Journal of Research of the National Bureau of Standards Vol. 54, No.6, June 1955 Research Paper 2593 Titrations of Bases with Diphenyl Phosphate in Some Aqueous and Nonaqueous Solvents! Marion Maclean Davis and Hannah B. Hetzer

Successful titrations of various bases with diphenyl phosphate in benzene, methanol, water, and mixtures of these solvents have been performed, in some cases potentiometrically (using the glass, saturated calomel electrode-system) and in other cases using indicator dyes. Several of the dyes (bromophthalein magenta E, 5-phenyl-9-diethyl Nile blue, and t he anhydro-base of tamarack green and Victoria blue B) have not been used elsewhere as acid­ base indicators. Diphenyl phosphate is shown to be a strong acid which is easily obtained and maintained as anhydrous crystals. It is not strong enough, however, to serve as a replacement for perchloric acid in titrations in acetic acid. Attention is called to the differing behavior of benzoic acid and diphenyl phosphate when the solvent is changed from water to methanol. Using a " nonleveling" solvent in place of water makes it possible to obtain separate, sharp end-points in titrating a mixture of two su ch acids. Other pertinent informa­ tion about diphenyl phosphate is summarized. 1. Introduction acidic behavior of this compound and of various related diaryl phosphates, even though they were As previously explained [1 ]2, we are engaged in first synthesized about 100 years ago and scattered developing methods for measuring acidity and article have appeared dealing with such topics as basicity in organic solven ts, with special emphasis their preparation, physical properties, and possible on so-called "inert" solvents. Two essential parts uses. of this research are (1) locating or preparing new chemical tools- indicator dyes and materials suit­ 2. Apparatus a nd Miscellaneous Materials able for use as standard acids or bases in such sol­ vents; and (2) finding out, by systematic studies, 2.1. Apparatus the behavior of different types of acids and bases in A titration apparatus similar to the one described various nonaqueous solvents. This paper deals with in [3], employing a gla s and a sleeve-type saturated a study of some of the properties of diphenyl phos­ calomel electrode, was used in the potentiometric phate, (C6HsO)2P(O)OH, with particular attention titrations. During titrations the temperature was to its use in titrations with bases in several aqueous usually within the range 25°±2° C. In most cases and nonaqueous solvents. 10-ml burets having 0.05-ml graduations were used. Interest in diphenyl phosphate first arose upon Melting points were determined by the capillary­ noticing in a catalog of chemicals available for tube method, using an automatically- tirred, elec­ experimental purposes [2] that, besides giving a trically-heated, high-boiling silicone oil and an strongly acid solution in water, it is easil.v soluble ASTM thermometer maintained at 3-inch immersion. in a variety of organic solvents, including benzene and carbon tetrachloride. The interest increased 2.2. Solvents after finding that the acid can be obtained and maintained as pure, anhydrous crystals without The organic solvents were all of a grade that wa difficulty. Very few studies have been made of the claimed to meet AC specifications. Acetic acid comparative effects of aqueous and nonaqueous and methanol were said to contain approximately solvents on the behavior of strong acids, and diphenyl 0.5 percent of water. Distilled water was boiled phosphate appeared suitable for such studies, as shortly before use, to remove carbon dioxide. well as a potentially useful standard acid for non­ aqueous titrations. Hitherto, picric acid and a few 2.3. Bases arylsulfonic acids have been the only commonly 1,3-Diphenylguanidine was recrystallized twice available strong, crystalline acids soluble in both from toluene and once from benzene-cyclohexane; water and a variety of organic solvents, and these the melting point was 149.0 0 to 149. ° C. Titration acids possess certain shortcomings. For instance, with perchloric acid in glacial acetic acid indicated picric acid is somewhat hazardou to handle unless a purity of at least 99.5 percent. Di-o-tolylguanidine 10 percent or so of water has been added, and the was recrystallized twice from 95-percent ethanol; yellow color of its salts is sometimes objectionable. melting point, 175.8° to 176.3° C. Titration in A further incentive to studying diphenyl phosphate aqueous alcohol with 0.1 N hydrochloric acid in­ is the meagerness of the existing Imowledge of the dicated its purity to be 99.9 percent; titration in glacial acetic acid gave the value 99.6 percent. A 1 Tbis rescarcb was supported iu part by the United Statcs'Air Force, through commercial high grade of guanidine acetate was used the Office of Scientific Research of the Air Research and Dcvelopmcn t Command. 2 Figures in brackets indicate the literature references at tho end of this paper without further treatment. Commercial high grad

309 tribenzylamine was recrystallized from ethanol; the commercial hydrochloride in 95-percen t ethanol, melting point, 92.8° to 93.2° O. A good grade of then pouring the ethanolic solution into water' the triethylamine was distilled from sodium hydroxide base, consisting of fluffy yellow needles, was wa'shed immediately before use ; the middle fraction was used. with water and dried at room temperature, first in Tris (hydroxymethyl)aminomethane was part of a puri­ the air and then in a vacuum desiccator containing fied sample that had been used in the detcrmination anhydrous alumina. of the aqueous dissociation constant [4]. For part of the work the indicator dyes commer­ Potassium acid phthalate was NBS Standard Sam­ ciall y available were not suitable, and three new ple for acidimetry. Sodium acetate in glacial acetic basic indicators were used. These call for some acid was prepared by dissolving anhydrous sodium additional comment. carbonate of analytical grade. The solutions con­ taining sodium methylate were prepared from the a. Tamarack Green Anhydro-Base solid compound. Tamarack green anhydro-base (fig. 1, formula II, 2.4. Indicator Dyes X = 01) has not been described previously, although the parent compound (X = H ) has been known for Both acidic and basic dyes were used in the titra­ many years. Both compounds belong to a class of tions. Bromophthalein magenta E (fig. 1, formula I, derivatives of triphenylmethane, known as "Homolka X =00002Hs) was prepared at the Bureau; its ap­ bases" or "anhydro-bases", the first one of which was plication in some nonaqueous acid-base titrations prepared by Homolka; later this class of compounds was described previously [IV The other acidic was studied by other workers, notably by Baeyer and indicator dyes- phenolphthalein, p-naphiholbenzein Villiger [5,6),4 The parent compound was first pre­ (a-naphtholbenzein), and the various sUlfonephthaleins pared by Meldola, who prepared its hydrochloride were all commercial materials; the sulfonephthaleins by the reaction of diphenylamine with benzotri­ were used as their sodium salts. chloride in the presence of anhydrous, po,,-dered zinc i Most basic indicator dyes arc supplied commer­ chloride and named it "diphenylamine green" [7]. cially as a salt (generally, the chloride). Azo dyes, An impure specimen of the hydrochloride which he however, are often available only as the free base, had prepared earlier by the oxidizing action of ail' on though sometimes also available as a salt. For a mixture of diphenylamine and benzyl chloride had reasons of solubility, the salt is more convenient for been introduced into commerce under the name aqueous titrations, and the free base is necessary for "viridine." Meldola did not isolate the pure an­ titrations in benzene and similar solvents. Either hydro-base, although he observed the red-brown form of the dye dissolves readily in alcohols. In the color, due to its presence, which resulted when a work described here all basic dyes were used as the solution of the salt was shaken with benzene and free bases with the exception of a-Naphthyl red sodium hydroxide solution. In later, more complete (4-phenylazo-l-naphthylamine hydrochloride) and studies, Baeyer and Villiger obtained the anhydro­ the three Nile blues. base in an analytically pure condition; they applied p-Ethoxy ch ry s oidin, 4-(p-ethoxyphenylazo )-m­ the name "viridine" to the pure dye [5].5 phenylenediamine, was isolated by adding an excess Tamarack green hydrochloride was prepared with of lO-percent sodium hydroxide to a hot solu tion of the substitution of o-chlorobenzotrichloride for benzotrichloride, but otherwise essentially as de­ 'This indicator is now obtainable from Distillation Products Industries, scribed by Meldola [7]. 6 The hydrochloride was listed as Eastman Organic Chemical No. 6810 (tetrabromophenolphthalein cthyl ester). converted to the picrate which, in turn, was con­ ys Hs verted to the anhydro-base, as outlined in [5], O-H N except for the use of benzene in place of ether. The I properties of tamarack green (both hydrochloride sr-a-sr and free base) resemble closely those described for the parent compounds [5, 7, 8]. 611 0 -c C>-~ 'These seldom·mentioned compounds give yellow to red solutions in solvents I II like benzene. The nitrogen atom attached to ring 1 (see formula II) is rather strongly basic, and on the addition of a proton forms a resonant cation that is ~0 deeply colored (red to green, depending on the substituents on the nitrogens and s~-0-sr in the rings). ('l'his is a simplifi ed discussion; it is probably not correct to as­ II I snme tbat in solvents like benzene the proton is completely detached from the o N-H anion of the acid molecule of which it was originally a part.) Anhydro·bases I tend to add water to form the well-known colorless compounds called "carbinol CsHs bases" or "color bases." Alcohols, ammonia, and primary or secondary amines give similar colorless compounds, R adding to the nitrogen of ring 1, and tbe rest of the molecule (-OR , - OR, or -NR,) becoming linked to the methane I IT :m carbon atom. A similar tendency is responsible for the fading by alkali of fabrics dyed with triphenylmethane dyes, but it has been reduced by introducing a F I GURE 1. Formulas of indicators. bul ky substituent (for instance, CI, NO" or SO,H) in an o"-position [6]. The o"-CI in tamarack green was introduced with such an objective in mind. 1. Bromophthaleill magenta E (tetrabromophenolphthalein ethyl estcr); ' Rill, Branch, and Patapofi' have also prepared and studied the anbydro-base X~COOC,H5. and called it "viridine" [8]. II. Tamarack green anbydro-base (anbydro-N,K'-diphenyl-p,p'-diamino-o"­ , Tho dye was prepared almost 10 years ago by M . M. Davis and P. J. Schuh- I cblorotriphenylcarbinol); x~ Cl. mann, for exploratory studies. The base was not analyzed and may have con­ III. 5-Phenyl-9-diethyl Nile blue (5-phenyJimino-9-dicthylami nobenzo[a]­ tained a little carbinol. Since preparation, its behavior as an indicator has not phenoxazonium chloride); X~ Cl . visi bly changed.

310 r b . Victoria Blue Anhydro-Base phenyl pho phates; diphenyl phosphate \ as ob­ tained only as a colorless, viscous oil that wo uld not Comparison of (.he formula of Victoria blue solidify in a freezing mixture [15]. B anhydro-base, C6HsN = CIOH6 = C[CsH4N(CHa)2]z, Rapp, repeating J acobsen's procedure, succeeded with that of tamarack green anhydro-base reveals in obtaining crystalline diphenyl phosphate by seed­ several differences in their structures, but it can be ing an aqueous solution with a little sodium diphenyl seen that they should function similarly when used phosphate, and reported the melting point of the as acid-base indicators. Victoria blue B anhydro­ acid as 56° C [16]. Autenrieth used a sligh t modi­ base was fu'st prepared by N athansohn and Muller fi cation of the Jacobsen procedure and reported that [9], who erroneously believed it to .be the carbinol. N oelting and Saas, after later studlOS [10], r:eported diphenyl phosphate is easily obtained in the crystal­ that the anhydro-base cannot be converted mto the line state if not contaminated with phenol ; his final carbinol or to a carbinol ether or corresponding product, obtained from an aqueous solution as 7 shimmering leaflets, had the melting point 61 ° to amino derivative. 63 °-a value several degrees higher than that of To prepare the anhydro-base, 4 g of. the co 11.1- Rapp [17]. The mystery of the divergent melting mercial salt (Color Index No. 729) was dIssolved m points was cleared up by H ocflake's observation that abou t 800 ml of hot water, the solution was filtered diphenyl phosphate has two crystalline forms, one to remove insoluble material, and the filtrate was being anhydrous and the other, a dillydrate [18]. N s.o~ium stirred mechanically while 2 to 3 ml of 6 The products of Rapp and Autenrieth were pre­ hydroxide was added. The dark brown preCIpItate sumably different mixtures of the two forms. Hoe­ was filtered and washed with the aid of suction, flake's anhydrous diphenyl phosphate was obtained air-dried, and then recrystallized from benzene. by heating the crud e acid (prepared by Jacobsen's 0 c. 5-Phenyl-9-diethyl Nile Blue (5-Phenylimino-9-diethylamino­ method) in vacuo to 60 to 80° to remove water of benzo[aJphenoxazonium Chloride, C 26H2.ClN30 ·H20) hydration, followed by solu tion of the acid in cold chloroform, then reprecipitation with ligroin. ~he The general structure of the three N il ~ blues used described the anhydrous acid as fine needles meltmg in the work is shown as formula III m figure 1. at 70°. The dillydrate was obtained as small leaf­ The only importan t difference among them is in the lets, melting at 51°, by recrystallizing the anhydrous group R , which is !~ydrogen in Nile bJue A . (eI No. acid from water. I-Ioefl ake furthermore stated that 913), benzyl in Nile 1?lue 2B .(CI No. 914), and anhydrous diphenyl phosphate takes up two mole­ phenyl in 5-phenyl-9-dlethyl .Nile blue. All.of th.e cules of water in the air, but loses this water again at salts give blue solutions wInch tum red (pmk, If 30° also over sulfuric acid. Bernton [19] concurred dilute) when the free base is liberated. As would in Hoeflake's description of the two crystalline modi­ be expected, the free base is strongest when R equals ftcations of diphenyl phosphate, but reported th at H and weakest when R equals phenyJ.8 . the hydrate is very stable in the air. Phenyl diethylKile blue was a ve~'y pme mat9nal Various practical applications have been s uggeste.d prepared in the research laboratones of Amel'lcan for cliphenyl phosphate. Bemton found that It Cyanamid Company [12].9 com bi.nes in 1: 1 ratio with different types of bases (ammonia, methylami.ne, diethylamine, aniline, 3. Diphenyl Phosphate amino acids), forming well-crystallized sfl:lts [19]. Diphenyl phosphate, together 'with other dlaryl ~l.lld 3 .1. General Discussion monoaryl phosphates, has been tested as a pOSSible reagent for the determination of iron; aryl pllO s ph ~tes Diphenyl phosphate has been sLudied in ter­ of varioLls metals were prepared [20]_ Potassmm mittently since 1866. It was first detected by diphenyl phosphate, like the po tass! L!ffi salts of s?me Hlasiwetz and Grabowski in a mixture formed by related acids, increases the solubIhtws of vanous the reaction of phenol with phosphorus pentoxide organic compounds in water [21 ]. Diphenyl phos­ [13], and was first isolated, as. a brown oil, by Glu tz, phate has been used as a substrate m StudIOS of after treatment of phenol WI th phosphorus penta­ em~yme action, and potassium salts of closel:f l'cl ~ted chloride [14]. Next, Jacobsen studied the step\yise acids have been reported to destroy cer-tam kmds reaction of phenol with phosphorus oxychlonde, of bacteria. In industrial applications, diphenyl fo llowed by hydrolysis of the intermediate product ~, phosphate has been reported to retard the cOlTosion and thus obtained crystalline monophenyl and tl'l- of steel sheets improve lubricatin g oil for steel-light 7 Evidently a naphthyl group, like an o"-substituted phenyl group, impedes metal bearing~, and excel dibutyl phthalate in re­ the addition of a four th gronp to the methane carbon atom of triphenylmethane ducing the permeability of ethylcellulose to water. d~C~ee fig ure S and related discnssion. 'l' he pH range of tIle commercial Nile There is no evidence in the literature, or from our blue A used is given as about 10.2 (blue) to 13.1 (purple), and that of the com· mercial Nile blue 2B, as about 7.2 to 8.6. (Howcver, Kolthoff reported the pH experience in ueing diphenyl phosphate, that it . is range of Nile blue A to be 9.0 (blue) to lOA (red) rIll.) . , Phenyl diethyl Nile blue is one of a series of bcnzophenoxazlDc dyes prepared toxic. However considering th at certain commerClal for possible chemotherapeutic applications and was m ~de available, tllrough th ~ esters of pho s ph~r i c acid and p yrophosphoric a.cid, courtcsy of J _ J . Denton, for investigations of it~ utility as an acid-base indicator in nonaqueous media. As D r. Dcnton has po lilted out m corresponden ce, the as well as some related compounds, are known to compounds known commercially as Nile blue dyes ar{' dye mIx.tures, whereas t~e specimen of 5-plle nyl-9-diethyl Nile blue supplied was analytlCally pure, and lS interfere with normal physiological reactions, it therefore properly referred to by its strict chemical name. However, a shorter seems prudent to avoid contact of diphenyl phosphate name is needed for common usa,ge, and the name used in t.his paper was agreed on. For ordinary titrations a less pure material should, of course, be satisfactory. with the skin and inhalation of fin e particles.

311 3.2. Purification allowing ethyl acetate to reflux in the apparatus. Over a period of 9 days the sample was repeatedly a . Anhydrous Diphenyl Phosphate heated, cooled, and reweighed. During this time "Anhydrous" diphenyl phosphate was supplied by the weight fluctuated; the greatest gain, which the Dow Chemical Company, through the kindness occurred during rainy weather, was 2.2 percent. of C. W. Robertson. The material as received con­ Water was then placed in the absorption bulb tained a little water of hydration, was slightly tinged of the Abderhalden, the bulb was attached to the with pink, and had a faint odor of phenol. Diphenyl apparatus with partial evacuation, and the sample phosphate is fairly soluble in hot cyclohexane; the was remelted. In 3 days it gained more than 10 amount of cyclohexane needed can be reduced a good percent in weight; for complete conversion to the deal by adding 10 percent or so of benzene. Any dihydrate a 14.4-percent gain in weight would have water present in the sample is removed if the cyclo­ been necessary. An accurate measurement of the hexane or cyclohexane-benzene solution is boiled for total gain was impossible as the sample rapidly a while. The pink material goes into solution and lost weight while on the balance; the loss became tends to remain with the diphenyl phosphate, but less rapid after the weight had dropped to about a it can be removed by repeated digestions with de­ 5-percent increase over the original weight. During colorizing carbon. A good way to remove most of 3-days' storage in the Abderhalden apparatus over the pink contaminant is to add less than the amount phosphorus pentoxide at room temperature the of hot solvent needed for complete solution of the sample became crystalline once more and regained acid' the bulk of the solution is then nearly colorless, its initial weight. the ~ndissolved acid forming a pink oil which adh eres to the bottom of the vessel when the colorless portion b . Diphenyl Phosphate Dihydrate is decanted off. When the cooled solution is seeded, Some anhydrous diphenyl phosphate was dis­ the acid quickly crystallizes in colorless needle.s; solved in hot water and, after cooling in a refrigera­ without seeding, it usually separates first as an OJ1. tor the solution was seeded. Leaflets immediately The purified acid has no odor. The purest pro.duct sta;ted to form. After filtration these were partly obtained in our work melted to a colorless lIqUId at dried between filter papers, then kept in a desic­ 68.5 0 to 69.0° C, with slight softening from about cator at 60-percent relative humidity (see footnote 68.0°; the untreated acid melted about a degree lower, 10). After 12 days at 25 0 C, during which the yielding a pale pink liquid. Titrations in water (s~e weight of the crystals decreased about 1.6 percent, section 4.1) indicated that the best anhydrous aCId the molecular weight found by tit.ration corre­ was about 99 .9 percent pure (see also section 4.5). sponded closely to the theoretical molecular weight The crystals of anhydrous diphenyl phosphate of the dihydrate (286.2). The same result, within showed no tendency to pick up water at room tem­ the limits of error, was obtained after further ex­ perature (which was usually in the range 20° to posure of the material to 60-percent relative hu­ 25° C). In one experiment a gram was exposed midity for 25 days. On the assumption of an i,nert to the laboratory atmosphere for 36 days, during impurity (most likely, occluded water) the tItra­ which the relative humidity probably averaged tions indicated the dihydrate to be approximately about 30 percent; the sample was reweighed :;tt 99.8 percent pure. A sample quickly liquefied when various intervals, 7 times in all, and no change m the tube containing it was placed in the melting­ the weight could be detected. In a. se.cond . ex­ point bath at 49 ° to 50 ° C; bubbles formed, but periment a 0.4- g sample showed no gam III weIght they soon disappeared, leaving a clear, colorless during a total exposure of 7 days to about 60-percent liquid. However, the dihydrate did not melt relative humidity at 25° C, or duting a subsequent sharply when heated gradually from a lower tem­ 8 days' exposUre to about 80-percent relative perature. For example, a sample placed in the humidity.10 These experiments do not support bath at 40° showed noticeable softening at about Hoeflake's assertion [18] that anhydrous diphenyl 41.5° and gradually liquefied, mostly at about phosphate takes up ~wo molecules of water .in the 49° to 50°. air. It was not possIble to repeat the expenments As mentioned in section 2.2.a, a sample of fused, by which Hoeflake reached her conclusion, as they anhydrous diphenyl phosphate that had gained were not described. Our experiments indicate that weight in a moist atmosphere returned to its original no difficulty should be experienced in maintain­ weight within 3 days when stored over phosphorus ing diphenyl phosphate as anhydrous crystals pentoxide. Slightly moist crystals of the dihydrate under usual laboratory conditions. (about 2 g) likewise lost weight when stored in a However when in the fused condition anhydrous desiccator over anhydrous calcium sulfate. Within diphenyl phosphate was found to gain weight 100 days, while kept at a temperature that was temporarily in a moist atmosphere. About a gram automatically maintained dose to 25° C, the sample of the acid was placed in a porcelain boat in an lost 13.2 percent in weight; if the sample had been Abderhalden drying apparatus which was open completely air-dry the loss during conversion to the to the air, and the sample was heated to fusion by anhydrous acid should have been 12.6 percent.

!O Approximately 6O.percent.relative humidity was adlieved b~ keeping the The molecular weight (by titration) of the dried sample in a desiccator contBinmg saturated aqueo us sodIUm bromIde ID coutaet acid was the same as that of anhydrous diphenyl witb sodium bromide dihydrate. For 80-percent reJative humidity. saturated aqueous ammonium chloride in contact with the solid salt "as used. phosphate, within the limits of error. 312 Further study of the interconversion of anhydrous phoric acid and monophenyl pho phaLe in waLer and dihydrated diphenyl phosphate would be of at 18° C. that replacing one of the hydrogen atoms interest. However, the evidence seems sufficient of orthophosphoric acid by a phenyl group in creases to justify the conclusion that the anhydrous form the acidity. Subsequently, Kumler and Eiler [23] of the acid can be maintained at constant composi­ m easured the relative acidic strengths of four mono­ tion more conveniently and reliably than the di­ alkyl ester (methyl, ethyl, n-propyl, and n-bu tyl) hydrated form. and the corresponding dialkyl esters of orthophos­ phoric acid in water at 25 ° C. Their data showed 3.3. Solubilities in Various Solvents that the mono alkyl esters are stronger acids than phosphoric acid, and that the dialkyl esters are In studies being made at the Bureau it is useful stronger still. Morton's data indicate that mono­ Lo know the approximate solubilities of anhydrous phenyl phosphate is stronger than the monoall\:yl diphenyl phosphate in benzene, cyclohexane, and phosphates investigated by Kumler and Eiler, al­ water . The results of some measurements are though the experimental conditions used in the two summarized in table 1. Other available information sets of investigations were too different to permit a numerical comparison . TA llLE 1. Approximate solubilities of diphenyl phosphate in various solvents at 25° C It was therefore to be expected that diphenyl phosphate would show strongly acidic behavior in

g of solute prJ' 100 ~ water, and this proved to be the case. In fact, the Eo l\'('ni of soh'l'nl curves obtained from potentiometric titrations with 0.1 N alkali in water were practically iden tical with Acrj,olH' ______. ______100 • Benzcnr ______. 170 b /' curves obtained for hydrochloric acid of equivalent Carbon trtmoh lorid r ..... 4 " strength . Moreover, th e measured pH values of Cyclohexane ______1.5 b d e Ether...... _...... __ __ 100 • several aqueous solutions 11 of diphenyl phosphate "ra ter._ ------.------:3 II. b d having concentrations in the range from about 0.005 M to 0.1 M were in very good agreement with the • 'l'his valur, given in [2J, was evidently obtained using hydrlted diphrnyl phospha te. pH values calculated on the assumption of complete b Determined in this laboratory by a titration proced ure, using anhydro/t$ ionization. dlphenyl pbosphate. e The value gi ven in [2J is 10; see footnote a. As will be shown in section 4, diphenyl phosphate d The sarno value was obtained by weighing the resid ue left after evaporation of solvent. also behaves as a strong acid in benzene, and appears • Infrared measurement~ by James Stewart of the Bureau indicated a benzene to be stronger in methanol than in water. However, content less than 0.01 percent. in acetic acid it behaves like a rather weak acid. on solubilities is included in th e table. Early 3 .5 . Stability in Aqueous and Alcoholic Solutions investigators reported that diph enyl phosphate is easily soluble in alcohol, benzene, chloroform, and Experiments to be described in section 4 showed ether [14 to 17]. that diphenyl phosphate can be titrated with bases The solubility of diphenyl phosphate in benzene very successfully in benzene, methanol, water, and is difficult to measure because supersaturation tends mixtures of these solvents, and therefore meri ts at­ to occur and a slight ch ange in the temperature has a tf'ntion as a poss ible acidimetric standard in some or marked eff ect. However, gooc agreement was all of these solvents. However, if it underwent obtained in duplicate experiments. Partial hydra­ rapid chemical change in solution, such as hydrolysis tion of the sample used may account for the much or alcoholysis, its usefulness would be limited. lower solubili ty reported in [2] . The literature contains fragmen tary and partially The earliest investigators [14 to 16] stated that conflicting statements abou t the stability of diphenyl diphenyl phosphate is practically insoluble in water, phosphate and closely related esters in aqueous and but later [17] 1 part of the acid was reported to be alcoholic solu tions. Glutz [14] stated that only one soluble in 35 parts of water. This value is close to phenyl group is split off when triphenyl phosphate is the one measured at the Bureau and also reported treated with strong potash, and Autemieth [17] re­ in [2]. Crystals of anhydrous diphenyl phosphate ported the same result from heating various triaryl liquefy as soon as they are brought in contact with phosphates with an equivalent of alcoholic potash. a little water. On shaking, the liquid separates into As further evidence of the stability of diaryl phos­ minute, scarcely visible drops. As a result of this phates, Autemieth claimed that they are not much behavior, in some early aqueous titrations of diphenyl saponified in dilute aqueous acidic or alkaline olu­ phosphate incomplete solution of the samples oc­ tions, even on long heating. CUlTed without detection, even when the quantity However, Plimmer and Burch [241 concluded that used was much below the amount required for a all three phenyl esters of orthophosphoric acid are saturated solution. It proved advantageous to completely hydrolyzed on heating with acid or alkali di ssolve the crystals in a little ethanol, then dilute to 85° for periods as long as a week; 5-percent solu­ to the desired volume with water. tions of the esters in 1 N hydrochloric acid or so dium hydroxide were used, and at appropriate intervals 3.4. Acidic Strength analyses for phenol and inorganic phosphate were Some years ago Morton [22] co ncluded from com­ made. The comparative behavior of some alkyl parative potentiometric titrations of orthophos- II Each solution contained about 10 percent of ethanol by volume.

313 phosphates was also studied. The general conclu­ T A B L E 2. Stability of 0.1 NI diphen yl phosphate in water at sions were as follows: (1) All three phenyl esters of 25 ° C ' orthophosphoric acid can be completely hydrolyzed I b y dilute alkali, though a trialkyl phosphate is not T ime ~ I a p sed since \ Molarity by titration h ydrolyzed beyond the dialkyl ester ; (2) all phenyl preparation with 0.1 N alkali and alkyl esters can be completely h ydrolyzed in Days dilute acid. Plimmer and Burch also mentioned that 0______0.0987, 0.0989. 12______n.0985. 0.0986, 0.0985. at 40° diphenyl phosphate was appreciably h ydro­ 98 ______0.0989, 0. 0989. lyzed by alkali in 24 hours. 156 ______0.0988, 0.0988. Smith and Stohlman [25] found that I-percent • T he acid was dissol\' ed with the aid of 20 ml of 95-pereen t ethanol, the final triphenyl phosphate in half-saturated aqueous so­ volume being 200 m!. dium hydroxide is completely h ydroly zed to phenol and inorganic phosphate after about 5 hours' reflux­ ing ; the three isomeric tricresyl phosphates behaved 4 , Use of Diphenyl Phosphate in Acid-Base analogously, except that they seemed to be more Titrations in Various Aqueous and Non­ resistant to hydrolysis. In absolute ethanol tri-a­ aqueous Solvents cresyl phosphate was hydrolyzed to di-o-cresyl phos­ phate at 25° 0 to an extent that was proportional 4.1. Titrations of Diphenyl Phosphate With Sodium to the concentration of potassium h ydroxide presen t, Hydroxide in Water but no removal of o-cresol from either di-o-cresyl phosphate or mono-a-cresyl phosphate was detected , The resul ts of some titrations of two differen t lots of purified diphenyl phosphate are presented in t able even when the concentration of alkali was as high 12 as 2. 5 N and the experiment lasted 9 days. 3. According to the resul ts shown in the table, both The conclusions of Breusch and K eskin [26] ar e lots of diphenyl phosphate were approximately 99 .9 partially in conflict with those of Plimmer and Burch . percent pure. Theil' resul ts from refluxing various triaryl phos­ phates (0.01 Nl) in aqueous sulfuric acid or sodium T A B LE 3. Titm tions of di phenyl phosphate in water' hydroxide, with periodic analysis of small samples for inorganic phosphate, may be summarized as fol­ NaO H b Weight of lows: (1) Triar yl phosphates are easily saponified Pu r it ~T sample c aleu- i Found in hot 1 N alkali, but are stable toward sulfuric acid lated (1 to 5 ]I'!), even on hea ting fo r hours; (2) tribenzyl phosphate is saponified by both acid and alkali. In Titrations using phenolphthale in explanation of th e results, the authors suggested that 9 ml ml % tribenz:vl phosphate is a true ester, but triaryl phos­ ' 1. 0025 35.68 35.69 100. 03 , O. 9898 35. 23 35. 18 99.86 phates are rather to be regarded as acid anhydrides. , 1. 0437 37. 15 37. 12 99. 92 ' 1. 0510 37. 41 37.36 99.87 Enz ~-ma tic h ydrolysis of diphenyl phosphate and ' 1. 0966 39. 03 39. 00 99. 92 other esters of phosphoric acid has been studied , ' 1. 0307 36.68 36.59 99. 75 ' 1. 0193 36. 28 36. 23 99.86 mostly in J apan; th e extent of h ydrolY3is reportedly ------Ave rage. __ __ 99.89 varies with the nature of the phosphatase and the --- - - I ----- substrate and also with the pH. An interesting study has been m ade of the mechanism of h ydrolysis P otentiometric titrations of trimethyl phosphate in alkaline, acidic, and ini. tially , O. 3576 12. 75 12. 74 99. 92 neu tral aqueous solu tions, with the use of water con­ '. 3222 [1. 49 11. 47 99. 83 d.3962 14. 13 14. 15 100. 14 t aining an excess of 0 18 [27] . The data indicate that d. 42 13 15. 02 15. 01 99. 93 in alkaline solution hydrolysis proceeds en tirely d. 4838 17. 25 17.24 99.94 through rupt ure of the P - O linkage, whereas in • In titrations with phenolphthalein, the samples were dissolved in 15 ml of acid solution rupture involves mainly , if not exclu­ 95-percent ethanol and 85 ml of water; ill the potentiometric titrations, the ce· sively, th e 0 - OH3bond . specti ve amoun ts of ethanol and water were 10 and 90 mt. b '[' he n ormality (about 0.1 N) was determined by standardization against While recrystallizing diphenyl phosphate from N B S Standard Sample of potassium acid phthalate. water we observed that protracted boiling of an " d The superscripts c and d indicate t wo different lots of purified diphen yl aqueous solution produces a slight milkiness and an phosphate. odor of phenol. On the other hand, no change in 4.2 . Titrations of Bases With Diphenyl Phosphate in the acidity of an essen tially aqueous (see foo tnote 11 ) Glacia l Acetic Acid 0.1 M solution was observed during 156 days' storage at 25° 0 (see table 2). It seems clear that, while high At the present time 70-percen t perchloric acid is temperatures must be avoided, diphenyl phosphate the only generally satisfactory reagen t for preparing is sufficien tly stable in neutral aqueous or alcoholic standard acid solutions for titrations of bases in gla­ solutions to warrant further investigations of its cial acetic acid. Acetic anhydride must be added to u tility as an acidimetric standard at room tempera­ destroy water introduced with the reagen t, and the ture. More rigorous tests, including intercompari­ concen tration of the solution must always be deter- sons with other acidimetric standards, will of course be necessary before it can be recommended for highly 11 See also table 2 and section 4.3. As ' \"as pointed out in section 3.3 it is ad visa­ ble to add a little alco hol to ensure complete solution of samples of diphenyl aecurat e m easurements. phosphate used in aqueous titrations. 314 ------~ mined by titration against a primary basic standard, 4.3 . Comparative Titrations of Diphenyl Phosphate such as diphenylguanidine, potassium acid phthalate, a nd Carboxylic Acids With Sodium Methylate in or sodium carbonate. Clearly, it would be very grati­ Solvents of Various Compositions fying to find a solid, anhydrous acid, comparable in strength with perchloric acid. Diphenyl phosphate a . Potentiometric Tilrations would not be expected to equal perchloric acid in In preliminary potentiometric titration of 5-. or strength, but it was desirable to ascertain the com­ 10-ml portions of 0.1-.111 diphenyl phosphate wlth parative behavior of the two acids in glacial acetic sodium methylate (approx 0.11 5-1\1) in me Lhanol, acid. the precision was very good (usually ± 0.01 ml), The titration curves presented in figure 2 show that even though the temperature of the laboratory diphenyl phosphate was not satisfactory as a replace­ was subject to fluctuations (mostly in the range ment for perchloric acid in the potentiometric titra­ 25 0 ± 20 C) and methanol has a rather high coefficient tion of di-o-tolylguanidine and guanidine acetate in of thermal expansivity. Carboxylic acids, of which glacial acetic acid. Curves obtained in the titration acetic, benzoi c, and palmitic acids arc typical, arc of pota sium acid phthalate and tribenzylamine with often dealt with in nonaqueous titrations. Solu­ diphenyl phosphate were very similar to those shown.l 3 tions of these three acids were also titrated wi th There is a recognizable break in all of the titration sodium methylate in methanol, under conditions curves in the expected region (8.4 ml), bu t in the paralleling those used in titrations of diphenyl titrations with diphenyl phosphate the break is not phosphate. The titration curves indicated that steep eno ugh for any but crude analytical proced ures. l4 cliph enyl phosphate is a much stronger acid than Dyes (for example, crystal violet) which arc com­ the three carboxylic acids, in h armony w'ith the monly used as indicators fo r titrations in glacial ace­ evid ence that diphenyl phosphate is a strong acid tic acid change color in the emf region - 450 to -550 in water. mv, and plainly would be completely useless in titra­ The ready solubility of both diphenyl phosphate tions with diphenyl phosphate. and benzoic acid in benzene, metha nol, and watcr In a few miscellaneous titrations of sodium suggested the possibili Ly of observing what alLera­ acetate with diphenyl phosph ate in glacial aceLic tions in the titration curves wouldl'esult from chang­ acid mixed with other solven ts-benzene, dioxane, ing the solvent gradually from methanol Loward ethanol, water-the results were much the same. benzene on the one hand, and toward water, on the In general, the microvoltmeter readings were very other. ' For this, 5-ml samples of 0.1 1\£ acid in steady excepL for mixtures con taining a high pro­ methanol were diluted Lo approximately 50 ml wiLh portion of benzene.15 various proportions of methanol and benzene 01' methanol and waLeI'. The samples were then 13 The two bases referred to in figure 2 did not give a preCipitate with either diphenyl phosphate or perchloric aCid; the perchlorates, but not diphenyl phos· tiLrated with sodium methylate, dissolved in a phates, produced UPOIl titrating potassium acid phthalate and tribenzylamine preCipitate under the conditions usecl. sol vent of approximately the' same composition as "It should be noted that the titrations were onl y of au explo ratory character, that of the solution being ti trated. The results and tbe results with diphenyl phosphate would probably have been a little better if the solutions had been completely anhydrous. '1'he solu Lio ns were arc shown in figure 3. 16 estimated to c011tain about 0.6 percent of water. " It is not unlikely that diphenyl phosphate associates with acetic acid through It is evident from the figure that for an aqueous bydrogen bonding. No one seems to have made any studies of aryl phosphates solution of either acid a negative value on Lh e emf which throw any direct light on their tendency to undergo self-association or otherwise partiCipate in hydrogen bonding. but trlbutyl phosphate has been scale would correspond to a 10\"1 pH valu e, and an fo und useful for extracting acetic acid from aqueo us mixtures with hydrochloric acid [28]. and several other stuclies indicate thai oxygen attached to phosphorus increasingly positive emf value would indicaLe a is sometimes a very effective electron donor and that some acids of pbosphorus gradual increase in the pH of the solution. I n the show a tendellcy to form aggregates in benzene [I(a) (1'.257),29 to 32]. case of diphenyl phosphate, the initial acidity­ -300 as registered by the emf reading- was decidedly reduced when the solvent was changed from nearly pure methanol to predominantly water. Hardly any initial effect on the acidity was observable on -400 changing from methanol to 17: 3 benzene-methanol.

> In the case' of benzoic acid, changing the solvent Eo from nearly pure methanol to g : 1 water-methanol E 2 :2 '" -500 again produced a striking effect on the initial acidity, this time a decided increase in the acidity. As with diphenyl phosphate, comparatively Ii ttle. change in the initial acidity resulted from changIl1 g from Di -o-tolyl­ Guanidine -600 guanidine Acetate me thanol to 17: 3 benzene-me thanol.

!6 '1'he titrant was prepared by diluting 1.I5]VI sodium methylate in methanol to 0.115 ]v[ in a volumetric flask, first adding any methanol and then dllutmg to o 5 10 0 5 10 volume with eitber benzene or water. 'I'he samples of 0.1 M aCid m methanol 0 .119 M Acid, ml were transferred to the titration beakers by means of a volumetric pipet; the added solvents, totalling 45 ml in volume hefore mixing, were. measured in gradu~ted FIGURE 2. Compamtive behavior of (1) diphenyl phosphate cylinders. The solvents containing tbe sample and the titrant were thus Slmllar, although not identical, in composition. A little water was. of course, mtroduced and (2) perchloric acid in the potentiometric titration of di-o­ into each solution witb tbe methanol, and a minute quanttty could also enter tolylguanidine and of guanidine acetate in glacial acetic acid. througb tbe sleeve· type connection of tbe calomel electrode with tbe soluti?n. 'L' he quantities of water thus introduced were adequate for proper functlonmg Electrodes: glass, saturated calomel. Initial solutions contained 0.001 mole of of tbe glass electrode. The solutions tbat. yielded curves 5 had such a blgh water tbe base in 30 m!. content as to be practically the same as aqueous solutions. 315 ", 300 . Diph.nyl Phosphot. Plus 300 200 Benzoic Acid

100

200 0

-100 > E " 100 300 Diph.nyl Phosphot. > 2 E., . E 200 E " 0 , 100 , " 0 4 5: I " -100 I " -100 ,. .... " ". "-::',,, -200 ___ ;..--,.r~:::: . .. .. 1m! -200 , / I 0.1 15 M Sodium M.thylat. , m!

FIGURE 3. Effect of composition of the solvent in the poten· tiometric titration of ben zoic acid and of diphenyl phosphate o 3 4 5 6 7 8 9 10 with sodium methylate. 0 . 115 M Sodium M.thylot., m! Electrodes: glass, saturated calomel. Initial solution: 50 ml of O.Gl -M acid. Composition of solvent: (1) benzene-methanol, 17:3 (by volume); (2) benzene­ FIGU RE 4. Effect of composition of the solvent in the poten­ methanol, 4:1; (3) benzene-metbanol, 1:1; (4) metbanol containing 0.5 weight­ tiometric titration of a mixture of diphenyl phosphate and percent of water; (5) water-methanol, 9:1. Curves 2 to 5 have each been sbifted ben zoic acid with sodium methylate. 1 ml to the right of the preceding curve. E lectrodes: glass, satnrated calomel. Initial volume: 50 ml. Initial concen­ tration of each acid : O.OI-M. Com position of solvent: (1) benzene·methanol, 4:1 (by volume); (2) methanol containing 0.5 weight-percent of water; (3) water­ Figure 4 shows the results of titrating a mixture of methanol,9:1. equivalent quantities of diphenyl phosphate and benzoic acid in three different solventsY The indicators, 5-ml samples of O.l-M diphenyl phos­ buffering effect of benzoic acid on the end-point for phate were repeatedly titrated potentiometrically the titration of diphenyl phosphate in water (curve with sodium methylate, each time with the addition 3) was absent in the titrations in methanol or ben­ of a different indicator dye, and the approximate zene-methanol, thus malting possible a sharp deter­ emfs at which the color changes of the various indi­ mination of the end-point for each acid. The results cators occurred were noted. For the purposes of support Hantzsch 's assertion that water "levels" the experiment, it was satisfactory to regard the the strengths of acids, malting strong acids weaker indicator end-point as the emf midway between the and weak acids stronger [33). An "unleveling" of emf values obtained immediately before and after the strengths of two such acids as benzoic acid and addition of the drop of allmli that produced the diphenyl phosphate can be brough t about by chang­ most striking color change.ls The results thus ing to a nonaqueous solvent like methanol or benzene­ obtained with 17 different indicator dyes are shown methanol. While benzoic acid is thereby reduced in figure 5. It will be noted that three of the symbols in apparent strength, the reduction is not suffi cient are enclosed in parentheses (a- NR, p- EO , MR); to impair seriously the accuracy of the titration. the parentheses are used here to signify that these A few exampJ es of a related character have been basic azo dyes did not undergo a sharp color change reported previously by others [34, 351. at the end-point. For all of the other indicators, the color ch ange at the end-point was very sharp. h. Titrations in 4 :1 Benzene-Methanol Using Indicators In the cases of the sulfonephthaleins (BPB, BCG, BOP, PR, BTB, CR, OP, TB) and phenolphthalein, Benzene-methanol, 4: 1 by volume, is currently there was a roughly linear relationship between the popular as a solvent for nonaqueous titrations of emf values in 4: 1 benzene-methanol at our arbi­ carboxylic acids [36), particularly if the acid is one trarily set end-points and the mid-points of ~he pH of high molecular weight. As was shown in figure 4 ranges as commonly given for aqueous solutlOns of (curve 1), when an equimolar mixture of benzoic these indicators. Nile blue A and Nile blue 2B did acid and diphenyl phosphate in this solvent was not fi t into this linear relationship- their color titrated potentiometrically with sodium methylate, change in benzene-methanol occurred in too nega­ th e separate end-points could be located satisfactorily. tive an emf region. In practice, it would often be more convenient to 18 It sbould be m ade clear tbat the end-points thus arbitrarily established are locate both such end-points by the use of indicators. only approximate. For t be most part, only one titration of a kind was performed. As a preliminary step in finding a suitable pair of Indicators were dissol ved in metbanol, benzene, or benzene-methanol; the con­ centrations of t be indicator solutions ranged from 0.1 percent to 1.0 percent. The amounts added (1 to 5 drops) were sucb tbat one drop of sodium methylate 17 Note that distances along the vertical scale are twice as great as in figure 3. produced a vivid color change at the end-point. 316 magcnta E ch anges color just a little too late (and

PP therefore too gradually). The 5 topmo t indicator 300 should be u able for the second end-point. The one PoNS that we liked best was p-naphtholbenzein, the indicator that was adopted for an important ASTM 200 procedure for determining acidity in petroleum NS-A products dissolved in benzene-isopropanoll38). The NS-29 .~ - .. -. (MR) com bination of phenyl diethyl Nile blue and p­ naphtholbenzein was very satisfactory for titrating 100 the mixture of diphenyl phosphate and benzoic > E SPM-E acid ; the first color change was greenish blue to E violet pink and the second color change, salmon Q) 0 PEoNS pink to violet. It is preferable to add p-naphthol­ benzein after reaching the first end-point; both indicators can be added at the start, but the first end-point is th en a little less striking. The same - 10 0 combination of indicators should, of course, be usable for other pairs of acids comparable in strength to diphenyl phosphate and benzoic acid, and the -200 method used fo r selecting suitable indicators is of still wider application.

o 2 3 4 5 6 7 8 9 10 4.4 Comparative Titrations of Diphenylguanidine, 0 . 115 M Sodium Methylot., ml Triethylamine, and Tris(hydroxymethyl) amino­ methane With Diphenyl Phosphate in Solvents FIGURE 5. Titration of (1) diphenyl phosphate, (2) benzoic of Various Compositions acid, and (3) a mixture of diphenyl phosphate and benzoic acid with sodium methylate in 4:1 benzene-methanol. Initial vol ume, 50 ml; initial concentration of each acid, 0.01 M . Electrodes: 1,3-Diphenylguanidine, triethylamine, and tris­ glass, saturated calomel. Indicators: a·NR, a-naphthyl red; p-EO, p-ethoxy (hydroxymethyl)aminomethane were titrated with chrysoidin; PE·NB, phenyl diethyl Nile hlue; BPM-E , bromophthalein magenta E; BPB, bromophenol blue; BOG, hromocresol green; MR, metbyl red; NB-2B, diphenyl phosphate in the arne series of solvents Nile blue 2B; BOP, bromocresol purple; NB-A, Nile blue A; PR, pbenol red; BTB, bromothymol blue; p-NB, p-naphtholbenzei n; OR, o·cresol red; OP, as those listed under figure 3 and also in 1: 1 water­ m-cresol purple; PP, I)henolphtbalein; TB, thymol blue. methanol. The titration curves are shown in figure 6. The conditions were closely similar to the ones used I t is clear from figure 5 that all of the indicators in the titrations of diphenyl phosphate and benzoic tried, except a-naphthyl red, p-ethoxychrysoidin, acid. 20 The values obtained in titrating triethylamine and methyl red, should be usable for titrating were a little low and not closely reproducible, diphenyl phosphate under the conditions chosen. probably indicating some volatilization of the Bromophthalein magenta E is one of the best. amine; the lo ss seemed to be least in the solutions p-Naphtholbenzein and the three Nile blu es are containing the highest proportion of methanol. also excellent. The relative vertical po sition of the These bases were selected as representing a large ile blues is that which would be predicted from group of organic bases having aqueous dissociation the differences in structure (see section 2.4.c) . constants, expressed as pKb' in the approximate Some of the sulfonephthaleins (PR, OR, OP, TB) range 3 to 6, 21 The tendency for triethylamine .to showed a dual color change, as in water; however, volatilize permits only a tentative compari on of It the color change in the acid region was very gradual. behavior with that of the other two ba es. However, Other sulfonephthaleins (BPB, BOG, BOP, BTB) the curves suggest that triethylamine is more aff ected were practically colorless at first, then gradually in strength than the other two bases when the water became a bright yellow as sodium methylate was content of the solvent is progressively reduced and added, eventually undergoing the characteristic the benzene con ten t is increased. 'There are also vivid color change at the end-point.19 indications that diphenylguanidine is stronger than Two additional curves (lightly dashed) are shown triethylamine in methanol and benzene, although in figure 5- the titration curve for benzoic acid weaker in water. Such a conclusion would be in alone and the titration curve for an equimolar harmony with 1'e ults obtained in spectrophotometric mixture of diphenyl phosphate and benzoic acid. studies of the comparative reactivities of triethyla­ For the mixture of the two acids, as well as for mine and diphenylguanidine with bromophthalein benzoic acid alone, the choice of indicators is greatly magenta E in benzene [1, 39]. narrowed. Phen yl diethyl Nile blue is the only possible indicator in the group studied which can '" In tbe titration of diphenylguanidine a solu t ion O.OI-.i\i in 9:1 water-methanol could not be prepared. 'rhe sample was dissolved in 100 Inl of solvent (Instead oC be used for the first end-point. Bromophthalein the usual 50 ml), with t he aid of about 4 ml of t he standard diphenyl phospbate solution and the titration was then continued from t hat point (curve 6). l' Analogous changes iu color of these dyes in benzene were reported earlier, 21 Val~ es of pK, for triethylamine, dipbenylguanidine, and tris(hydroxy­ and the probable structllral cbanges involved were discussed [37] . methyl)aminomethane are 3.3, 4.0, and 5.9, respectively [I, 4].

~40407-55--2 317 can easily be located by adding a minute amount of an \ 200 " \ indicator dye, which likewise forms a salt according ---'::~:~.:..~~:..:.-::..- .. .. to eq (1). The dye can, of course, be either an ...... , ...... 10 0 , amine or an acid, but it must undergo a striking \ \ \ I i \ change in color when converted to a salt. o 1\ 2. 3: 1 ! \ Vorliinder [42] was the first to perform titrations \ \ I -100 Tris(hydroxymethyil-" ',_ " involving the chemical reaction shown as eq (1); he om inomethone .... _- - . _' ...... referred to this type of titration as "aminometry" or "alkalimetry without hydroxyl ions." 23 In Europe 300 and Australia aminometTY has been applied suc­ cessfully to a large variety of organic bases, and the 200 m ethod would probably find more widespread appli­ cation if acids of suitable strength and purity were > 100 E_ rea,dily obtainable.24 E 0 .. b. Titrations of Diphenylguanidine and Di-o-tolylguanidine -100 With Diphenyl Phosphate in Benzene Few solid organic bases of strengths comparable with those of the aliphatic amines ( pKb~3 to 5) are 300 readily available. ImpOTtant among the few are 200 1,3-diphenylguanidine, C 6H 5NH (C = NH) NHC6H 5, 1 2 3 100 and the closely related compound, 1,3-di-o-tolylguani­ I cline. Previously benzene solutions of these bases 0 51 6: have bcen titrated satisfactorily with d-camphor­ sulfonic acid, picric acid, and trichloroacetic acid in -100 \ ~ """ ':.:" '" this laboratory [1]. As shown in tables 4 and 5, the two bases can also be titrated successfully with 0 .115 M Oiphenyl Phosphote, ml diphenyl phosphate in benzene. Table 4 presents two titrations of 0.05-M diphenyl­ FIGL- RE 6. Effect of composition of the solvent in the poten­ guanidine with 0.05-1l1 diphenyl phosphate with tiometTic lilmlion of Iris (hydToxymethyl) amino methane, 1,3- diphenylguanidine, and triethylamine with diphenyl phos­ bromo phthalein magenta E as the indicator. In phate an acid solution the indicator is yellow; the color in a basic medium is magenta to blue, depending Electrodes: glass, saturated calomel. Initial solution: 50 ml of O.Ol-M base (for exceptioos, see text). Oomposition of solvent: (1) benzene-methanol, 17:3 on the concentration and on the solvent [1 , 39]. (by volume); (2) benzene-methanol, 4:1 ; (3) benzene-methanol, 1:1; (4) methanol containing 0.5 weight-percent of water; (5) water-methanol, 1:1; (6) water-meth­ Experiments 1 and 2 were performed under slightly aool,9:1. Oun'es 2 to 6 have each been s],ifted 1 illl to the rigbt of the preceding curve. different conditions, as indicated in the footnotes. In experiment 1 the solu tion gradually increased in its content of diphenylguanidinium diphenyl phos­ 4 .5. Titrations of Diphenylguanidine and Di-o-tolyl­ phate, becoming about O.Ol-M at the end of the guanidine With Diphenyl Phosphate in Benzene, titration; the concentration of bromophthalein ma­ Using Indicators genta E was about 60 percent as great at the end of a . General Discussion the titration as at the beginning. Experiment 2 was performed in such a way as to maintain the Acid-base titrations using only benzene as the content of diphenylguanidinium diphenyl phosphate solvent differ from the titrations discussed in preced­ approximately constant (O.O.l-M) at each end-point; ing sections in that ions play a very minor role. the concentration of bromophthalein magenta E The reaction in benZ('lle or other aprotic solvents was also the same at each end-point. consists essentially of the combination of the base The change from yellow to magenta was always with the acid to form a salt. An illustration is the sharp. After this end-point had been reached, suf­ combination of an amine with a monobasic, non­ ficient acid was added to restore the yellow color. associated acid, as follows: This second end-point ought to be a little more ac­ curate than the magenta end-point because to R 3N + HA p R 3NH+ .. A- . (1) obtain the magenta color sufficient diphenylguani­ (amine) (acid) (salt) dine must be added to combine with both diphenyl phosphate and bromophthalein magenta E. The As indicated in the equation, the major part of the data do not indicate any advantage in performing salt is not separated into ions when the solvent is 23 A review of the literature and additional examples from " 'ork i.n thi.s labora­ benzene or a like solvent; instead, the salt behaves tory were given in ll]. more like a highly polar molecule.22 The end-point ,. For successful titration tbe reaction (eq 1) must, of course, proceed to comple­ tion. Benzoic and acetic acids, for example, are much too weak to be used for such titrations. VorUinder nsed hydrogen chloride, dissolved in benzene,ligroin, 22 It sbould be noted, bowever, tbat in spite of the minute concentration of or chloroform, and had to restandardize it daily. p·'l'oluenesu!fonic acid is quite free io ns in sucb a solution, LaMer and Downes successfully performed certain strong but the commercial monohydrate mnst be debydrated by individnal aeid·base tit rations pot.entiometrically and conductimetricall y in benzene [40], workers and is not of higb pnrity. Picric acid bas found some use, but is not and M aryott performed some additional acid-base titrations conductimetrically as strong as would be desirable; other objectionable features are t he yellow color in benzene [411 . of its salts and the hazards in handling the acid and its salts. 318 TABLE 4. Titrations of 0.05-1\11 di phenylguanidin e with T A BLE 5. Titrations oj 0.02- J1I ditolylguanidine with O.OiJ- JIf 0.05-M diphenyl phosphate in benzen e, using bromo­ diphenyl phosphate in benzene, using dijJel'e nt indicators phthalein magenta E

Experiment 1. bromophthalein magenta E" I. ---'------.-----'------'------,---=---- Usin~ Acid Strength of base in terms of acid Acid SL rength of base in ?vrolariLy Base terms of acid Base of sali [---,----[ Yellow I M agenta Yello\l' I l\Iagenta Yellow Magenta Yell ow 1 j'vlage nta ---~--~------1------' Experiment 1a ml ml rnl 1---· 1. 02 1.00 1.01 I. 01 0 0.990 ml ml rnl 2.03 2.01 2.02 I. 00, .995 1.11 1.09 1.08 0.0016 I. 02s 1.00, 3.02 3.01 3.02 1. 000 .99, 2.03 2.01 2.02 .0029 1.00, .99, 4.03 4.01 4.02 1. 00, .99s 3.03 3.00 3.01 .0042 1.00, .99, 5.03 5.01 5.01 1. 001 1. 000 4.04 4.01 4.02 .0053 1.00, .99, 6.03 6.00 6.01 1. 003 .99s 5.06 5.04 5.04 .0063 1.00, 1.000 7.02 7.00 7.01 1. 00, .99, 6.03 6.01 6.02 .0071 1.00, .99, 8.06 8. O~ 8.05 1. 00, . 99, 7.07 7.0.> 7.06 .008 1.00, .99, 9.0-1 9.02 9.02 1.00, 1.000 8.03 8. 00 8.03 .0087 1.000 .996 9.99 9.97 9.98 1. 00, .99, 9 .02 9.00 9.02 .0094 1.000 .99s Ayerage 9 .99 9.96 9.9i .01 1.00, .99, ------I 1. 003 .99, ------A voragc. ______1------I 1.00, I ~ Experiment 2. Using tamarack green (base)b --- Experiment 2 b Strength of base in Base terms of acid Acid I 1.03 1.01 1. 03 0.01 1.000 0.98, 2.02 2.00 2.02 .01 1.000 .990 Yellow Green Yello\\' I Green 3.04 3.02 3.05 .01 .990 I I 4.04 4.02 4.0tl .01 ..99, 99, I .990 , 5.07 5.05 5.08 . 01 .99, .99, ml ml ml 6.03 6.01 6.03 .01 1.000 .99, 1. 01 1. 00 0.99 1. 01 0 1. 020 7.02 i.OO 7.01 .01 1.00, .99, 2.05 2.07 I 2.06 .990 .995 8.03 8.00 8.03 .01 1.000 .996 3.05 3.07 3.06 . 993 .99, 9.02 9.00 9.02 .01 1.000 .99s 4.01 4.02 4. Ol . 99. 1. 000 10.00 9.98 9.99 .01 1.00, .99, 5.01 5.05 5. O~ .99, 1.00, 7.01 7.01 i.OO 1. 000 1.00, 8.52 8.51 8.50 I. 00, J. 00, 9.89 9.90 9.88 .99, I 1. 00, I I --- Ayerage __ __ n. re hirty ml of benzene and 2 drops of O.l-percent hromophLhaiein magen ta E - - ._------.99, 1. 00, in benzene were ad(led to the titration Oask at the beginning of the titration. ---I Acid and base were acid(l(i to the so!w'nt in approximately 1-1111 incrcllIents and ---- cnd-points werG dClermilwd for the sllccrssi\'C' incrcmrl1ts. "I'h(' approximate l~xpcrill1('nt 3. Using Yicioria blue B (base)' calcu lated molarity of diphenylguanidiniulll diphenyl phosphate at the end­ point is given in the fourth co lumn. Strength of base in tel illS of b rJ' hirty m1 of benzene and 2 drops of O.l-percent hromophthalein magenta E Base aei(l in benzene were mixed, and 3 ml of this solutio n was added from a buret at the beginning of the experiment and before each succeed ing end-point was deter­ mined. 'J'he concentration of diphenylguanidinium djphenyl phosphate was therefore O.O I-.M at each end-poin t .

ml Tnl ml rn' I. 00 1. 02 1.00 0.980 1. 000 the titration bv the more cumbersome procedure 2.00 2. O~ 2.02 2.01 .985 0.990 .995 used m 9xperiment 2. ~\.n a lo gous results were 3.00 3.02 3.01 . 993 .99, 4.02 4.06 4.03 4.01 .990 .99s 1.00, obtained m titrating O.02-M di-o-tolylguanidine 5. 00 5.03 5.00 .99, 1.000 6.01 G.O-l .995 with O.02-Af diphenyl phosphate by the same two 7.01 7.03 7.01 .99, 1. 000 proced ures. 8.03 8.06 8.0-1 8.02 .996 1. 00, 9.05 9.08 9.06 9.04 . 99, 1. 00, Table 5 presents examples of titrations of O.02-M 9.97 10.00 9.99 9.97 .99, 1. 000 di-o-tolylguanidine with O.02-M clipheny l phosphate Average __ in benzene using three different indicator d)Tes- an acidic indicator, bromophthalein magen ta E, and a ~ I 'hhty ml of benzene and 2 drops of O.I-percent bromophthaleLn ma.genta in bellzene were addcd at the beginning of the titratioll. t,wo basic indicators, tamarack green (base) and b One drop of O.l-pereent tamarack green in benzene was adeled at the start. Victoria blue B (base) . The tit.ration using bro­ e rr hirty ml of benzene and 2 drops of Ill-percent Victoria. blue B in benzene mophthalein magenta E was analogous to experi­ were added at the beginning. ment 1 of table 4 and_ needs no flll'ther discussion. In addition to these t itrat ions O.02- J'£ cliphenyl­ In the use of tamarack green (base) the initially guanidine and O.Ol - M di-o-tolylguanidine were yellow color becomes green when sufficient diphenyl titrated in benzene with diphenyl phosphate of phosphate has been added to combine with both equivalent concentration. The precision and ac­ ditolylguanidine and the indicator; the yellow color curacy were excellent, especially in view of the small was restored by adding a little ditolylguanidine. volumes and low concentrat ions used. Diphenyl phosphate should be satisfactory for t itrating other In titrations wi th Victoria blue B (base) as the incli­ moderately strong bases in benzene and other ('ator, three di stinct colors- orange, blue, and aprotic solvents. It is probably a weaker acid than nn intermediate gray- were observed. There is p-toluenesulfonic acid and other sulfonic acids, but scarcely any difference in the results obtained with is superior to the sulfonic acids in being more soluble the till'ce indicators, although on the whole we in aprotic solvents and more ea il~T obtainecl in a liked bromophthalein magenta E the best. pure anhydrous state.

319 4.6. Concluding Remarks [11] 1. M. Kolthoff, Acid-base indicators (The Macmillan Company, New York, N. Y ., 1937; 4th German Edi­ tion, 1932, t ranslated by C. Rosen blum). In the preceding sections it has been shown that [12] M. L. Crossley, R. J . Turner, C. :\1. Hofmann, P . F . accurate titrations with diphenyl phosphate can be Dreisbach, and R. P . Parker, J . Am. Chern. Soc. 74, performed in solvents as different as benzene and 578 (1952). [13] (a) H . Hlasiwetz and A. Grabowski, Liebigs Ann. 139, water, as well as in solvents of intermediate char­ 95 (1866); Zeitschrift fUr Chemie 9, 390 (1866). acter. In view of its high acidic strength, its rela­ (b) O. R embold, Zeitschrift fUr Chemie 9, 651 (1866). tively high equivalent weight (2 50.19), the ease with [14] L. Glutz, Liebigs Ann. 143, 181 (1867). which it can be obtained and maintained as anhy­ [15] G. Jacobsen, Bu . deut. chern. Ges. 8, 1519 (1875). drous crystals, and its apparent; resistance to hydrol­ [16] M. Rapp, Liebigs Ann. 224, 156 (1884) . [17] W. Autenrieth, Bel'. deut. chern . Ges. 30,2369 (1897). ysis and alcoholysis at ordinary temperatures, it [18] J . M. A. Hoefiake, R ec. trav. chim. 36, 24 (1916). promises to be a useful material for acid-base titra­ [19] A. Bernton, Bel'. deut. chern. Ges. 55B, 3361 (1922). tions in both aqueous and nonaqueous solutions. [20] F. Zetzsche and M. Nachmann, H elv. Chim. Acta 9, 705 (1926). More rigorous tests must be applied before it can be [21] S. Forss man, Biochem. Z. 263, 135 (1933), t hrough Chem. recommended as an acidimetric standard for work Abstracts 27, 4726 (1933) and Brit. Chem. Abstracts of the highest accuracy. 1933A,897. The properties of diphenyl phosphate suggest its [22] C. Morton, Quart. J. Pharm. Pharmacol. 3, 438 (1930) . [23] W. D. Kumler and J . J. Eiler, J. Am. Chem. Soc. 65 , use as an acid catalyst. Furthermore, it offers a 2355 (1943). means for extending knowledge of the behavior of [24] R . H . A. Plimmer and VV. J . N. Burch, J. Chem. Soc. acids. The great majority of the studies of acidic (London) 1929, 279. behavior made so far have been performed with [25] M . 1. Smith and E. F. Stohlman, U. S. Public H ealth Repts. 48, 734 (1933). aliphatic or aromatic carboxylic acids. However, [26] F. L. Breusch and H . Keskin, Rev. fac. sci. univ. Istanbul studies of the sort illustrated in figures 3 and 4 show 7A, 182 (1942). that generalizations reached through observations of [2 7] E. Blumenthal and J. B. :\1. Herbert, Trans. Faraday the behavior of carboxylic acids in various solvents Soc. 41, 611 (1945) . [28] (a) H . A. Pagel and F. W. McLafferty, Anal. Chem. 20, cannot safely be applied to all types of acids. The 272 (1948); investigations with diphen:vl phosphate are sugges­ (b) H. A. Pagel, P . E. Toren, and F. W . McLafferty, tive of interesting work still to be done with acids Anal. Chem. 21, 1150 (1949). which as yet are not widely known, such as bisaryl­ [29] c. S. Marvel, M . J . Copley, and E. Ginsberg, J . Am. Chem. Soc. 62, 3109 (1940) . phosphonic acids. [30] (a) W . Gordy and S. C. Stanford, J. Chern. Phys. 8, 170 (1940); 9, 204 (1941). (b) W. Gordy, J. Chem. Phys. 9,215 (1941). The authors are indebted to James Stewart for (c) S. C. Stanford and W. Gordy, J . Am. Chem. Soc. 63, measuring the benzene content of cyclohexane 1094 (1941). [31] L. F. Audrieth and R. Steinman, J. Am. Chem. Soc. 63, (table 1), and to the following companies for donating 2115 (1941). materials: American Cyanamid Company, for 5- [32] (a) G. M. Kosolapoff and J. F . YfcCullough, J. Am. phenyl-9-diethyl Nile blue; Dow Chemical Company, Chem. Soc. 73, 5392 (1951), and other papers by for diphenyl phosphate; Heyden Chemical Corpo­ Kosolapoff and coworkers. (b) G. M. Kosolapoff, Organophosphorus compounds ration, for o-chlorobenzotrichloride. (John Wiley and Sons, Inc., New York, N. Y., 1950). [33] A. Hantzsch, Z. E lektrochem. 29,221 (1923). 5. References [34] (a) L. Lykken, P . Porter, H . D. Ruliffson, and F. D . Tuemmler, Ind. Eng. Chem., Anal. Ed. 16,219 (1944). (b) L. Lykken, Lubrication Eng. 2, 23, 64 (1946). [1] (a) M. M . Davis and P. J. Schuhmann, J. Research [35] S. Palit, Ind. Eng. Chem., Anal. Ed. 18,246 (1946). NBS 39, 221 (1947) RP1825. [36] (a) J . S. Fritz and N. M. Lisicki, Anal. Chern. 23, 589 (b) NBS Tech. News Bull. 36, 168 (1952) . (1951) . [2] Dow Special Products, p . 88 (The Dow Chemical (b) J. S. Fritz, Acid-base titrations in nonaqueous Company, Midland, Michigan, 1st edition, March solvents (G. Frederick Smith Chemical Company, 1939). See also the 1950 edition of this catalog. Columbus, Ohio, 1952). [3] C. J. P enther and F. B. Rolfson, Ind. Eng. Chern., Anal. [37] M. M. Davis, P . J . Schuhmann, and M . E. Lovelace, J . Ed. 15, 337 (1943). Research NBS 41, 27 (1948) RP1900. [4] R. G. Bates and G. D . Pinching, J. Research NBS 43, [38] H . F erguson, Anal. Chern. 22, 289 (1950). 519 (1949) RP2043. [39] M. M . Davis and H . B. 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