Titrations in Nonaqueous Solvents C . A. Streuli, Americcrn Cyanamid Co., Stamford, Conn.
A
in the understanding of acid-base phenomena in nonaqueous solvents has been marked in the past tv,o years b.,. a significant increase in papers dealing with the more fundamental aspects of the problem. There have been publications on scidbase behavior in acetonitrile (28, 78, 80, 1 IO), ethylenediamine ( I % ) , tert-butyl alcohol (W), and dimethylsulfo\ide (81), together with a number of articles on conductivity in a v u i e t y of solvents. These investigations demonstrate again that a number of factors must be considered when discussing acid-base behavior in any solvent. Following the established order (147) the present review covers papers published between Oct2ber 1961 and October 1963. The above-mentioned topics together with papers dealing with new or improved titrants, electrodes, solvents, ndicators, redox titrations, Lewis acics, and methods and procedures are included. Kot all papers published in t lis time interval are reviewed, of cour:e, some through oversight, others because it was felt that they did not add appreciably to the present body of knowli.dge. N ADVA?;CF.
conjugation in acetonitrile. This is defined as conjugation of an anion, Ai-,with a hydrogen bond donor, HR. In a solvent which is not hydrogen bonded and has a dielectric constant of 40 or less, the degree of dissociation of the salt B H + - i - is greatly increased by addition of HR. The effect can be illustrated by titration of sulfuric acid with pyridine in acetonitrile in the presence of resorcinol or 3,5-dinitrobenzoic acid with triethylamine, again in the presence of resorcinol. Coetzee et al. have outlined methods for the purification of acetonitrile (26) and Coetzee and Padmanabhan have determined the autoprotolysis constant of acetonitrile to be 3 X (28). Measurements were made on picric acid-tetraethylammonium picrate 1,3-diphenylguanidine-diphenyland guanidinium perchlorate buffers. Muney and Coetzee (110) also measured the degree of ionization of sis organic bases in acetonitrile and showed thisproperty for all six bases to be approsimately 4 powers of 10 smaller in acetonitrile than in water. The equivalent conductivity of a base in this solvent is highly concentration-dependent. For diphenylguanidine the predominating equilibrium reaction with the nitrile is:
FUNDAMENTAL ACID. BASE STUDIES
hcid-base equilibria in acetonitrile have been investigatrd by Kolthoff, Bruckenstein, and Chantooni (78) and Kolthoff and Chantooni (80). They found that, inaceton trile, perchloric acid is completely dissociated but other acids are weak. Sulfuric, nitric, hydrochloric, and hydrobromic acids dissociate according to the equation
Values for the stability constant, E S well as values for the simple disFociation constant, KRA. The order of acic strength in this solvent was found to be HCI04>HBr> H,SO,>HKOa, HCl, Hl'i, where H P i is picric acid. The relative relationships of these acids to one another and the relation between dissoriation of the acids in wetonitrile 2nd acetic acid must be roncerned with 3olvat'ion power, dielectric3 strength, a n 6 basicity of the solvent as well as formation and dissociation of ion pairs. Kolthoff and Chantooni (79) have also discussed the phenomenon of hetero-
KARA, were calculated
at concentrations of the guanidine greater than 5 X 10-2M. At lower concentrations, however, the predominating equilibrium is: R
+ CH,CW
~
HB+
+ CHZCN-
The behavior of the other five bases has not as yet been completely elucidated. Coetzee and McGuire (27) have calculated overall dissociation constants for perchloric acid in a'series of five nitriles and in acetone as solvents. Perchlorate ion appears to be virtually unsolvated in all six solvents. Values for the dissociation constant were of the general order of 10-3. Schaap et al. (1%) have published an eltensive article on electrochemistry in anhydrous ethylenediamine with a detailed discussion of salt and solvent effects on potentiometry. The treatment is mainly from a theoretical point of view, similar to that done by other authors for the acetic acid system. The discussion also includes informa-
tion on a reference electrode consisting of zinc amalgam-zinc chloride, and ethylenediamine saturated with lithium chloride. A\cid-base equilibrium in tert-butyl alcohol and application of such knowledge have been reported by hlarple and Fritz. The overall dissociation constants for perchloric, picric, and benzoic acids were calculated as well as that for 2,4-dinitrophenol. The effect of salts upon these values and upon hydrogen ion activity has been analyzed (97). A practical application of these ideas has been made in the addition of tetrabutylammonium bromide to the butyl alcohol solutions of weak acids. Addit'ion of the salt changes the proton concentration in such a manner that more recognizable end points in potentiometric titrations are obtained (98). Potentiometric titration curves of mineral and carboxylic acids as well as phenols have relatively little slope in the buffered region in this solvent and thus provide advantage for differentiating titrations (47). Crabb and Critchfield have also noted that tert-butyl alcohol appears to be the best of nine solvents for performing differential titration of phenols ( S I ) . Brown and Ives (16) have measured tmhe dielectric constant's of water-twt-butyl alcohol mixtures and concluded that this property does not correlate with compositions in a manner similar to that shown by mixtures of water with other lower alcohols. Acid-base behavior in dimethyl sulfoxide was studied by Kolthoff and Reddy (81). They found the hnsic strength of the solvent comparable to that of water, and that perchloric, hydrochloric, and sulfuric acids were completely dissociated in the solvent. The order of degree of dissociation of uncharged acids is not the same as that shown in water, however. Picric acid is more dissociated in the sulfoxide than in water, while acetic and benzoic acids are much less dissociated. The autoprotolysis constant for the sulfoxide Mackle was found to be 5 X lO-'S. and O'Hare have published data on t,he thermodynamic properties of this sulfoxide and on dimethylsulfone (95). The selfdissociation constant for acetic anhydride has been reported as 3 X a t 20" C. by Jander and Surawski (72) while that for 2-aminoethanol a t the same temperature is 5.14, according to Jacquinot-Vermease VOL. 36, NO. 5 , APRIL 1964
363 R
and Schaal ( 7 1 ) . They obtained this value qpectrophotometrically ; a previous value of 5.2 was obtained potentio. metrically. An investigation of dimethylformamide by Teze and Schaal (149) indicates that the amide has a range of 18 pH units. (A number of analytical applications are suggested besides the theoretical studies.) Paul, Guraya, and Sreenathan have also considered autoprotolysis in dimethylformamide and have noted high conductivities for protonic acids, which indicates hydrogen bonding n ith the Eolvent molecule^. Ionizstion constants for various acids show that fluorosulfonic acid is the strongest of the acids studied (110). The heats of neutralization of primary, secondary, and tertiary aliphatic amines with trichloroacetic acid in benzene were determined by Mead (104) using the thermometric titration technique. The order for A H , does not follow pKh(HzO) values or Taft g * values unless structure of the amines is also included in the consideration. Solvation and the importance of the referenre acid are considered to be factors in determining relative order of base strength. Chatten and Harris (21) have attempted correlations between halfneutralization potentials and pK,(H20) value for amines and phenothiazines in five organic solvents: acetic acid, acetone, acetonitrile, isopropyl alcohol, and nitromethane. Multiple relationships were shown by all solvents but acetic acid. The dependency of the relationships appears to be primarily due to molecular structure and hydrogen bonding. h’elson and Iwamoto (112) studied four redox systems to evaluate liquid junction potentials in nonaqueous solvents. They found the couple 4,idimethyl-1,lO-phenanthroline iron(I1) , 4,i-dirnethyl-l,lO-phenanthroline iron (111) to be the most generally useful. The reliability of the potentials is given as 5 mv., and the authors conclude that liquid junction potentials, when nonaqueous solvents are used, “are not as large as erroneously believed by many.” Mather and -4nson (99) showed that the anode and cathode reactions occurring when a current is passed through solutions of sodium perchlorate in acetic acid-acetic anhydride mixture produced hydrogen gas and acetate ions a t the cathode and hydrogen and acetylium ions a t the anode. The acetylium ion condenses with the acetic anhydride to form basic products that effectively remove hydrogen ion from solutions. The occurrence of the condensation reaction causes coulometric generation of hydrogen ions a t a platinum anode to be less than 100% efficient in this solvent system. 364 R
ANALYTICAL CHEMISTRY
I n a spectrophotometric study of proton exchange between o-chloroaniline and perchloric, hydrochloric, and sulfuric acids in dioxane-mater mixtures Koskikallio and Ervasti (83) found in confirmation of previous kinetic experimentq that the apparent acid strength of perchloric acid increases greatly, that of sulfuric acid increases slightly, and that of hydrochloric acid decreases as the amount of water in the dioxanewater mixture is decreased. A number of men have determined protonation constants by measurements made in nonaqueous solvents. Edward and Wang (39) give p K B R +values of -6.8 and -0.9 for propionic acid and propionamide from measurements made in sulfuric acid. Feakinq, Last, and Shaw (45) determined values for aminophosphazenes in water and nitrobenzene. The fully substituted compounds are as strong as amines. Joesten and Drago (73, 7’4) measured equilibrium constants for the formation of S,i?‘-dimethylamides. Gramstad (51) has studied phenol-pyridine association constants. I n formamide solutions Mandel and Decroly’ (94) have determined dissociation constants for formic and acetic acids: - logK, values were 5.49 and 6.82, respectively, a t 25OC. Davis (35) determined the relative strengths of all six isomeric dinitrophenols in benzene in terms of association with triethylamine. The relationship of strength in benzene and water is not a linear one. This is believed due to hydrogen bonding between the phenolic hydrogen and o-nitro groups. Gillespie and Robinson (48) in neutralizing sulfuric acid solutions of tetra(hydrogen sulfate) boric acid with various metal hydrogen sulfates concluded that the salts formed are best formulated as polysulfatoborates containing a six-membered ring of boron, oxygen, and sulfur.
CONDUCTANCE MEASUREMENTS
The equivalent conductances of lithium, sodium, potassium, rubidium, and cesium perchlorate in acetonitrile at 25’ C. were measured by Mine and Weblan (107) and limiting equivalent conductances were obtained by extrapolation. The limiting conductances generally increase with molecular weight, with the exception of the potassium compound. This behavior is also noted in wateralcohol solutions. Precise conductance measurements have been made on 12 salts dissolved in dimethylformamide. Conductance results for the perchlorates of eight univalent ions and for potassium iodide indicate complete dissociation. Lithium chloride and silver nitrate are in-
completely dissociated in this medium. Anions appear to be unsolvated, while cations move with a solvation sheath (12f). The conductance of lithium and ammonium iodide in 1-butanol a t several temperatures was measured by Venkatasetty and Brown (163). Both salts are found to behave as weak electrolytes. Esamination of conductance data for potassium thiocyanate and trimethylphenylammonium iodide in sulfolane indicates normal conductance by the salts in this solvent as well as in 2,4-dimethylsulfolane (41). Tetramethylammoniuni and triphenylmethylaminonium chlorides have been found to be strong electrolytes when dissolved in antimony trichloride. C1-. The ionization is R C I e R + However in the case of bornyl, n-decyl, cyclohexyl, diphenylmethyl, cinnamyl, and benzyl chlorides ionization occurs according to the scheme 2RC1 e R2Cl+ C1-, except in dilute solution, where the normal mode is observed (34). The conductivity of ethyl ether solutions of ethyl magnesium bromide were reinvestigated by Dessy and Jones (37). Bombara and Troyli have shown that sodium acetylide is a fairly strong electrolyte in liquid ammonia (15). Conductance measurements in 80 and 90% acetone for HCI, LiCl, NaCI, HC104, and XaC104 indicated that the last two compounds are completely dissociated (5). Measurements in 40, 60, 80,and 98.4weight 97, ethanol in water on the conductance of perchloric acid showed the equivalent conductance depended on tempernture, solvent composition, dielectric constant, and viscosity (49). Dawson et nl. (36) have shown that small quantities of acetic acid and acetates remain persistently in Nmethylacetamide even after extensive purification. I n the above solvent acetic acid was found to have a dissociation constant of the order of lo-* a t 40’ C. Conductance data on potassium chloride, potassium acetate, and mono- and dichloroacetic acid in this solvent were also obtained. Electron spin resonance, optical absorption spectra, and electrical conductivity of alkali metals in ethylenediamine were measured by Windwer and Sundheim (167). Evaluation of the data on potassium indicated ahsorbing and conducting species were not the same, the former being in a higher state of aggregation. Measurements in sulfuric acid-sulfury1 chloride mixtures on the degree of dissociation of weak bases gave data on nitrobenzene which agree with cryoscopic and spectrophotometric data. Standard heats and entropies of dissociation were calculated together with the autoprotolysis constant for the solvent (91).
+
+
STANDARDS A N D TITRANTS
TRIS, tri(hydrosymethy1)methylamine, has been used by Williams and Harley (155) as a standardizing reagent for perchloric acid in acetic acid solutions. They found it compares satisfactorily with biphthalate and diphenylguanidine. 130th potent'iometric and visual titrations were used. The solubility of the compound is reported as 1.7 moles per liter in acetic acid. -4new approach to titrants was taken by Maricle (96), who used coulometrically generated biphenyl radical anions in a totally nonaqueous solvent as a titrant for ant'hrawne, nitrobenzene, nitromethane, benzophenone, and azobenzene. Titrations of these compounds were quantiktive. Other hydrocarbons such as pyrene and halogenated hydrocarbonj did not yield quantitative results, although reaction obviously occurred, Side reactions were regarded as the cause of the difficulty. End poin.:s were detected potentiometrically and can also be appraised visually. Dimethylformamide was the solvent csed. Fluorosulfonic acid is reported to be a somewhat stronger acid bhan perchloric acid in acetic acid medium. This claim is supported by conductometric data (118). The acid forms a highly conductive solid with acetic acid. Lane (88) notes that although trifluoromethane sulfonic acid is of comparable strength to perchloricas a titrant inacetic acid, its greater expens8sprobably eliminates it as a competitor. Harlow and Wyld (68) have presented a new preparation of tetraalkylammonium hydroxide from isopropyl alcohol solu bions of pot:issium hydroside and the tetraalkyl ammonium chloride. Cundiff and Markunas (32) also offer a new preparation using freshly prepared silver oside. Afarple and Frit'z (96) prepared a very pure reagent using a silver oxide process arid an elaborat'e procedure to get rid of impurities. The potentiometric studies :arried out with this reagent gave titration curves which were reproducible within 2 to 5 mv. Stability of the base in water, isopropyl alcohol, tert-butyl alcohol, and pyridine was determined. Harlow has determhed the stability of 12 different tetraalkyl ammonium hydrosides and discusse'j this in relation to cation structure, solvent composition, water content, and teinperature (57'). He found cations containing a masimum of methyl groups to be the most stable, those containing etbj.1 and phenyl groups least stable. Harlow (56) has also related the conductometric and pot,entiometric titration curves of acids titrated with quaternmy ammonium hydroxide to cation structure, hydrogen bonding, and ion pair formation. Salvesen (184) notes the superiority of
benzalkonium hydroxides over methoxides as the titrant for weak acids in dimethylformamide. ELECTRODES AND INDICATORS
A modified glass electrode was constructed and tested by 13adoz-Lambling et al. ( 6 ) as the indicating electrode for the titration of bases in acetonitrile solution. The electrode contained 0.lM lithium perchlorate, 0.01M silver ion, 0.01X picric acid, and 0.005V diphenylguanidine, all dissolved in acetonitrile. The authors report that this electrode attains a more rapid and reproducible equilibrium with the solution than is achieved with conventional electrodes. A commercial combination glasscalomel electrode is reported to function satisfactorily in both water-miscible and water-immiscible solvents (60). Cheng, Howald, and Miller (22) compared the behavior of a conventional glass electrode with that of the platinum chloranil electrode in cells without liquid junctions. "Acid errors" of about 70 mv. in the glass electrode were due to chloride ion absorption. The effect of water on the potentiometric titration of hydrochloric acid with sodium acetate was also measured. Finally Harlow has noted that traces of potassium ion in solutions of tetraalkyl ammonium hydroxides can have profound effects on the shape of the titration curve for weak acids in nonaqueous solvents. The potassium tolerance of the glass electrodes appears to be ehtremely low (10 p.p.m.) and depending upon a particular electrode, potassium content of titrant, and other factors, unusual titration curves may be obtained which are related to interaction of the metallic ion and the electrode (56). I n the field of indicators, Kolling and Stevens (7'7) have studied the composition effect of acetic anhydrideacetic acid mixtures upon the halfneutralization number (pL,i 2) of the indicator bases acridine orange, brilliant green, rhodamine 13, and thioflavine T. Rhodamine 13 appears to be a suitable indicator for photometric titration of leveled and intermediate bases in a solvent containing 70/30 acetic anhydride-acetic acid. Pyridine-2-azo-pdimet hylaniline has also been suggested as a n indicator in the titration of leveled amines in acetic acid solution. The precision attainable is claimed to be superior to that attainable with crystal violet (20). Other indicators suggested for titration of bases in acetic acid and acetic anhydride are dicyanobis-( I ,lophenanthroline) iron(I1) and dicyanobis-(2,2-bipyridyl) iron(I I). .\minps, heterocyclics, and salts were titrated (139). Di-(2,4-dinitrophenyl)methane as a 1% diosane solution has been used
for the determination of organic acids by titration with alkali in nonaqueous solvents ( 8 2 ) . Kolling and Stevens studied azine, acridine, thiazine, and thiazole dyes as possible indicators for acid-base titrations in acetic acid. Of the nine dyes tested, Safranine 0, neutral red, and thioflavine T are satisfactory visual indicators for bases having pK, values greater than 7.9. The acridine dye, phosphine, is suitable for photometric titrations ( 7 6 ) . Palit has used a technique of eatracting a dye from a buffer with benzene and then using this solution to test the acidity or basicity of detergents and polymers at, micronormal levels of materials such as stearic acid. The color chmges observed are believed to be a result of salt formation (114). X uniform overlapping set of Hammett indicators consisting solely of primary anilines has been established by the addition of eight new indicators. pK values of two former indicators were reevaluated and employing the new p K values the Hammett acidity function ions were redetermined for 60% sulfuric acid (75). SOLVENTS
Reynolds, Little, and Pattengill have proposed a commercial solvent consisting principally of iV,N-dimethyllaurylamide and S,N-dimet~hylmyristamide as a suitable medium for the titration of mineral acids, carboxylic acids, and phenols. The solvent is readily purified by passage through an alumina column and most acids appear to show satisfactory titration characteribtics in this medium. Resorcinol and hydroquinone do not. The solvent appears amenable to differentiating titrations (126). The relatively strong. base 1,1,3,3tetramethylguanidine has also been proposed as a solvent for the t'itration of phenolic compounds. The commercially available solvent may be used without further purification. It' appears that the best titration curves were achieved with a solvent containing from 1 to 3y0water (154). For the titration of weak organic bases (pKb-l 1) dichloromet'hane or chlorobenzene has been recommended, together with a chloroform solution of p-toluenesulfonic acid. Dimethyl yellow was used as an indicator (138). Concentrated salt' solutions have been used for the titration of acids and bases which are weak in water. Rosenthal and Dwyer (130, 131) have studied acidity in 4 and 8 M lithium chloride solutions and eztahlished an acidity function, H,. pMH-Ho wa.* shown to be a constant for a given salt solution where p3IH equals -log total strong acid concent'ration. I t was al,\o shown VOL. 36, NO. 5 , APRIL 1964
365
R
+
that -log a H + = pH AE~/0.059 where A E L ~is the liquid junction potential and p H is the value measured by the glass electrode. 130th papers deal with other salt solutions as well and note the abilities and limitations of this type of solvent from both theoretical and pract'ical points of view. Leithe has used supersaturated salt solutions as the solvent for titration of HCO,-, HP04-*, HCS, and phenol with 0.5.11 S a O H , using azo violet as the indicator. The procedure is useful of mixtures of inorganic materials (89). Leyden, Smith, and I-nderwood have t,itrated aldehydes and ket'ones as bases in sulfuric acid. They used a conductometric technique. Anomalous stoichioniet,ry is observed in certain diketones. The method is limited by instability of many compounds in sulfuric acid solutions (-90). Acid-base reactions in molten potassium nit'rates have been studied by Shams El Din (140,141). Potassium dichromat? wa? used as the acid and potassium carbonate] sodium peroside, and sodium bicarbonate were used as the bases. The reaction for the carbonate is
An osygen-indicating electrode was used to follow the reaction] while a silver wire was used as the reference electrode. Theoretical results were achieved on all three bases. Potassium hydroside was also titrated with dichromate and gave excellent potentiometric titration curves (140). APPARATUS A N D TECHNIQUE
Xrnett and his collaborators (2, S, 4) have used a combination of solvent extractions from strong acid solut'ion and gas liquid chromatography to determine the basicit,y of a number of very weak organic bases including aliphat,ic and cyclic ethers. Results from the method are compared to values obtained by other techniques and a number of theoretical considerations advanced for the behavior of the osygen bases. Thermometric titration using a differential technique is advocated by Tyson, McCurdy, and Ilricker (161). They describe an apparatus which measures differences in temperatures between the reaction vessel and blank solution. With this device heats of react,ion can be measured rapidly and simply. 13ruckenstein and Gracias ( 1 7 ) have used a manual differential spectrophohrnetric technique to determine end point,s for aqueous and nonaqueous arid-base and chelometric titrations. The method is advocated for use in titrations where solvents unsuitable t'o 366 R
0
ANALYTICAL CHEMISTRY
potentiometry are employed or where reversible indicator electrodes are not available. Hanna and Sipgia (64) have applied their method of differential reaction rate analysis technique to the determination of primary, secondary, and tertiary amines and mixtures of aliphatic and aromatic amines. The method involves the reaction between amines and isothiocyanate to form thiourea. Connors and Higuchi (29) have presented three equations for the linear extrapolation of photometric indicator titration data to yield the end point and equilibrium constant for the titration reaction BHA
+I
$
IHA
+B
which governs the color change of indicator I in the titration of base B with acid HA. Higuchi et al. (62) have also considered the titration of amides and other weak bases in acetic acid. The data obtained are used to evaluate the practical and theoretical limitations for the quantitative and qualitative determination of such compounds. Svoboda has applied the general technique of constant current potentiometry to the nonaqueous titration of organic amines; two platinum indicator electrodes polarized by a constant 1-p a. current are employed. Peak-shaped t,itration curves are obtained which permit direct indication of the end point from meter readings. The electrode system gave reproducible results without pretreatment. A 1 to 1 mixture of m-cresol and acetonitrile was used as the solvent, perchloric acid as the titrant (148). Dimbat and Harlow have noted the display of a lumometric end point' in some nonaqueous titrat.ions (38). At the end point there is a sharp rise in luminescence intensity which is measured by a sensitive scintillation counter. .4 circuit for a simple coulometer using a constant applied potential and an integrating motor has been described by Head and Marsh (59),who have used it for both aqueous and nonaqueous acid-base titrations. Rough reaction kinetics for reactions of very weak acids and bases were obtained through the use of nonaqueous titrations. The compounds studied were acylamidines and related compounds (12). Zinc chloride and stannic chloride in acetic acid solutions have also been used as acidic solvents for promoting the acetylation of 2-naphthol with acetic anhydride. The kinetic form of the reaction is reported, analyzed, and compared with that found for ErGnsted acids in acetic acid (1%).
Ishidate et al. report a new titrimetric method based on difference in dielectric constant between acid, base,
and salts in a nondissociating solvent of low dielectric constant such as helane. The technique is similar to high frequency titration. Satisfactory rebults were obtained for the titration of amineb with picric acid, p-toluenesulfonic acid, and trichloroacetic acid (67-701. REDOX TITRATIONS
Bivalent copper, iodine, ferric ion, bromine, and antimoniate have been titrated potentiometrically with titanium(II1) chloride in dimethylformamide solutions. Viqual, spectrophotometric, and potentiometric end points were obtained. Chromous chloride was also used as a titrant for the entities mentioned above a$ \\ell a. for titanium(1V) and iodine monochloride (63)'
Azides in organic solventq may be determined by precipitating the compound as the silver azide, filtering and washing to remove all trace of solvent, osidizing the precipitate with excess ceric ion, and back-titrating with ferrous ion (52). The effect of various solvents on the formation of silver azides is considered. Cerimetry in nonaqueous solventb was discussed by Rao and Murthy, wh6 carried out titrations in acetic acid and acetonitrile. Potassium iodide could be satisfactorily titrated in a mixture of the two solvents, as could quinol. d platinum indicator electrode and either a glass or antimony reference electrode was employed (f 83). The same titrantsolvent-electrode system was used for the titration of ascorbic acid. Four equivalents were used (124). Ferroin, diphenylamine, methyl red, and Janus green were shown to be satisfactory indicators with quinol in acetonitrile, although the last two are not reversible (122). A direct titrimetric procedure for bromination of phenols is described by Huber and Gilbert. The end point IS measured by constant current potentiometry using platinum foil electrodes. Titrations were carried out in glacial acetic acid and the method is applicable to phenol and substituted phenols (66). LEWIS ACIDS
Hitchcock and Elving have used the high frequency technique to titrate organic bases with both aluminum and stannic chloride. Titrations with aluminum chloride were carried out in acetonitrile. Nitrogen bases were titratable and formed both 1 : l and 1 : 2 adducts. Oxygen bases could not be titrated, undoubtedly because of the basicity of the nitrile ( 6 4 ) . When stannic chloride is the titrant, benzene solutions may be used and under these conditions both osygen and nitrogen
bases may be titrated. The former include alcohols, ethem, ketones, and esters (66). Nitrobenzene has also been used as the solvent' ior high frequency titrations of Lewis acids. I Cefola, >I., A n a l . Chinz. Acta 29. 127 11963'1. (21) Chatten, L.' G.. H.&ris; L. E., ANAL. CHEM.34, 1405 (1962). (22) Cheng A. T. Howald, R. A,, Miller, I>.L., J . Phys. Cheni. 67, 1601 (1963). (23) Chiang, H. C., J . Pharm. Sci. 50, 885 (1961 ). (24) Choy, T. K., Quattrone, J. J., Jr., Elefant, AI., Anal. Chzrn. Acta 29, 114 119621. (22) Cluett, M . L., ANAL.CHEM.34, 1491 (1962). (26) Coetzee, J. F., Cunningham, G . P., McGuire, 1). K., Padmanabhan, G. It., Ibid.,34, 1139 (1962). (27) Coetzee, J. F., LlcGuire, D. K., J . A m . Chem. SOC. 67, 1810 (1963). (28) Coetzee, J. F., Padmanabhan, G. R., J . Phys. Chem. 66, 1708 (1962). (29) Connors, K. A . , Higuchi, T., Anal. Chim. Acta 25, 309 (1961). (30) Corsini, A., Yih, I. M., Fernando, Q., Freiser, H., - 4 N A L . CHEM. 34, 1090 11962). (31) Crabb, K. T., Critchfield, F. E., Talanta 10, 271 (1963). (32) Cundiff, R. H., Markunas. P. C.. ANAL CHEM.34, 584 (1962). (33) Daftary, R. D., Haldar, B. C., A n a l . Chzm. Acta 25. 538 1 19611. (34) Davies, A.'Gl,- eaughan, E. C., J . Chem. Soc. 1961, p. 1711. (35) Ilavis, 51. hI., J . A m . Chem. Soc. 84, 3623 (1962). (36) Ilawson, L. R., 1-aughn, J . W., Pruitt. 11. E., Eckstrom, H. C., J . Phys. Chem. 66. 2684 119621. ( 3 7 ) I)essy, R E., Jones, R 31, J . Org. C h e w . 24, 1685 (1959). (38) I)lnibat, 11, Harlon, G A , , !%SAL. CHEV 34,450 (1962) (39) Edward, J. T., Wana, I. C . Can. J . ('hem 40, 066 (1962). (40) Ellert, H., Jasinski, T., Weclawska I< =lcta Polon Phoriri 19. 75 (19621 (41) Eliassaf, S , FUOES, R S i . Lund, J I.:, J r , J Phys Chein 67, 1724 (1963). (42) Emelin, E. A,, Svostunova, G . P., Zovodsk. L a b . 27, 971 (1961). 143'1 Emelin. E. A4.. Pvostunova. G . P.. ' Tsarfin, I-: A . , Ibad., 27, 283 (1b61). (44) Icrshow, B. S . , Pokrovskaya, V. L., Plast. Xassy 1961, p , 65. (45) Feakins, I)., Lmt, W . A,, Shaw. R.A,, Chem. I n d . (London) 1962, p . 510. (46) Feuer, H., Vinvent, B. F., Jr., . ~ N A L . CHEST. 35, 598 (1963). (47) Fritz, J. S., llarple, L. W., ?bid., 34, 921 (1962). (48) G'illespie, R J., Robinson, E. A , , Can. J . Chem. 40, 1009 (1962). (49) Goldenberg, K . , Arnis, E. S., 2. Physzk. Chein. 30, 65 (1961). 150) Goldstein. G.. A N A L CHEK34. 1169 (1962). (51) Gramstad, T., Acta Chem. Scand. 16, 807 (1962). (52) Grove, E. L., Brarnan, R. S., Combs, H. F., Sicholson, S. B., A N A L . CHEM. 34, 682 (1962). (53) Gutterson, >I., AIa, T. S., Microchem. J . 5 , 601 (1961). (54) Kanna, J. G., Siggia, S., ANAL.C H s l l . 34, a47 (1962). (5.5) Harloy, G . A , , ?bid.,34, 148 (1962). (561 l b i d . . D . 1482. 1
I
I
-
(All 1
,
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Pharnt. Franc. 19, 81 (1961).
(121) Prue, J. E., Sherrington, P. J., m-
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(122) Rao, G. P., hlurthy, A . R. V., Z. A n a l . Cheni. 180, 169 (1961). (123) Ibid., 182, 358 (1961). (124) Ibid., 187, 96 (1962). (125) Rashbrook, R. B., Analyst 87, 826 (1962). (126) Reynolds, C. A . , Little, J., Pattengill. SI..ANAL.CHEM.35. 973 f 1!)63). (lz7)' Rink, SI., Lux, R.,& u t . .ipothe&r Ztg. 101,911 (1961).
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