Electroanalytical Study of the Reduction of Chromate in Molten Lithium Chloride-Potassium Chloride Eutectic H. A. Laitinen and R. D. Bankertl Unioersity of Illinois, East Chemistry Building, Urbana, Ill. Chronopotentiometry of chromate in molten LiCI-KCI eutectic revealed a single, irreversible, diffusioncontrolled, three-electron reduction step. The diffusion coefficient of chromate ion at 450° is estimated to be 1.7 x 10-6 cm2 sec-1. Current-reversal chronopotentiometry revealed abnormally short reverse transition times, which decreased markedly with rising temperature, but increased somewhat in the presence of added oxide ion. It is concluded that the primary reduction product i s the ion Cr04+, which decomposes to yield Cr03-3 and O-2. At a sufficiently high surface Concentration of CrO4-5, solid Li,Cr04 is deposited. In the presence of MgClz or CaCI,, the reduction potential is shifted 0.8 and 0.6 volt, respectively, in the positive direction. Extraordinarily stable deposits are produced. No reverse electrochemical oxidation of these deposits occurs.
THEREIS substantial disagreement in the literature as to the mechanism and even the stoichiometry of the cathodic reduction of chromate in molten salts. Black and DeVries ( I ) found two voltammetric waves for chromate in LiC1-KC1 eutectic at 408" using platinum microelectrodes, and concluded that the two waves corresponded to successive reduction steps of Cr(V1) to Cr(II1) and Cr(II), respectively. On the other hand, Ferguson ( 2 ) and Bhatia (3) observed only a single reduction wave at 450". Controlled potential coulometry indicated that three faradays of electricity were consumed per mole of chromate. Analysis of the black, insoluble electrode deposit was inconclusive because of entrainment of the molten salt solvent. Treatment of the film with water yielded a solution containing Cr04-2, Lif, K+, el- and OH-, and hydrated CrzOs as a precipitate. A material balance indicated that the film contained an excess of Li+ and Cr(II1) in the approximate ratio 5 LizO+Cr203.It was not clearly established, however, whether this material was a single compound or a mixture. Chromate has been used as a depolarizer in thermal battery systems using molten LiC1-KCL as the solvent. Selis and McGinnis ( 4 ) studied the galvanic cell
I
I
Mg LiC1-KC1 - K2Cr04 Ni. From plateaus on potential-time curves observed upon slow discharge, they postulated the formation of various oxides of nickel formed by its oxidation by chromates. Similar systems using Ca, Mg, or Ni as anodes were studied by Selis and McGinnis (51, and by Selis, Wondowski, and Justus (6) 1
Present address, American Cyanamid Co., Stamford, Conn.
-~
(1) E. D. Black and T. DeVries, ANAL.CHEM., 27, 906 (1955). (2) W. S. Ferguson, 15th monthly rept., Contract No. DAI-49-186502-ORD(P)-187, University of Illinois, 1955. (3) B. B. Bhatia, Ph.D. Thesis, University of Illinois, 1959. (4) S. M. Selis and L. P. McGinnis, J. Electrochem. Soc., 106,
900 (1959). (5) S. M. Selis and L. P. McGinnis, Ibid., 108, 191 (1961). (6) S. M. Selis, J. P. Wondowski, and R. F. Justus, Ibid., 111, 6 (1964). 1790 *
ANALYTICAL CHEMISTRY
using silver as a cathode. Although it was concluded that silver is oxidized by chromate, the reduction process of chromate was not characterized, nor was the direct electrochemical reduction of chromate considered. The present investigation was undertaken to elucidate the reaction mechanism with the aid of current-reversal chronopotentiometry, and to establish the composition of the primary reduction product formed during the first moments of electrolysis. For this purpose, direct emission spectrochemical analysis of the first few micrograms of solid deposit was undertaken. EXPERIMENTAL The furnace assembly and temperature control system, along with the melt container and associated equipment, have been described previously (7, 8). Current-voltage curves were obtained with a Sargent Model XV polarograph (E. H. Sargent and Co., Chicago, Ill.). A Sargent Model IV coulometric current source was used to generate the Pt(I1) reference electrode. The constant current source used for chronopotentiometry and currentreversal chronopotentiometry has been described (9). Automatic current reversal was effected with a thyratron-relay circuit, the trigger pulse being obtained from the Delayed Gate output of a Tektronix 535-S2 oscilloscope (Tektronix, Inc., Portland, Ore.). The switching time was from 3 to 5 mseconds. The ratio of the forward to reverse current was variable over a 12-fold range. Potential-time curves were recorded with a Tektronix 53542 oscilloscope, or in some cases, with a Houston Instruments Model HR97 X-Y recorder. The oscilloscope traces were photographed using a Tektronix type C-12 camera with a Polaroid back. All emission spectrographic analyses were performed using an N.S.L. Spec Power High Voltage Spark Source [National Science Laboratories, Inc., Cleveland, Ohio) in conjunction with a Bausch and Lomb Littrow Spectrograph (Bausch and Lomb Optical Co., Rochester, N. Y.) and an A.R.L. Universal Arc-Spark Stand (Applied Research Laboratories, Los Angeles, Calif.). Spectra in the wavelength region of 2500 to 3500 A were recorded on Kodak Spectrum Analysis 1 plates, while those in the region 3550 to 7500 A were recorded on Kodak 1-N plates. Line intensity data were obtained with an N.S.L. Spec Reader Model TM-102 densitometer. The copper spark emission spectrographic method (IO) was used to analyze the electrode deposits. A detailed account of the sampling procedure and preparation of calibration curves is given elsewhere ( 1 1 ) . Solvent. The solvent system consisted of a eutectic mixture of 41 mole % KC1 and 59 mole % Lie1 and was obtained from Anderson Physics Laboratories, Inc., Champaign, Ill. (7) H. A. Laitinen and K. R. Lucas, J . Electroanal. Chem., 12, 553 (1966). (8) H. A. Laitinen and D. R. Rhodes, J. Electrochem. SOC.,109, 413 (1962). (9) S. Deron, Ph.D. Thesis, University Of Illinois, 1964. (10) N. H. Nachtrieb, M. Fred, and F. S. Tomkins, J . Opt. SOC. Am., 37,279 (1947). (11) R. D. Bankert, Ph.D. Thesis, University of Illinois, 1966.
I
1
1
2.5
2.c
1.:
ul
50
', ' 1.c
o o w
-025
-050
- 075
-10
APPLIED POTENTIAL VOLTS
- 1 25
(vs P i b ) REFERENCE)
Figure 1. Current-voltage curve for chromate reduction A = 2.4 X 10-3 an2,t Curve A. C = 5.16 X Curve B. C = 2.51 X
=
0.5
450"
mole cm-3, residual MgClz mole em+, no MgClz
0.2 S E C . 4 C
The eutectic was purified using the procedure described previously (12) except that Mg was not added to displace heavy metal cations. Electrodes. Platinum microelectrodes were prepared from 26 B and S gauge platinum wire as described earlier (8). Indicator electrodes for chronopotentiometric measurements were made by spot welding platinum foil (0.05 mm) to 26 B and S gauge Pt wire. The counter electrode was either a Pt flag or a 118-inchdiameter graphite rod. The Pt(II)/Pt reference electrode was used in this study. Details on the generation and use of this reference electrode in LiC1-KC1 have been described previously (12, 13). Chemicals. Reagent grade chemicals were used in all cases. Those chemicals containing water of hydration were vacuum dried at 110" C. before being added to the melt. RESULTS AND DISCUSSION
Preliminary voltammetric experiments (Figure 1) showed strikingly different behavior depending upon whether or not the melt had been treated with magnesium to displace the residual heavy metal impurities. Addition of an excess of anhydrous MgClz caused the first wave to become fully developed and the second one to disappear. Calcium ion behaved qualitatively the same. Evidently, the double wave of Black and DeVries ( 1 ) can be ascribed to the fact that in this early work no special pretreatment had been used for the elimination of water from the melt. As a consequence, the borosilicate glass vessels underwent serious etching and presumably enough calcium ion dissolved to produce the first wave. The investigation was therefore carried out in two parts, first in the absence of added divalent ions, and then in the presence of added calcium or magnesium chloride. Reduction of Chromate in Absence of Divalent Metals. The voltammetric curves were not highly reproducible, as might be expected from the fact that a solid deposit is formed. As with solid metal deposition (14) the current was found to increase with time upon continuous polarization at - 1.2 volts. A black deposit could be observed microscopically on the platinum microelectrode surface. (12) H. A. Laitinen, R. S. Tischer, and D. K. Roe, J . Electrochem Soc., 107, 546 (1960). (13) H. A. Laitinen and C . H. Liu, J . Am. Chem. Soc., 80, 1015 (1958). (14) H. A. Laitinen, C . H. Lin, and W. S. Ferguson, ANAL.CHEM., 30, 1266 (1958).
TI ME,
SEC.
Figure 2. Chronopotentiogram for chromate reduction in absence of CaClz or MgC12 C = 8.48 X t = 450"
mole cm-a, ZO = 1.32 mA
Chronopotentiometry allowed this difficulty to be circumvented, by proper choice of concentration and current density. A typical cathodic chronopotentiogram is shown in Figure 2. Diffusion control was demonstrated by the constancy of Z T ~ as / ~a function of I. Because of limitations in the useful range of transition times to avoid the disturbing effects of convection and residual impurities, only a limited range of current densities could be studied for a given chromate conmole/ centration. For a solution containing 8.48 X cm3 K2Cr04,and for six current densities ranging from 9 to 14 mA/cm2, at 450" the value of Z T ~ was ' ~ 9.06 =t0.16 mA sec"* cm-2. From the Sand Equation, the value of n 2 D is correspondingly 1.55 i. 0.2 X cm2 sec-'. Using n = 3 from the controlled potential coulometry studies of Ferguson and Bhatia, we calculate D = 1.73 f 0.05 cmz sec-1 for chromate ion at 450". This value appears reasonable in comparison, for example, with previous estimates of 1.7 X cm2 sec-l and 2.02 x 10-6 cmz sec-l for the diffusion coefficient of Cd(II), present as CdC14-2, in the same melt at 450". Plots of E US. log ( 7 1 ' 2 - t 1 / 2 ) / t 1 / 2 were nonlinear, but as shown in Figure 3, the plot of E 1;s. log ( ~ 1 1 ~t 1 I 2 ) was linear. Equating the slope of the linear plot to 2.3 RTIaNaF, we calculate aN, = 0.45, suggesting a totally irreversible charge transfer step in which the transfer of the first electron is ratecontrolling with a transfer coefficient of 0.45. From the intercept, the cathodic charge transfer rate constant k , is calculated to be 9.6 x 10-7 cm sec-I, referred to the zero of the potential scale [ l M platinum (11) I platinum]. At E = - 1.18 volts, k , = 4.73 x 10-3 cm sec-l. The chronopotentiometric results also are of interest in connection with the question of the stability of chromate ion in LiCI-KCL melts. Although previous results in this laboratory ( 2 , 3 ) , as those of Smith and Boston ( 1 4 , indicate that the chromate is a stable species, Delarue (16) has stated that ~~
(15) G. P. Smith and C. R. Boston, Ann. N . Y.Acad. Sci., 79, 930 (1960). (16) G. Delarue, J . Electroanal. Chem., 1, 285 (1959). VOL. 39, NO. 14, DECEMBER 1967
0
1791
I
I
'
"t
I
03
om
9
os
0,s
OZC .LOG
c25
7'2
"XI
-7'2
,
on0
"35
045
050
050
060
lEECll2,
Figure 3. Potential-time analysis mole cm-8, l o
1.13 X t = 450"
C
=
=
Figure 4. Effect of oxide ion on reverse transition time
11.0 mA
= 1.79 X mole cm-a, l o = 14.0 mA cm-2 Curve A. No added oxide, l o ' = 3.0 mA Curve B. 6.27 X PO-5 moIe LLO added, IO' = 4.0 mA
C
chromate dissociates to yield C r 0 3and O-*, giving an orangecolored solution in which C r 0 3 is the electroactive species, The chronopotentiometric results showed no indication of complications due to an equilibrium prior to the charge transfer step. A slow equilibrium would have been expected to lead to an abnormally low apparent diffusion coefficient,which would vary with added oxide ion concentration. The transition time, however, like the voltammetric wave height (3), was unaffected by the addition of Li20 to the system. If a rapid equilibrium existed, the reduction potential should have responded to added oxide, but again, there was no effect. In addition, there was no indication of an anodic wave due to oxide unless LizOhad been purposely added. Current-reversal chronopotentiometry showed an anodic transition upon reversal of the current before reaching the cathodic transition (Figure 4). Although the back-transition time increased somewhat upon the addition of excess LizO, it was always shorter than that expected for reoxidation of a diffusing reduction product. The ratio of 10'7'/Io'tf, where laand Io' are the forward and reverse current densities, t f is the forward electrolysis time and 7' is the back transition time, was found to decrease with rising temperature, as shown in Table I. The theoretical relationship is given by
which yields, for the experimental IO = 5.8, IO' = 3.2 mA cm-2, a theoretical ratio Io'r/Idf = 0.392. The abnormally short reverse transition times, coupled with the effect of added oxide suggest a decomposition of the primary reduction product to yield oxide ion. The simplest such scheme is CrO4-2
+ 3e-
t/
kl
Cr03-3
Cr04-6 7'
k-
+
0 - 2
1
Table I. Dependence of l o ' ~ p / I oon t f Temperature mole/cmaK2Cr04at 450" C 2.01 X l o' / l o "C t j sec T ' sec Io'+'/Iat, 0.553 4% 1.96 0.64 0.177 500 1.96 0.43 0.121 550 1.96 0.38 0.107 600 1.95 0.22 0.062
1792
a
ANALYTICAL CHEMISTRY
(2)
18
I
I
4
1
tfO.:SECL
Figure 5. Dependence of r' on t f Curve A . Theoretical dependence for ZO'/ZO = 0.5 Curve B. Experimental for C = 1.49 X 10 -5 mole ZO = 8.16 mA ~ m - Io' ~ , = 4.08 mA C I I I - ~ Curve C. Theoretical dependence for 10'/10= 1.0 Curve D. Experimentalfor C = 2.13 X IO-: mole c m P I' = IO' = 6.0 mA c m P
Quantitative interpretation was complicated by the formation of an insoluble film whenever the forward coulombs exceeded a few millicoulombs per cm2. Figure 5 shows the dependence of the reverse transition time upon the forward electrolysis time at 450" for two different current density ratios, in comparison with those calculated from Equation l. The reverse transition time, although always shorter than theoretical, shows a tendency to increase more rapidly than theoretically expected whenever the total forward charge exceeded 4 to 5 mC/cm2. This increase is undoubtedly associated with the finite solubility rate of a solid film formed only upon appreciable cathodic electrolysis. By keeping the forward coulombs I d f constant and below that for film formation, it would be expected that the reverse transition time would deviate more and more from the theoretical with increasing forward electrolysis time. This is indeed the case, as shown in Figure 6. In an attempt at a quantitative kinetic analysis, reverse transition times were measured in the presence of excess added
Table 11. Data for Rate Analysis [O-l]= 6.27 X mole/cm3; [CrOa-2]= 1.79 X mole/cm3 IO Io ’ (Tt)-I/z mA/ mA/ cm2 cms t f (sec) 7’(sec) 1 t f / ~ ’ (sec-1’2)
14-
12-
M
-+
10-
16 16
08-
16
.t
16 16 10 10 10
08-
04-
-
4 4 4 4 4 12 12 12
1.94 0.98 0.76 0.55 1.47 1.92 1.43 0.91
2.53 1.16 0.88 0.595 1.82 0.31 0.26 0.15
1.33 1.36 1.37 1.39 1.34 2.495 2.57 2.65
0.63 0.93 1.07 1.30 0.74 1.65 1.98 2.51
0,z
0
02
I 06
d4
I
08
$ 2
I O
14
I
I
I
le
I8
2 0
2 2
tpor ISECI
Figure 6 . Dependence of
7’
on tf at constant and low lot,
Curve A. Theoretical dependence for Io’/Io = 0.5 Curve B. Experimental for C = 1.49 X lom5mole cm-2, l o ’ / l o
=
Table 111. Analysis of Electrode Deposit Cr Coulombs found, Cathodic Anodic Sample Pg 0.43 1 34.4 X 10-3 None 2 3 4 5*8 gYb 7 8 9
0.5
oxide ion to approach a pseudo first order reverse chemical reactio?. The data are given in Table 11. Dracka (17) has derived the relationship between forward electrolysis time and reverse transition time for the case of a reversible first order chemical reaction following an electrochemical reaction. Using our notation this is given by
where k = kb
+
kf sum of forward and reverse chemical rate constants
Ka K = - equilibrium constant of chemical reaction Kf For kT‘ > 6, erf -I I
dk7 s 1 and Equation 3 becomes 4; lo’ tr
/In’
’)
+ 2K
dc 10
(4)
For the five experiments in Table I1 at ZO’/ZO= 0.25, a plot of 4 1 tf/7‘ cs. 1/47; was linear, with an intercept very close to the theoretical value of 1.25, and a slope of 0.096 secllz, corresponding to K d k = 2.31. Using Equation 3, and the data for ZO’/ZO= 1.2, we estimate k = 1.28, 1.87, and 3.04 sec-1 for 7 ’ = 0.37, 0.27, and 0.15 second, respectively. This trend, although consistent in direction with an increase in surface oxide concentration and a smaller apparent decomposition rate constant with increasing total forward coulombs lot,, is more plausibly associated with the beginning of the formation of an insoluble film. The value k = 3.04 sec-’ is probably the best estimate, because it corresponds to the least amount of forward electrolysis. Taking K = 2.31 and k = 3.04, we estimate k, = 1.3 and kb = 1.7 sec-l. Although we recognize that these estimates are valid only to a factor of two or three, they do appear of the right order of magnitude to be consistent with other observations.
+
dz
(17) 0. Drarka, Collection Czech. Chem. Commun., 25,338 (1960).
20.6 X 10-3 53.4 x 10-3 34.6 X 39.8 X 80.7 X 36.0 x 10-3 38.0 X 22.2 x 10-3
None
None None None None
12.89 x 10-3 12.99 X 7.55 x 10-3
0.37 0.20 0.40 0.24 0.70 0.145 0.195 0.126
Li/Cr 5.6 4.8 4.6 4.5 5.1 2.9 10.9 8.0 9.1
* Allowed to stand on open circuit. Q
for 5 minutes for 10 minutes
Analysis of Surface Films. Because of the previous difficulties due to entrainment of melt constituents, it was desired to produce the deposit under chronopotentiometric conditions and to prepare it for analysis in a manner that would not necessitate its treatment with water. Accordingly, electrolysis was carried out using a 0.5-cm2 platinum flag as cathode for about 2 seconds at 16 to 20 mA/ cm2. The cathode was washed in a clean melt, and the excess solvent was evaporated off by heating for 24 hours under vacuum at 450”. The deposit was dissolved in 1 ml of hot, concentrated HCl, concentrated to ca. 0.5 ml and evaporated, along with 20 hg of Mn as internal standard, onto 6.25-mm Ag rods, which had been machined lightly to remove any oxide coating before being used as spectrographic electrodes. Electrode blanks, receiving the same treatment but without electrolysis, showed no measurable amount of lithium, chromium, or potassium. Results of several analyses are given in Table 111. Samples 1 to 4, which were prepared with forward electrolysis only, and without standing on open circuit, showed an average Li:Cr ratio of 4.9, indicating a film composition LijCrOd. Samples 5 and 6, which were allowed to stand on open circuit indicate negligible decomposition in 5 minutes, but an appreciably lower Li:Cr ratio after 10 minutes of standing. Samples 7, 8, and 9 indicated that when reverse electrolysis was carried out the Li:Cr ratio increased substantially, and that an appreciable amount of chromium was lost from the deposit. It should be noted that in all the samples, the efficiency of film formation was low; only 10 to 1 5 % of the cathodic current went to formation of the solid product. The amount of anodic current passed in preparing samples 7, 8, and 9 was several times that required to oxidize the chromium found in the deposit in comparable experiments without reVOL. 39, NO. 14, DECEMBER 1967
* 1793
3.01
i
25T
i i
c 51
1
t
+05SEC
Figure 8. Effect of CaClz on chronopotentiograms of chromate
C = 1.12 X 10-6 mole ~ m - ~ Curve A. No CaCh mole cm.-8 CaClz Curve B. 2.26 X mole cm-3 CaClz Curve C. 1.08 x
verse current (samples 4, 5 , and 2). Therefore most of the anodic current was consumed in reoxidation of chromium(II1) present in the solution rather than in the film. The reason for the abnormally high Li:Cr ratio in the film remaining after reoxidation is not clear, but it should be noted that the analytical method used would not reveal what ani0n-e.g.) chloride or oxide-might be retained with excess lithium ion upon partial removal of chromium from the film by reoxidation. An x-ray powder pattern of the deposit was obtained. D spacings, along with approximate relative densities, are listed in Table IV. The sample was prepared by controlled potential electrolysis of 273 mg of K2Cr04,followed by heating at 450’ under vacuum for 20 hours to remove excess melt. No lines for LiCl or K2Cr04 were found in the diffraction pattern. Lines attributable to KC1 were observed but are not included in Table IV. The diffraction data do not agree with any known chromium compound listed in the ASTM files. Since no d spacings attributable to LizO were found, we may conclude that the x-ray data are consistent with the postulation of a new compound of formula LijCrOq rather than mixtures of LizOwith Cr203or lower lithium chromites.
Table IV.
X-Ray Data for KrCr04Reduction Products
Electrolvsis. no CaClz d, A
5.40
IIIo, red 100 80 53 47 46 40 35 33 33 33 33 20 20 20
2.11
15
3.48
6
2.72 2.04 2.79 3.88 2.47 4.11 1.99
4.76 1.93 1.45 3.65 2.38 1.97
1794
0
Electrolysis, with CaCh Intensity d, A order 4.81
1 3 4 2
4.76 3.03 2.48 2.30 2.06 1.92 1.88
1 3 4
5
1.46
5
None 2.48 2.36 2.06 1.89 1.86 1.47 1.39
ANALYTICAL CHEMISTRY
Galvanic cell (5) Intensity d, A order
1.39
2
Reduction of Chromate in Presence of Magnesium or Calcium Ions. Chronopotentiometry fully confirmed the
results qualitatively anticipated from voltammetry. Addition of small quantities of MgCl, (Figure 7) or CaC1, (Figure 8) caused the appearance of an early reduction step, which became fully developed at sufficiently high concentrations of MgClz or CaC12, with quarter wave potentials of -0.2 and - 0.4 volt, respectively, rs. the Pt(II)/Pt electrode. The new step was shown to be diffusion-controlled by the ’ ~ varying l o . In the presence of MgC12, constancy of 1 0 7 ~with Z O T ~ was ’ ~ 5.07 =k 0.13 mA secl’Zcm-2 for C = 5.63 X 10-6 mole/cm3 corresponding to D = (1.25 i 0.06) X 10-6 cm2 sec-I for n = 3, whereas in the presence of C a G , Z O T ~was ’~ 2.52 i 0.09 cm2 sec-1 for C = 2.63 X low6mole/cm3 corresponding to D = (1.39 + 0.04) X cm2 sec-l. Considering the inherent inaccuracies in measurement of electrode areas for the platinum flag electrodes, especially because film formation began with the first moments of electrolysis, these results are consistent with expectations for a three-electron, difiusion-controlled reduction. Upon reversal of current, the potential dropped immediately to that of the Pt(II)/Pt reference electrode, showing that the film cannot be reoxidized electrochemically. It was not possible to obtain a reliable value of n by controlled potential electrolysis, because as soon as the cathode became covered with an appreciable amount of deposit the current decreased, evidently because the film is somewhat resistive. Analysis of the Deposits. The insoluble reduction products formed in the presence of MgClz or CaClz proved to be extraordinarily stable, Both products resisted all forms of aqueous acid attack, even by aqua regia or boiling concentrated HC104. This behavior is remarkable in view of the fact that a standard procedure for quantitative oxidation of Cr(II1) in solution is boiling with concentrated HCIOe (18). The magnesium product formed such a tightly adherent film on platinum that it could not be scraped off without removing appreciable quantities of platinum, enough to interfere with an x-ray powder pattern, The calcium product, however, formed a loosely adherent film that could be scraped off. (18) J . Jordan, J. Meier, E. Billingham, and A. Pendergrast, Nature, 187, 318 (1960).
Its x-ray powder diffraction data are listed in Table IV, and compared with similar data for a compound prepared in 1952 (19) by short-circuiting the following galvanic cell:
I
I
Ni LiCl-KC1-KzCr04 Ca
(5)
Several products seemed to be produced, but the most stable was a black deposit that resisted all types of acid attack. It could be dissolved only by alkaline fusion. Chemical analysis of the compound showed 20.7 % LizO, 21.O % CaO, and 55.7z CrzO3, in reasonably close agreement to the composition of a compound 2Li20.CaO .Crz03,or Li4Ca(Cr0&, which would contain 22.4% LizO, 20.9% CaO, and 56.7x Crz03. The comparison of x-ray data in Table IV is suggestive that the two materials are identical. An attempt was made to analyze the first few micrograms of magnesium or calcium products spectrochemically by producing the deposit on 6-mm-diameter platinum rods which could be mounted directly in a spectrograph and analyzed by the copper spark emission method. Blank experiments were run by the same procedure, omitting the electrolysis step. Unfortunately, substantial amounts of Li, Cr, and Mg were (19) H. A. Laitinen, unpublished data, 1952.
invariably found in blank experiments whether the electrode was rinsed with distilled water or treated by heating in vacuum at 450". In either case, potassium was present only in trace amounts whether electrolysis was performed or not. Similarly, in the presence of CaCL, blank experiments showed high and erratic amounts of Li, Cr, and Ca, but no potassium. Although attempts were made to apply blank corrections to arrive at the compositions of the deposits ( I I ) , the results were variable and inconclusive. A systematic investigation is now under way to prepare larger quantities of these extraordinarily stable new materials, and to determine their compositions under a variety of conditions of concentration, temperature, and current density. ACKNQ WLEDGMENT
We are grateful to J. P. Walters for helpful discussions concerning the emission spectrographic work. RECEIVED for review August 7, 1967. Accepted September 18, 1967. This work was supported by the Army Research Office, Durham, N. C.
A Study of the Quantitative Nitration of Alcoholic Hydroxyl Grou George H. Schenk and Milagros Santiago Department of Chemistry, Wayne State University, Detroit, Mich.
48202
A study of the nitration of alcohols with various nitric acid-acetic anhydride reagents has shown that O-nitration of primary alcohols is quantitative at room temperature in 20 minutes in acetonitrile solvent. Most secondary alcohols and 2-methyl-2-propanol (tertbutyl alcohol) nitrate more slowly and incompletely. Infrared and chemical studies indicate that nitration rather than acetylation occurs in the nitric acid-acetic anhydride reagent. The extent of nitration can be found by sodium hydroxide titration of the unreacted nitric acid and a reagent blank. A study has also been made of the effects of other acids, solvents, and ratio of nitric acid to acetic anhydride.
These colorimetric methods succeed because they are reproducible, but they are not necessarily stoichiometric. Reagents less reactive than fuming nitric acid or sulfuric acidnitric acid are needed to restrict the many possible reactions to one primary reaction which is thermodynamically and kinetically favored. One such reagent is nitromethane, which can be used for stoichiometric nitration of tyrosine at pH 8 in aqueous solution ( 5 ) . A more promising reagent for waterinsoluble compounds is the acetic anhydride-nitric acid reagent (6)recently used to prepare nitrate esters of hydroxyl compounds. This reagent, acetyl nitrate, consisted of two drops of 70% nitric acid mixed with 0.3 ml of acetic anhydride at 0" C. It was then mixed with about 50 rng of hydroxyl compound at room temperature. The reagent will be anhydrous because of the nitric acid-catalyzed reaction of water with acetic anhydride
NITRATION of ORGANIC COMPOUNDS has been mainly utilized for synthetic work because nitration is not selective and is difficult to control. However, colorimetric methods using fuming nitric acid or sulfuric acid-nitric acid mixtures have been used for the detection of aromatic hydrocarbons like pyrene ( 1 ) and for the determination of various hydrocarbons ( 2 , 3) and bound styrene (4). Benzene has also been determined in the presence of toluene and xylene by oxidizing the latter compounds with chromic acid, nitrating the benzene, and condensing the resulting m-dinitrobenzene with butanone (4). (1) E. Sawicki and T. W. Stanley, Chemist-Analyst,49,77 (1960). (2) E. Berl and W. Koerber, IND.ENG.CHEM., ANAL.ED., 12, 175 (1940). (3) E. Berl and R. Raub, Ibid., 12, 177 (1940). (4) A. Muller, Rec.. Foc. Quim. Uiiiu. Nod. Mayor Sun Marcos, 7, 5 (1955); C.A. 50, 16563a (1956).
+
A c ~ O H20 -(H-)+
2HOAc
(1)
After this step, both acetyl nitrate and dinitrogen pentoxide (7, 8) form in equilibrium amounts
+ HONOz AcONOz + H'N03Ace0
$
+ HOAC e NzOs + HOAC
AcONOz
(2) (3)
( 5 ) J. F. Riordan, M. Sokolovsky. and B. L. Vallee, J . Am. Chem. SOC.,88,4104 (1966).
(6) D. C . Malins, J. C . Wekell, and C . R. Houle, ANAL.CHEM., 36, 658 (1964). (7) T.G . Bonner, J. Clzem. SOC.1959, p. 3908. (8) V. Gold, E. P. Hughes, and C. K. Ingold, Ibid., 1950, p. 2467. VOL. 39, NO. 14, DESEMBER 1967
e
9795
