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 Unfortunately, substantial amounts of Li, Cr, and Mg were (19) H. A. Laitinen, unpublished data, 1952.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
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
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.
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
Acetic anhydride also converts dinitrogen pentoxide to acetyl nitrate (7,8) A c ~ O NzO, 2AcON02 (4)
+
Nitration of 2,4-dinitrobenzyl alcohol (7) in dilute acetic anhydride-acetic acid and of benzene in “pure” acetic anhydride (9) is second order in nitric acid. In 0.01M sulfuric acid, the nitration of benzene becomes first order in nitric acid (9). One study (9) postulates that in “pure” acetic anhydride (ca. 8M anhydride and ca. 1M acetic acid from the water in the nitric acid), the nitronium ion is the reactive intermediate. Another study (10) postulates that protonated acetyl nitrate is the reactive intermediate in the same solvent. A third study (7) postulates that in dilute anhydride-acetic acid, dinitrogen pentoxide is the nitrating species. It is claimed that the nitronium ion intermediate does not appear to be consistent with product stereochemistry ( I O ) and that the dinitrogen pentoxide intermediate does not permit a simple rationalization of the sulfuric acid catalysis (9). At 0-nitration conditions close to systems in this paper, Bonner ( 7 ) found two rate maxima. One occurred at 0.68M anhydride and the other at much higher anhydride concentrations, The maxima at 0.68M anhydride could be accounted for in terms of dinitrogen pentoxide as the reactive intermediate and Reactions 2 and 4. (Once the anhydride concentration is increased above 0.68M, the additional anhydride removes dinitrogen pentoxide according to Reaction 4.) The maxima at much higher anhydride concentrations was accounted for in terms of dinitrogen pentoxide and Reactions 2 and 3. The increase in anhydride concentration would increase the ionization of nitric acid and its subsequent reaction with acetyl nitrate (Reaction 3). Regardless of the mechanism, the primary reaction can be written as the formation of a nitrate ester. HO-NO2
+ RQH
-+ H20
+ RON02
(5)
We were interested in two questions: Could experimental conditions be controlled so that Reaction 5 would be stoichiometric, and could the stoichiometry be measured accurately under these conditions? It was obvious that undiluted anhydride-nitric acid could not be used to answer either question. Inspection of Reaction 5 suggested that titration with standard sodium hydroxide might be used to follow the progress of nitration by neutralization of unreacted nitric acid, A study was then made of various reagents and conditions that could be used to answer the questions. The results are described below. EXPERIMENTAL
Apparatus. A Perkin-Elmer Infracord infrared spectrophotometer with standard sodium chloride cells was used to obtain the infrared spectra of the nitrate ester of 1-dodecanol. Reagents. Nitric acid was colorless B & A C.P. reagent grade. Acetic anhydride, pyridine, acetonitrile, and other chemicals were reagent grade. Potassium acid phthalate was primary standard grade. Some of the alcohol samples were distilled and others were recrystallized before use. 0.42M NITRICACIDREAGENT.Eleven milliliters of acetic anhydride was pipetted into 10 ml of acetonitrile in a 100-ml volumetric flask. The mixture was cooled in an ice bath and kept in the bath during the slow dropwise addition of 2.65 ml of concentrated nitric acid. About 50 ml of acetonitrile was added and the flask was agitated in the bath for about 5 (9) M. A. Paul, J . Am. Clzem. SOC.,80,5329 (1958). (IO) F. G. Bordwell and E W. Gorbisch, Ibid.,82, 3588 (1960). e
ANALYTICAL CHEMISTRY
minutes. The reagent was allowed to warm to room temperature and was then diluted to the mark with acetonitrile. It was kept away from direct sunlight or was kept in the dark, and it was prepared fresh daily. This reagent is about 0.5M in acetic anhydride. Other reagents with different amounts of nitric acid and acetic anhydride were prepared similarly, 0.6N SODIUMHYDROXIDE.About 100 ml of water and 5400 ml of absolute methanol were added to 240 ml of saturated aqueous sodium hydroxide. This titrart was standardized against potassium acid phthalate. MIXED INDICATOR. Three parts of 0.1 neutralized thymol blue and one part of 0.1Z neutralized aqueous cresol red were mixed. Infrared Identification of Product. To ensure that nitration was occurring, the following experiment was performed. Two millimoles of 1-dodecanol were nitrated with 10 ml of the 0.42M nitric acid reagent for 20 minutes. Ten milliliters of hexane were then added and the two-phase solution was stirred overnight with a magnetic stirrer. The solution was then poured into a separatory funnel and the upper hexane phase was subjected to infrared analysis. A reagent blank without 1-dodecanol was treated similarly. Study of Experimental Variables. Experiments designed to show whether nitration or acetylation occurred and to test experimental variables were carried out in the same way. Exactly 10 ml of whatever nitrating reagent was used was pipetted into a glass-stoppered flask containing the alcohol sample. The mixture was allowed to stand at room temperature in the dark or away from direct sunlight for the appropriate time. Then 2.5 ml of water was added to prevent volatilization of nitric acid, and 7.5 ml of pyridine was added for base-catalyzed hydrolysis of the acetic anhydride. After 10 minutes at room temperature, the flask is titrated as in the procedure below for primary alcohols. A blank is run using the appropriate nitrating reagent alone. The calculation is as described below in the procedure. Procedure for the Determination of Primary Alcohols. Weigh accurately a sample containing 2 mmole of primary hydroxyl group into a 125-ml glass-stoppered flask. Pipet in exactly 10 ml of the 0.42M nitric acid reagent, and swirl the flask gently. Allow the mixture to stand 20 minutes at room temperature in the dark or away from direct sunlight. Then add 2.5 ml of water, swirl, and add 7.5 ml of pyridine. Allow the flask to stand 10 minutes at room temperature. Add mixed indicator to the flask and titrate with 0.6N sodium hydroxide, using a 50-ml buret, to a color change from yellow to violet. Run a reagent blank by pipetting exactly 10 ml of the 0.42M nitric acid reagent into glass-stoppered flask and allowing it to stand 20 minutes at room temperature in the same light as the samples. Then add 2.5 ml of water and 7.5 ml of pyridine. Titrate as dexribed above for the sample. Use the difference between the blank titration and the sample titration (V, - V,) to calculate the per cent purity of the primary alcohol.
x
RESULTS AND DISCUSSION
Nitration GS. Acetylation. Since nitric acid is a weak catalyst of acetylation (11) and since the uptake of acetic anhydride would be measured as nitration by acid-base titration, it had to be established that a nitrate ester rather than an acetate ester was the main product. A sample of 1dodecanol was treated with the 0.42M nitric acid reagent in acetonitrile, the products were extracted into hexane, and the infrared spectrum of the extract was taken (Figure 1). The reagent itself was also extracted and the spectrum of its ex(11) J. S.Fritz and G . H. Schenk, ANAL.CHm., 31, 1808 (1959).
I 3t
I
4
[
S
"
I
I
6
7
Figure 1. Infrared spectrum of hexane extract of the nitrating reagent plus 1-dodecanol, indicating the presence of l-nitratododecane
tract was also taken (Figure 2) to ensure that the spectrum in Figure 1 was not that of the reagent alone. Judging from the appearance of bands at 3.7, 5.8, and 10.75 p in both figures, it would appear that the dimer of acetic acid (from acetic anhydride and water in the nitric acid) was extracted in both cases. It is difficult to rule out the formation of a small amount of acetate ester since any acetate band at 5.7 p would be obscured by the 5.8 p acetic acid band. The NO2band at 7.9 p also tends to obscure the characteristic acetate band at 8 p . Figure 1 does show a small shoulder in this region that might arise from acetate. There is, however, no doubt that a nitrate ester is formed. This is particularly evident at 6.1 (NO, asymmetric stretch), 7.9 (NOz symmetric stretch), and 11.7 p (0-N stretching). Weaker bands at 13.2 (out of plane deformation) and 14.4 p (NOz bending) also confirm nitrate ester. All of these bands have been reported by previous workers (6, 12). It seems unlikely that nitric acid would be extracted into hexane although its intense absorption bands at 5.95 and 7.6 p (13) may be masked by bands at 6.1 and 7.9 p. However, the 13-p band (13) is absent in both Figure 1 and Figure 2. Chemical evidence also exists for nitration as opposed to acetylation. In a careful study of the nitration of 2,4-dinitrobenzyl alcohol with 0.32M nitric acid-0.39M acetic anhydride in acetic acid, Bonner (7) failed to find a detectable amount of ester. These conditions were not greatly different from those used in this study: 0.42M nitric acid-0.5M acetic anhydride in acetonitrile. The data in Table I also appear to support nitration. As the ratio of anhydride to nitric acid is decreased (at a constant ratio of anhydride to 1-dodecanol), the per cent reaction decreases. If nitric acid were merely acting as a catalyst of acetylation, it would be unreasonable to expect that halving the anhydride and alcohol concentrations would decrease the per cent reaction from 90 to 10 (Table I). In fact, it has been shown that decreasing the concentration of anhydride from 2 to 0.25M in ethyl acetate in the presence of an effective acid catalyst (14) does not significantly reduce the per cent acetylation. The catalysis of O.15M nitric acid of acetylation by 2 M acetic anhydride in ethyl acetate has also been shown to be very weak (11). If nitration were occurring under the conditions of Table I, a decrease in the concentration of anhydride would certainly decrease the rate (regardless of mechanism) by shifting the point of equilibrium in Equation 2 to the left, Effect of Acids. Since it has been shcwn that sulfuric acid accelerates nitration (9), the effect of different acids on the (12) J. F. Brown, J. Am. Chem. SOC.,77,6341 (1955). (13) R. A. Marcus and J. M. Fresco, J . Chem.Phys., 27,564 (1957). (14) G. H. Schenk and J. S. Fritz, ANAL.CHEM., 32,987 (1960).
i
I
1
I
8 9 IO II W A V E L E N GTW I M I C R O N S )
WAVELENGTH (MICRONS)
L
1
I
12
13
14
1
15
Figure 2. Infrared spectrum of hexane extract of the n i ~ ~ a t i n reagent alone, indicating absence of 1 - n i ~ ~ ~ o ~ o a e c ~ n e
Table I. Effect of Changing the Anhydride to H N 0 8 Ratio at Constant Ratio of Anhydride to Alcohol AczO, M a Reaction in 15 minutes 1-Dodecanol,A4 0.42 0.42 90 0.20 0.42 0.32 0.32 55 0.15 0.42 0.22 0.22 10 0.10 0.42 a The molarity of acetic anhydride remaining after reaction with the 13% water in the nitric acid. ~
~
~
~
Table 11. Effect of Solvent on Nitration of 1-Dodecanol
Nitric acid,
Initial ml, AczO per
M
100ml
0.4 0.52
10 10
0.8
12
0.48
12
% Reaction after 30 minutes EtOAc (Et0)8P0 25 11 9 47 38 ,.. 63 58 99 95 86
CH3CN
.
i
.
nitration of 1-pentanol was investigated, In the presence of sulfuric acid, a white precipitate was formed upon titration of the reaction mixture with standard sodium hydroxide. Addition of phosphoric acid gave no appreciable catalytic effect. Perchloric acid should be an effective catalyst because of its inherently greater acidity in nonaqueous solvents, but erratic results were obtained with it. After 30 minutes of nitration, the following results were obtained with the stated concentrations of perchloric acid: O.I4rCf, 73% ; 0.164iM, 17%; and 0.187M, 46%. Since it has been found that perchloric and sulfuric acids accelerate the hydrolysis of isopentyl nitrate (15),it appears that the perchloric acid acts the same way in these experiments. This must occur during the short interval between the addition of water and the addition of pyridine before titration, and the per cent reaction probably varies with the time of this interval. This made it necessary to obtain optimum stoichiometry with nitric acid alone, acting as catalyst itself. Solvent Effects. Nitration definitely depends on the type of solvent, Ethers such as dioxane, diethylene glycol dimethyl ether, and tetrahydrofuran are poor solvents for nitration. For example, the per cent reaction of 1-dodecanol in (15) T. 6.Bonner and D. E. Frizel, J . Chem. SOC.,1959, p. 3902. VOL. 39,
NO. 14, DECEMBER 1967
e
1797
Table 111. Variation of the Nitric Acid-Acetic Anhydride Ratio in Acetonitrile Solvent Nitric acid, Anhydride, Reaction Of 1-dodecanol after: M Ma 15 min 45 min 0.48 0.44
0.34 0.40 0.53 0.57 0.68 0.79
0.42 0.32 0.32 0.32
27.4 84.3 99.7 93.5 93.2 104
40.4 98.4 99.3 98.9 97.4 104
0 Molarity of acetic anhydride after reaction with water in the nitric acid.
Table IV. Determination of Primary Alcohols Purity Av dev, found Alcohols (in 20 minutes) (3 to 12 results) Cyclohexyl carbinol 99.9 0.3 1-Decanol 99.9 0.6 1-Dodecanol 100.0 0.7 Methanol 101* 1 2.4 I-Qctadecanol 98.7 1.5 1-Qctanol 100.0 0.8 1-Pentanol 94.7 1.8 &Butanol 100.0 0.1
z
z
Table V. Other Nitration Results Alcohol or mixture Reaction Time, minutes SECONDARYALCOHOLS Benzoin 52-57.6 20 2-Butanol 100 20 Cyclohexanol 89-95 30-60 Diethylcarbinol 72-81 20 2,6-Dimethyl4-heptanol 91-97 20 OTHERS
2-Methyl-2propanol 1-Qctanol 0.8 mmole 1-octanal 1-Octanol 0 . 5 mmole 1-butanal I-Octanol 0.7 mmole 1-butanal
+
76-96
20-90
+
104-106
20
+
99.6
30
116
30
diethylene glycol dimethyl ether was only 18% using the 0.48M nitric acid reagent listed in Table 11. Triethyl phosphate was adequate (Table 11) but appeared to retard nitration compared to ethyl acetate and acetonitrile. With ethyl acetate, a white precipitate forms and redissolves during the hydrolysis with water and pyridine. The same thing happened with acetone and dioxane as solvents. Acetonitrile was therefore chosen as solvent for further work. Nitric Acid-Anhydride Ratio. The ratio of nitric acid to acetic anhydride was critical in acetonitrile, as shown in Table 111. As the ratio approached 1 : I , the rate of nitration increased. The optimum rate and stoichiometry occurred at a 1.2:l ratio of anhydride to nitric acid. The rate then decreased until the anhydride-acid ratio exceed 2.5 :1. This behavior can be rationalized in terms of Equations 2,3, and 4, just as Bonner (7) rationalized his results. The optimum rate 1798
0
ANALYTICAL CHEMISTRY
can also be rationalized in terms of acetyl nitrate (9) or protonated acetyl nitrate (10) as the intermediates. The 104% nitration may be the result of nitration of acetic anhydride itself (16). Other Variables. The concentration of the sample, and therefore the sample size was critical. Variation of the sample size from 1 to 3 mmoles of I-dodecanol with the 0.42Mnitric acid reagent indicated that best results were obtained with a 2-mmole sample. After 20 minutes reaction, nitration of either a 1- or a 3-mmole sample of 1-dodecanol gave results ranging from 93 to 97%. Since the reagent contains only 4.2 mmoles of nitric acid, the use of a 3-mmole sample probably decreases the concentration of nitric acid below that necessary for rapid reaction. Using a 2-mmole sample, the reaction time was less crucial. The per cent reaction of 1-dodecanol increased from 82% at 5 minutes to 91% at 10 minutes to 99.7% at 15 minutes. From 15 to 60 minutes, the per cent reaction did not alter significantly. Other primary alcohols reacted at similar rates but most secondary alcohols reacted more slowly and incompletely. A time of 20 minutes was chosen for later investigation of the stoichiometry of reactions of primary alcohols as a class with nitric acid. Because it was desirable to hydrolyze the acetic anhydride in the reagent to acetic acid before titration, the ratio of pyridine to water for the hydrolysis was important. Pyridine is necessary because it catalyzes the rapid hydrolysis of acetic anhydride (17). Water was added first after nitration was complete to hydrate the anhydrous nitric acid and prevent its volatilization when pyridine is added. At ratios of pyridine to water of less than 3 :1, 74 to 95% reaction of 1-dodecanol was obtained. At ratios of 4 :1 or more, the per cent nitration of 1-dodecanol was usually 97 to 98% but some lower results were obtained. A 3 :1 ratio of pyridine to water gave consistent results. Variation of the hydrolysis time after nitration of I-dodecanol gave 98.3% reaction after 5 minutes' hydrolysis but results which were consistently in the 99 to 100% region after 10 to 30 minutes' hydrolysis. Determination of Primary Alcohols. In general, the nitration of long chain primary alcohols exhibited a 1 : I stoichiometry. The stoichiometry could be measured accurately enough by sodium hydroxide titration to use nitration for the quantitative determination of primary alcoholic hydroxyl groups, The ease of nitration varied with the individual alcohols: longer chain primary alcohols were not as volatile as alcohols such as methanol and I-pentanol, and 2 mmoles of the former could be weighed and handled with less evaporation. However, as the chain lengthens, the alcohol becomes less soluble in the nitrating reagent. For example, l-octadecanol did not readily dissolve at room temperature but dissolved upon mild heating. The corresponding blank was similarly heated to correct for any side reactions. Table IV contains the analytical data for the determination of primary hydroxyl groups. Nitration of Secondary and Tertiary Alcohols. Unhindered secondary alcohols such as 2-butanol also reacted quantitatively in 20 minutes (Table IV). More hindered secondary alcohols reacted more slowly and did not appear to react stoichiometrically (Table V). Cyclohexanol in particular reacted in a peculiar fashion, being 89% nitrated after 30 minutes, but only 95% nitrated after 60 minutes and 97.5% (16) F. H. Cohen and J. P. Wibaut, Rec. Trau. Chim., 54, 409 (1935). (17) A. R. Butler and V. Gold, J. Chem. SOC.1961, p. 4362.
nitrated after 3 hours, It is possible that slow oxidation to adipic acid occurred in the later stages of the reaction. The only tertiary alcohol investigated, 2-methyl-2-propanol (tert-butyl alcohol), behaved somewhat like cyclohexanol. It was also more difficult to reproduce a particular result after a given time. Ninety minutes appeared to be the optimum nitration time; as the reaction time was increased, the titration showed a gradual decrease in per cent nitration. It is possible that decomposition to a nitroalkene occurred as has been observed with 98% nitric acid and 2-methyl-2-propanol at 0 to 5' C (18). When bromine was added to a nitration (18) C. A. Michael and G. H. Carlson, J . Am. Chem. Soc., 57,1268
(1935).
mixture after hydrolysis, decolorization of the bromine occurred, indicating the presence of some functional group such as a double bond. Interferences. Aldehydes, which are frequently impurities in alcohols, interfere in the nitration of primary alcohols by giving high results. A slight evolution of yellow fumes and formation of a yellow solution was observed upon addition of the reagent to octanal. By itself, octanal reacted about 25% with the nitrating reagent, Small amounts of aldehydes can be tolerated in the determination of primary alcohols, but amounts larger than 0.5 mmole cause high results (Table V).
for review September
1967* Accepted September
13, 1967.
Solubility and Complex Formation Equilibria of SiIver ChIoride in Propylene Carbonate James N. Butler Tyco Laboratories, Inc., Waltham, Mass. 02154 The equilibria of silver chloride i n propylene carbonate solutions containing excess chloride have been studied potentiometrically in a constant ionic medium (0.1N tetraethylammonium perchlorate) at 25" C. Equilibrium constants were fitted by a nonlinear least-squares pit-mapping technique. Only mononuclear complexes ASCI, were found. The overall formation constants for these complexes are: log p1 = 15.15 i: 0.15, log pz = 20.865 i: 0.015, log p3 = 23.39 i: 0.06 (for n = 1, 2, 3) and the solubility product of silver chloride is log K,, = 19.87 i: 0.02 (errors are standard deviations). The solubility of AgCl in excess chloride is approximately equal to the concentration of excess chloride, and the predominant soluble silver species is ASCI2-. The equivalence point of the titration of CI- with Ag+ occurs at [Ag+] = 1O-9*1, essentially independent of the concentration of reagents or supporting electrolyte. The thermodynamic equilibrium constants were estimated using the Debye-Huckel theory and neglecting ion-pair formation, and alternatively by assuming reasonable values for the ion-pairing constants.
PROPnENE CARBONATE (4-METHYLDIOXOLONE-2) has recently been used as a solvent for electrochemical studies relating to high-energy batteries ( I , 2) and free radical species (3). Its high dielectric constant (4) (64.4 at 25 C) permits ionic salts to dissolve at concentrations of several moles per liter, and its aprotic character permits electrode reactions with strongly reducing species such as radical anions or lithium metal to be carried out with little decomposition of the supporting electrolyte. One promising reference electrode system in this solvent is silver/silver chloride, and considerable effort has been expended in attempts to construct batteries using silver/silver (1) R. J. Jasinski, "High Energy Batteries," Plenum Press, New York, 1967. (2) R. J. Jasinski, J . Electroanal. Chem., 15, 89-91 (1967). (3) R. F. Nelson and R. N. Adams, J. Electroanal. Chem., 13, 184 (1967). (4) W. S. Harris, thesis, University of California, 1958; U. S. At. Energy Comm. Rept. UCRL-8381.
chloride cathodes with lithium anodes (5-7). This paper re ports studies of the silver/silver ion electrode, and of the solu bility and complex formation equilibria of silver chloride. EXPERIMENTAL
Propylene carbonate (Matheson, Coleman and Bell) was purified by distillation in a Podbielniak vacuum-jacketed column of approximately 50 theoretical plates, operated under total reflux for several hours before any distillate was withdrawn. The reflux ratio was then changed to 1O:l. The first 200 ml of distillate were discarded. From an initial charge of 3500 ml, approximately 2700 ml of purified solvent were collected, The distillation proceeded at a pressure of 1 mm Hg and a stillhead temperature of 78 to 80" G. Analysis for water and volatile organic impurities was made by gas chromatography (8), on a column of Porapak Q, using thermal conductivity detection and helium as a carrier gas. The only detectable impurity was 15 to 20 ppm of water. Organic impurities were less than 4 ppm. Silver perchlorate (Chemical Procurement) containing 0.5 % water, tetraethyl ammonium chloride (Eastman) containing 1.5 water, and tetraethyl ammonium perchlorate (Eastman) containing less than 0.1 water, were dried over anhydrous CaS04 in a desiccator before solutions were prepared but were not otherwise purified. Soltldions were analyzed for chloride ion by potentiometric titration with aqueous silver nitrate, and for silver ion by potentiometric titration with a nonaqueous chloride solution of known concentration. The cell consisted of two electrode compartments connected by a salt bridge. Coarse glass frits separated the two ends of the salt bridge from the electrode compartments, and one of the electrode compartments contained a Teflon-coated (5) J. E. Chilton, Jr., Tech. Documentary Rept. ASD-TDR-62-1
(April 1962), AD 277 171. (6) H. F. Bauman, J. E. Chilton, W. J. Conner, and G. M. Cook, Tech. Doc. RTD-TDR-63-4083, (October 1963) AD 425 876. (7) J. E. Chilton, Jr., W. J. Conner, G. M. Cook, and R. W. Holsinger, Tech. Rept. AFAPL-TR-64-147 (February 1965), AD 612 189. (8) R. J. Jasinski and S. Kirkland, ANAL.CHEM., 39, 1663 (1967). VOL. 39, NO. 14, DECEMBER 1967
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