC104 supporting electrolyte The anodic limit of the platinum electrode in nitromethane, 0.05 to 0.0831 in TJIAC. is $0.92 rolt a t 0.3 The source of this anodic background is probably the oxidation of chloride ion. I n nitromethane, 0.1M in I\Ig(C104)z, the anodic background occurs a t about $2.2 volts a t 0.3 ma. em.+ After a short depolarization time a t this potential, the cell resistance increase? abruptly such that the I R drop in the cell exceeds the voltage output of the current supply. a result, the cell current decreaaeq, and the platinum electrode polarizes slightly in the cathodic direction. Apparently, a highly resistive film is produced b y the anode reaction a t this potential. The effect of two oxidative pretreat-
Table II.
Iodide Oxidation
C iOT1/2"
Solvent (Electrolyte) 1. 5.62 X 10-3M Tributylethyl ammonium iodide A. CH8NOz (0. lizl LiCIOa) B. cH3so2 (0.1M LiC104) c. CH&T\TOz ( 0 , l M LiC104) D. CH3?i02 (0.1M LiC104) E. CH3NOz (0.03Zf TWIAC) F. CHsNOz (O.03M TRIAC)
Other Conditions
i0
E1 14
598.5
+0.205 $0.515
499 6 sec. precathodization-immediate current reversal 1-2.76 X 10-8M I2
+o. 18
+0.50
499
1-0.205 +0.49
347
+ O . 145 +0.43 +0.35 +0.53 +O.57(t) (t 30 min.) (t 105 min.) +0.39 +0.78
730 +2.76 X 10-3M IZ single oxidation step
Amp. Sec.1/21.
(pa. Cm. - 2 )
1133
t = time of first run
G. HzO
732
iodide H. CHsCS ( 0 ,Ifif SaC104.H20)
228
(0,051lf Tlf AC) 2. 2.16 X 10-3M Sodium
T
(Sec.)
3.84 ( l ) c 6.00 (2) 5.82 (1) 9.54 (2) 10.8 (1) 8.28 (2)
( RIoles Cm.2 ) 0.334 ( 7 1 0.347 ( T I
+
3 . 0 (1) 38.4 (2) 7 . 2 (1) 14.0 (2) 9.0 9.6
0.597 ( 0.605 0.625
+
10.2
0.645
6.6(1) 15.0 (2)
0.313
+ +
+
T ~
D X lobb (Cm.2 Sec.-l)
7*)1/2
1.5:3 ( l e - )
T Z ) ' ~ ~
1.65 ( l e - )
7*)lIZ
1.22 (2e-) 1.26 (2e-)
1.62 ( l e - )
$0.21 8.22 (1) +0.54 11.76 (2) 0.471 ( T + ~ T ~ ) ~ I 3.05 ~ (le-) T = 25' & 1" C. in all experiments The chronopotentiometric constant is calculated from 7 1 7 2 when this sum represents a diffusion controlled transition time for a total process involving one-electron transfer per iodide. b D,the diffusion coefficient, is calculated from equation ( A ) using the chronopotentiometric constant cited in the previous column. c The numbers in parentheses refer to the sequence of stepwise oxidation. 0
+
ments of the platinum electrode, namely tion for the current-time-concentration the dichromate-sulfuric acid cleaning relationship in chronopotentiometry (4), solution and a concentrated nitromethane solution of tetrachloroglycouril, one of the positive chlorine compounds The oxidation and reduction steps studied, was examined. The chromic were in all cases well defined (Figure 1) acid treatment resulted in a poorly except for benzoquinone reductions defined reduction step having a tangent potential of -0.80 volt. (The tangent in nitromethane solutions containing water. The chronopotentiometric conpotential is measured a t the intersection stant for ferric reduction provides an of straight lines drawn through the approximate reference value for an elecinitial charging and depolarization retrode process involving one electron gions of the potential-time curve.) transfer. The diffusion coefficient for This reduction becomes more facile upon addition of n a t m to the nitroferric ion calculated from the chronomethane solution. After the tetrapotentiometric constant for semi-infinite chloroglycouril treatment. no evidence linear diffusion and the electrode area of oxidatioii of the platinum electrode (1.3 sq. cm.) is 1.6 X 10-5 em.* see.-' was obser\ ed upoii cathodic electrolysis. a t 27' C. Although several errors The pretreatment of the platinum inherent in the experimental technique electrode throughout this study conmay result in a high value for this diffusisted of an occnsional chromic acid sion coefficient, it is nevertheless much wash follo~vedby precuthodization for larger than the value, 0.57 x 10-5 cathodic runs and acetone and n-ater cm.2 see.-', cited for 0.3M Fe (S03)3 washes in between niost runs. The in water (sa). The possible sources of error are cylindrical diffusion (3, 12) effect of precathodization on the anodic oxidation of iodide is discussed in a and electrode surface roughness n hich subsequent section. increases the effectii e electrode area General Reductions and Oxidations with respect to the measured geometric in Nitromethane. Several representaarea (2'). Hoivever, these errors do not tive, electroactive compounds were account for a significant difference besubjected t o anodic or cathodic, tween the diffusion coefficients in the constant current electrolysis in nitrotwo s o l ~ e n t s . The large diffusion comethane-TMAC solution. The efficient in nitromethane is apparently chronopotentiometric results are a manifestation of a very weak interacshown in Table I. I n Table I, t h e tion of the diffusing species with the , chronopotentiometric constant, i o ~ ~ / ~ / C solvent molecules. is a combination of measurable or conThe EIMvalue for ferric reduction in trolled variables from the basic equaTable I is much less anodic than the
reduction potential of the ferric, ferrous couple in noncomplexing aqueous solui similar El,, was observed for tion. ; the reduction of FeC13 in nitroruethane, 0.1114 in LiC104. I n light of the large diffusion coefficient of ferric in nitromethane, it seems unlikely that strong solvation of ferric is responsible for this shift in reduction potential. The experimental evidence is not sufficient to provide an adequate explanation of this phenomenon. The oxidation-reduction study of 1,4benzoquinone, 1,4-hydroquinonewas carried out primarily to test the behavior of a redox system involving protons in a solvent containing a low level of labile protons. Kawzonek et al. (17) report that hydroquinone does not oxidize in neutral acetonitrile solution because i t does not ionize, ionization being a prerequisite to electron transfer. Hen-ever, in nitromethane, 11)-droquinone appears to oxidize by n two-electron change to the quinone (See chronopotentiometric constant in Table I.). The chronopotentiometric ronstant for the cathodic reduction of 1,4-benzoquinone in both n e t and dry nitromethane is indicative of incomplete reduction to hydroquinone. An explanation for this fact is that protonation of the cathodically inactive semiquinone anion formed in the one-electron reduction step is very slow, such that no appreciable concentration of the neutral, reducible semiquinone radical is produced a t the electrode surface during the chronopotentiogram. As VOL. 34, NO. 8, JULY 1962
941
water and/or dilute HCl are added, the rate of protonation of the semiquinone anion formed a t the electrode is increased such that some complete reduction to hydroquinonp can occur. The reduction of a dithiolium salt (8) (Table I, bottom) appears to proceed through a one-electron step to a neutral dithiolium radical. Iodide Oxidation. T h e anodic oxidation of iodide to iodine in nitromethane solution occurs in t F o steps (Figure 1, Table 11). A similar double step mas observed by Popov and Geske ( 1 3 ) for the iodine, iodide redox system in acetonitrile and is explained in terms of a stable triiodide intermediate, 61- += 2134e- (2/3e- per I-) (1) 213- += 312 2e- (1/3e- per I-) (2)
+ +
The electrochemical evidence for this mechanism is a 2 : 1 ratio of id(l)to id(21, millicoulometry of the total process and the disappearance of wave (1) upon addition of sufficient iodine to convert the iodide to triiodide. The chronopotentiometric data in Table I1 are indicath e of a similar mechanism in nitromethane. The ratio of T~ to T $ for tributylethyl ammonium iodide (TBE-11) in nitromethane, 0.1M in LiCIOl (Table 11, -4 and I?), is 0.61 to 0.64 while the evpected ratio for a stepwise process involving 2i and elec-
Table 111.
trons per iodide ion is 4 to 5 or 0.8 ( 4 ) . The corresponding ratio for ?;a1 in acetonitrile a t low T (no convective error) is 0.70 (Table 11, H). Also, the addition of about mole of I2 per mole of I - reduces the transition time of the first step (Table 11, D). The resulting transition time accounts for slightly less iodide than expected on the basis of complete converpion of the iodine added to triiodide. ilnother similarity of the iodine, iodide redox couple in acetonitrile and nitromethane is the slow increase in concentration of iodide in solutions containing iodine (Table 11, F ) . I n one experiment with 6.62 X 10-3X Iand 2.76 X 10-2A1112,the concentration of iodide doubled in tlvo days. This is very likely the result of iodination of a n enolic iaci) form of these two solvents with the simultaneous release of HI. Several investigators ha1 e discussed the possible role of adsorption of iodide and iodine, specific double layer effects, and the condition of oxidation of the electrode in the iodine, iodide redox process ( f , 9 ,f f , f S ) . Electrode surface phenomena also play a role in the anodic oxidation of iodide in nitromethane. Short precathodization of the platinum electrode with immediate current reversal to the anodic run results in an increase in T~ and a decrease in 7 2 with respect to previous and subsequent
runs a t the same current density (Table 11, C). The enhancement of the first step may be caused by Oxidation of reduced specie' formed during precathodizstion or of the platinum itself. The decrease in the second step cannot be explaind. The oxidation of I- in nitromethane containing chloride (TX.4'2) is described in Table 11, E and F . The chronopotentiometric constant from the total transition time of the tm-o steps, E , or the single step. F , and the coniparable data from the nork of Popov and Geske are indicatile of an over-all tn o elwtron oxidation evpressed by
I-
+ C1-
-+
IC1
+ 2e-
(3)
In experiment E , the first step n.ns very sensitive to the condition of the platinum n ith respect to precathodization, prewashing, and addition of I?. The T~ cited in Table 11, E represent? the largest and best defined first step. In some cases, step (1) is almost eliminated, but the chronopotentiometric constant based on ( T ~ T ~ ) * is independent of T ~ . K h e n 2 mole of I2 per mole of I - is added (Table 11, F ) , the first step is completely eliminated A comparison of step (1) involving the oxidation of I- in 0.1Jf LiCIOl and 0.05V T L I C .upporting electrolytes seems to indicate a difference in electrode surface phenomena in addi-
+
Cathodic Reduction of Some Positive Chlorine Compounds in Nitromethane iorl ~
Compound" Trichloroisocyanuric acid (41% active C1) Hexachloromelamine (52%) active Cli Chlorine (42 mg. 5O/ml.) ( t = time of soh. preparation) Tetrachloroglycouril (44 % active C1)
Current Density (pa. Cm.-2) 719
298
(See.) 20.1
719
1153
719 719 298
E tb +O.GY
$0.64
Moles Cm.0.38 (30" C.)
+O.76 +0,65 +o, 71
719
298
c
( q p . Sec.1'291.)
T
Very poorly defined reduction steps +O.G5 + 0 . 3 (1)c 63. -4.2 ca.
2.0 10.3
10.24
+ O . 28
- 0 . 1 (2) +O.lSd
0 . 4 8 i t + 1 houri 0.20 ( t 2 . r5 hours)
+
0.24 0.24
0
H-~-C-H
1
1
7.8 +0.78 + o . z (1)c 0.03 38.4 +0.11 +0.06 (2)d 0.15 147 12.9 298 +o.os ( 2 ) 10.2 -0.21 -0.28 0.26 719 1,3-Dichlorourea -0.11 -0.17 0.21 298 34.8 (437, active C1) a Test solutions contained about 25 mg. of active chlorine in 50 ml. of solution based on the analysis cited. * Et is the tangent potential a t which the first evidence of depolarization due to K-Cl reduction or some other reduction process is observed. The numbers in parentheses refer to the sequence of reduction steps. These Ella values were measured from fused double waves for both glycouril compounds. Diphenyltetrachloro glycouril (31y0 active C1)
942
ANALYTICAL CHEMISTRY
tion to the effect of chloride on the electrode process. For instance, the inhibition of T I upon addition of 12 is greater in 0.0531 TLIXC than in 0.1M LiC104. It is possible t h a t the specific adsorption of TRIAC results in a blocking of the electrode surface to I-. This would explain the general inhibition of T~ in T X i C and the shift in the of step (1) from +0.2 in 0.1M LiC104 t o +0.35 in 0.05X T X A C . Cathodic Chronopotentiometry of N-Chloro Compounds. The results
of cathodic chronopotentiometry of chlorine. four chloramide compounds, and hevachloromelamine are shown in Table 111. T h e purpose of this work was to examine the reactivity of t h e S-C1 bond in a series of similar compounds in ternis of different potentials of cathodic reduction. In this EtudyI precautions were taken to exclude all halide ions and maintain a low level of water ( 5 0.03y0). Thus, hydrolysis or disproportionation reactions of the S-chloro compound to give chlorine or. in the absence of chloride, hypochlorous acid n ere prevented or made insignificant. The reduction steps observed for most of the X’-chloro ronipounds in Table 111 n ere poorly defined probably because of a complicated, s l o ~c,!ectrode process occurring a t a low leiel of available protons which nould be involved in a two-electron reduction of a n X-Cl function. T n o possible steps in the general electrode process for S-C1 reduction are --T--Cl
+ e-
-f
=N,
+ C1-
(4)
Although the chronopotentiometric constants for the compounds described in Table I11 indicate electron change values in the 1 to 3 range, in some cases, further reduction of the test compound was observed at more cathodic potentials. However, these further reduction steps were so poorly defined in most cases t h a t no meaningful measurements could be obtained. Also, because of the poor definition of the first cathodic waves, the tangent potential (Et’)values are the most characteristic potentials for comparison purposes. Trichloroisocyanuric acid shon ed the best n-are shape and the most reproducible potential and io+?C values. On the basis of the chronopotmtiometric constants shown in Table I, the value of 0.58 for trichloroisocyanuric acid represents a two-electron reduction. Both chlorine and 1,3-dichlorourea, a n extremely unstable compound, showed reduction steps a t - 0.1 to -0.2 volt which may be accounted for b y chlorination of the solvent. However, cathodic tests on model compounds such as I-chloro-1-nitroethane and 1,ldichloro-1-nitroethane shon ed no evidence of reduction. Both tetrachloroglycouril compounds h a r e apparent reduction potentials ( E t ) very much loner than trichloroisocyanuric acid or hexachloronielamine. Both compounds show a fused double n a v e with a rather low chronopotentiometric constant as the major reduction step. The diphenyl substituted compound also shon-s a small reduction step which may be caused by a small amount of chlorine or other impurity in the test compound. These low reduction potentials of the chlorinated glycourils may be an indication of greater stability
of the K-Cl bond in these compounds with respect to trichloroisocyanuric acid and hexachloromelamine. CITED
. CHEM.33.
1123
Halogenation,” Chap. S and 9, Butterworths, London, 19,59. (6) International Critical Tables, McGrar-Hill, Xew Turk, 1926. a. T’ol. 5, p. 65. b. Vol. 6, p. S3 c. Yo1 7, p. 213. ( 7 ) .Inamoto, R. T., Private Coinmunication and Paper S o . 41 Division of Bnalytical Chemistry, 140th Meeting, ACS, Chicago. September 1961. (8) Klingsberg, E., J . d n i . Cheni. SOC. 83, 2934 (1961). (Second paper in press.) (9) Lorenz, IT., Muhlberg, H., Z. Elektrochem. 59, 736 (1953); %. PhUszk. Chena. (S.F.) 17,129 (1958) (10) Selson, I. IT., In amoto, R. T., A s . 4 ~CHEW . 33, 1795 (1961). (11) Newson, J. D., Riddiford, A . C., J . Electrochem. SOC.108,699 (1961). (12) Kicholson, bl. AI., J . -4111 Chem. SOC. 76,2539 (1954). (13) Popov, A. I., Geske, D. H., Ibzd., 80, 1340 (1958). (14) Rapoport, L., Smolin, E. >I., “sTriazines and Derivatives,” pp. 330333, Interscience, Yew York, 1959. (15) Reilley, C. S , Everett, G. IT., Johns, R. H., A \ A I . CHEV. 27, 483 (1955). (16) Streuli, C. A , , American Cyanamid Co., Stamford, Conn., private communication, 1961. (17) Wawzonek, S., Berkey, R., Blaha, E. IT., Runner, hl. E , J . Eltactrochem. SOC.103, 456 (1956). RECEIIEDfor review Januarv 16, 1962. iiccepted May 10, 1962. Metropolitsn Regional Meeting, ACS, KeTv York, January 1962.
Unusual Matrix Effects in Fluorescent X-Ray Spectrometry L. S. BIRKS U. S. Naval Research laborafory, Washingfon 25, D. C.
D. L. HARRIS American Chrome Co., Nye, Monf.
b Certain systems of elements may b e postulated in which the x-ray intensity from some element, A, will decrease rather than increase as the concentration of A increases. The requirement for such a situation i s that as the concentration of element A increases, the concentration of another element, B, which has a high x-ray absorption coefficient for radiation from element A, increases a t an even faster rate. Iron and chromium in the mineral
system chromite-olivine form such a pair of elements and the predicted phenomenon i s observed; i.e., FeKa intensity decreases steadily as the iron content increases from 8 to 17%.
I
fluorescent x-ray spectrometry, x-ray intensity is not linear with concentration ( 1 ) . Matrix absorption and enhancement are the causes for nonlinearity and are well understood. N
It has usually been assumed that a n increase in the Concentration of a given element would result in an increase rather than a decrease in ite x-ray intensity; that is, i t has been assumed that calibration curvee n-ould ah-ays have a positive slope. Recently, a practical mineral system was found in n hich this is not true. As the iron content of the system increases, the F e K a intensity decreases. The conditions for such a negative slope are rather restrictive as VOL. 34, NO.
a,
JULY 1962
943
