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Volumetric analysis of metals in glacial acetic acid. Casey, Allan Terence 1953

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VOLUMETRIC ANALYSIS OF METALS IN GLACIAL ACETIC ACID by ALLAN TERENCE CASEY A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of Chemistry We accept t h i s thesis as conforming to the standard required from candidates for the degree of MASTER OF SCIENCE. Members of the Department of Chemistry. THE UNIVERSITY OF BRITISH COLUMBIA A p r i l , 1953 ABSTRACT Recent work has demonstrated that certain metal s a l t s of inorganic acids may e a s i l y be t i t r a t e d with perchloric acid i n g l a c i a l acetic acid. To test the accuracy of the method, a series of t i t r a t i o n s was run on the acetates of twenty-three representative metals. The r e s u l t i n g curves show that accurate reproducible end points may be obtained f o r the acetates of the more ele c t r o p o s i t i v e metals, while those of aluminum, chromium and iron y i e l d anomalous r e s u l t s . An explanation f o r t h i s i s offered. The acetates of t i n , bismuth, antimony, uranium and copper are too weak to be determined by t h i s method. Metals, oxides, or other more d i f f i c u l t l y soluble com-pounds may be conveniently introduced into acetic acid solution by f i r s t d i s s o l v i n g i n n i t r i c acid (or i n certain cases aqua regi a ) , followed by the addition of s u f f i c i e n t acetic anhydride to react with a l l the water and decompose a l l the n i t r a t e present. Quantitative a n a l y t i c a l separations may be effected using such common reagents as oxa l i c , n i t r i c , and t a r t a r i c acid, hydrogen iodide and hydrogen sulphide. Hhey were used i n analysing eleven metal mixtures, with accuracy comparable to the more lengthy gravimetric process often required i n aqueous solution. ACKNOWLEDGEMENTS I wish to express my sincere apprecia-t i o n to my research d i r e c t o r , Dr. K. Starke, f o r his invaluable aid and encouragement during the course of t h i s i n v e s t i g a t i o n . I also wish to thank the International Nickel Company fo r f i n a n c i a l support of t h i s work, and Major Aluminum Products of Vancouver f o r a sample of aluminum-magnesium a l l o y ; TABLE OF CONTENTS Page I INTRODUCTION 1 II EXPERIMENTAL A. Materials and Apparatus Solvents 12 Preparation of 'i'itrant 13 Indicator. 14 Potentiometer........ 14 Standardization of the Ti t r a n t 15 B. T i t r a t i o n of Metal Salts Sodium Acetate 16 Acetates of Potassium, Lithium, 16 Barium, Strontium, and Ammonium Magnesium Acetate 17 Acetates of Calcium, Manganese, and Cadmium 17 Acetates of Cobalt and Nickel 17 zinc Acetate IS Acetates of Lead and S i l v e r IB Acetates of Iron, Chromium, and Aluminum IB Acetates of Tin, Antimony, Copper, Bismuth, and Uranium 1<3 T i t r a t i o n of Chloroacetates 19 T i t r a t i o n of Chlorides 19 C. Separation and Analysis of Metal Acetate Mixtures S o l u b i l i t i e s of Metal Salts i n Acetic Acid 20 Separations Effected Zinc and Magnesium... 22 Zinc and Copper 23 Barium and Calcium 23 S i l v e r and Aluminum *. 23 Lead and Tin 24 Magnesium and Antimony; Magnesium and Lead, Magnesium and Copper 24 Nickel and Chromium 25 Cadmium and Magnesium 25 Tin and Magnesium. 25 D. Introduction of Metals Into Acetic Acid Solution 26 TABLE OF CONTENTS (continued) Page I I I RESULTS AND DISCUSSION A. Determination of Potentiometric End Point 23 B. T i t r a t i o n of Metal Salts '30 C. S o l u b i l i t i e s i n Acetic Acid..... 33 D. Separations 36 E. Introduction of Metals into Acetic Acid Solution. 3& IV CONCLUSIONS 42 V BIBLIOGRAPHY 44 i v TABLE OF vCONTENTS FOR TABLES Page Table I T i t r a t i o n of Metal Salts. 3 1 Table II P r e c i p i t a t i o n Reactions i n Acetic Acid... 3 4 Table I I I Results of S o l u b i l i t y Tests With Sulphates 3 7 Table IV Results of T i t r a t i n g M e t a l l i c Mixtures... 3 9 Table V Reactivity of Amine Acetates, Acetic Acid, and the Chloroacetic Acids Toward Metals 4 0 Table VI Acids Tested Which Attack the Metals Considered With S u f f i c i e n t Rapidity to be of P r a c t i c a l Use 4 1 TABLE OF CONTENTS FOR FIGURES Figure I Figure I I A Figure I I B Figure I I I A Figure I I I B Fi g u r e IV A Figur e IV B Figur e V A Figur e V B Figure V I A Figure V I B Figure V I I A Figure V I I B Figure V I I I A Figure V I I I B Page R e l a t i v e Basic Strengths of C e r t a i n Acid Ions, Solvents and Organic Bases.. 10 P o t e n t i o m e t r i c T i t r a t i o n of Sodium Acetate w i t h P e r c h l o r i c A c i d 4# P r e c i s e Determination of the End Point f o r Figure I I A 49 P o t e n t i o m e t r i c T i t r a t i o n of Barium Acetate w i t h P e r c h l o r i c Acid 50 P r e c i s e Determination of the End Point f o r Figure I I I A..... 51 P o t e n t i o m e t r i c T i t r a t i o n of L i t h i u m and Strontium Acetates with P e r c h l o r i c Acid 52 P r e c i s e Determination of the End Poin t f o r Figure IV A 53 P o t e n t i o m e t r i c T i t r a t i o n of Magnesium Acetate w i t h P e r c h l o r i c t A c i d 54 P r e c i s e Determination of the End Point f o r Figure V A 55 P o t e n t i o m e t r i c T i t r a t i o n of Calcium Acetate w i t h P e r c h l o r i c Acid 56 P r e c i s e Determination of the End Point f o r Figure V I A 57 P o t e n t i o m e t r i c T i t r a t i o n of Cadmium and Manganese Acetates with P e r c h l o r i c Acid 53 P r e c i s e Determination of the End Poin t f o r Figure V I I A 59 P o t e n t i o m e t r i c T i t r a t i o n of Cobalt Acetate w i t h P e r c h l o r i c A c i d 60 P r e c i s e Determination of the End Point f o r Figure V I I I A 61 v i TABLE OF CONTENTS FOR FIGURES (Continuted) Page Figure IX A Potentiometric T i t r a t i o n of Lead Acetate with Perchloric Acid.. 62 Figure IX B Precise Determination of the End Point f o r Figure IX A 63 Figure X A Potentiometric T i t r a t i o n of Aluminum, • Chromium, and Iron (III) Acetates with Perchloric Acid 6 4 Figure X B Precise Determination of the End Points f o r Figure X A 65 Figure XI Curves Obtained from the Attempted Potentiometric T i t r a t i o n of Copper (II) Antimony ( I I I ) , and Tin (II) Acetates-with Perchloric Acid 6 6 Figure XII Curves Obtained from the Attempted Potentiometric T i t r a t i o n -of Bismuth (III) and Uranyl Acetates with Perchloric Acid 67 Figure XIII A Potentiometric T i t r a t i o n of Barium Chloroacetate with Perchloric Acid.... 68 Figure XIII B Precise Determination of the End Point f o r Figure XIII A 6 9 1 I . INTRODUCTION U n t i l r e c e n t l y , w a t e r has been the s o l v e n t employed f o r most a n a l y s e s , b o t h q u a l i t a t i v e and q u a n t i t a t i v e . However, i t i s i m p o s s i b l e t o c a r r y out r e a c t i o n s , such a s n e u t r a l i z a t i o n , i n v o l v i n g a c i d s (or b a s e s ) a p p r e c i a b l y weaker t h a n w a t e r , s i n c e w a t e r , b e i n g a s t r o n g e r a c i d ( b a s e ) , w i l l r e a c t p r e f -e r e n t i a l l y w i t h t h e a d d e d r e a g e n t . F o r t h i s r e a s o n , i n t h e l a s t t h r e e d e c a d e s , n o n - a q u e o u s s o l v e n t s have become i n -c r e a s i n g l y i m p o r t a n t i n a n a l y s i s , p a r t i c u l a r l y i n q u a n t i t a -t i v e d e t e r m i n a t i o n s , s i n c e a c i d - b a s e t i t r a t i o n s i n t h e s e m e d i a p o s s e s s the a c c u r a c y , s p e e d , and s i m p l i c i t y r e q u i r e d f o r good a n a l y t i c a l m e t h o d s . I n o r d e r t o c h a r a c t e r i z e t h e p r o p e r t i e s o f i n t e r e s t i n nonaqueous s o l u t i o n s , i t i s u s e f u l t o r e v i e w t h e p a r t i c u l a r p r o p e r t i e s o f w a t e r w h i c h a r e r e s p o n s i b l e f o r t h e u n i q u e p o s i t i o n i t o c c u p i e s among s o l v e n t s ( 3 4 ) . W a t e r , a v a i l a b l e i n l a r g e q u a n t i t i e s , i s an e x c e l l e n t s o l v e n t f o r most c l a s s e s o f s u b s t a n c e s , b o t h e l e c t r o l y t e s and n o n - e l e c t r o l y t e s . The p u r e s o l v e n t i s a p o o r c o n d u c t o r o f e l e c t r i c i t y , but aqueous s o l u t i o n s o f e l e c t r o l y t e s have h i g h c o n d u c t i v i t i e s . M o r e -o v e r , w a t e r has t h e a b i l i t y t o c o n v e r t p o t e n t i a l e l e c t r o l y t e s i n t o t r u e e l e c t r o l y t e s . Many s u b s t a n c e s , such as s u l p h u r d i o x i d e , • a m m o n i a , hydrogen c h l o r i d e , and t e r t i a r y a m i n e s , w h i c h by t h e m s e l v e s a r e n o n - e l e c t r o l y t e s , r e a c t w i t h w a t e r t o g i v e s o l u t i o n s w h i c h a r e e x c e l l e n t c o n d u c t o r s . N e u t r a -l i z a t i o n r e a c t i o n s between s t r o n g a c i d s y i e l d i n g h y d r o n i u m 2 ions and bases y i e l d i n g hydroxyl ions occur e a s i l y i n water, a process made possible by i t s l i m i t e d s e l f - i o n i z a t i o n . The phenomenon of hydrolysis i s a c h a r a c t e r i s t i c reaction of water. Not only do dissolved s a l t s undergo hydrolysis, but many sub-stances such as acyljchlorides, esters, and s a l t s of very weak acids react chemically with the solvent i n such a way that the o r i g i n a l solutes cannot be recovered d i r e c t l y . The s u p e r i o r i t y of water as a general solvent i s re-f l e c t e d more p r e c i s e l y through i t s exceptional and anomalous physical properties. The high values of i t s melting point, b o i l i n g point, heat of vaporization, and d i e l e c t r i c constant, and i t s low v i s c o s i t y , d i s t i n g u i s h i t as a solvent of the highest general u t i l i t y . However, there are properties of nonaqueous media which indicate that they are, i n c e r t a i n spec i a l i z e d circumstances, superior t o water as a solvent. One of the most s t r i k i n g of these properties i s the high a c i d i t y displayed by solutions of HCIO^, H^SO^, and HCI, f o r example, i n aprotic l i q u i d s such as benzene or chloroform, or i n acid solvents, such as g l a c i a l a c e t i c acid. These compounds are strongly a c i d i c and completely ionized i n aqueous solution, but are even more a c i d i c i n solvents of l e s s e r i o n i z i n g power. Moreover, the acid strengths of HCIO^, H^SO^, and HCI d i f f e r widely in these media, whereas i n water they are indistinguishable f o r equal concentrations. That such phenomena occur may at f i r s t seem surprising, but, as the ensuing discussion shows, they arise as a l o g i c a l consequence of the Bronsted-Lowry concept of acids and bases. The concept of electrolytic dissociation has dominated the f i e l d of acid-base chemistry for over half a century. Until 1926 i t was considered to be the f u l l explanation of acid-base phenomena, although i t was realized that the theory could not be applied to solvents other than water. There were scattered reports of titrations in nonaqueous media (19,20,21^ but i t was not u n t i l Hantsch (30) published his series of investigations that a departure from the prevailing theory became evident. The appearance of the Bronsted-Lowry theory (2, 3, 4, 5, 42, 43) and the classic investigations by Conant, Hall, and Werner (7, 8, 26, 27, 28) on acidity in glacial acetic acid, put the following analysis on a sound experimental and theoretical basis. It has been shown by Bronsted that a logically con- , sistent general measure of acidity i s the hydrogen-ion pot-ential, or the closely related hydrogen-ion activity. It has .been found that the use of acti v i t i e s instead of pot-entials i s convenient, but i t should be noted that this does, not imply any knowledge of the concentration. Hydrogen-ion act i v i t y i s defined as an exponential function of the hy-drogen-ion potential, which represents the work concerned in the reversible removal of hydrogen ions from a given system, or their addition to i t . High activity means loosely bound hydrogen ions; low activity implies strongly bound hydrogen • ions. A reaction proceeds by the transfer of hydrogen ions 4 from higher to lower a c t i v i t y , although hydrogen ions as such do not exist to a measurable extent i n the solution.' Since only the r a t i o s of a c t i v i t i e s may be measured, one must always set up an a r b i t r a r y scale of numerical values. Usually d i f f e r e n t conventions are adopted f o r d i f f e r e n t solutions such that the a c t i v i t y c o e f f i c i e n t approaches unity at i n f i n i t e d i l u t i o n , but t h i s i s not sa t i s f a c t o r y when s o l -utions of the same solute i n various solvents are considered. I f the same scale i s used f o r a set of solvents, then two solutions with the same hydrogen-ion potential ( i . e . the same true a c i d i t y ) w i l l have the same hydrogen-ion a c t i v i t y , no matter what the solvent, but i t also follows that the a c t i v i t y c o e f f i c i e n t w i l l not approach unity f o r i n f i n i t e l y d i l u t e solutions. With t h i s understanding, following Hammett (29) we may consider a mathematical approach to the effect of solvent upon a c i d i t y . Let HX be the acid concerned; S the solvent; C the concentration; a the a c t i v i t y ; f the a c t i v i t y c o e f f i c i -ent; and C a the stoichiometric concentration of HX. Regard-l e s s of the solvent, we have C x - J ~ but i f the solvent i s basic we also have i - K5 C S H * (2) If, as seems probable, Cy+. may be considered negligibly small d= CHX +CKr = C„K -f CSH+ (3) since the reaction between the acid and the solvent i s HX + S ;F=± X" •+ SH+. A combination of 1, 2 and 3 gives 2 TP ( V KHX I T S^HX / Since au+ increases with increasing K y this equation re-n o cords the principle that decreased basicity of the solvent (greater Ky) means greater acidity, although i t means less ionization. In order to see more clearly the effect of changes of the dielectric constant, we may consider two limiting cases of 4. When the second term under the radical sign i s large compared with one, we obtain • & Sx~ This i s the case f o r a weak acid or.weakly basic solvent; the.reaction HX + S ^  ' X" + SH+ i s incomplete and the acidity depends on both and Kj^, i.e. upon the properties of both acid and solvent. On the other hand, when the same term i s small the radical may be expanded and using only the linear term we have *H+ c & Ks SSH* ( 6 ) Here ionization i s large and the acidity depends only on 6 the p r o p e r t i e s of th e . s o l v e n t . This represents the l e v e l l i n g e f f e c t of the s o l v e n t , mentioned by Hantsch and by Conant and H a l l . I t i s now p o s s i b l e t o make a t l e a s t a q u a l i t a t i v e pre-d i c t i o n of the e f f e c t of the d i e l e c t r i c constant from t h e known e f f e c t of t h i s property upon the a c t i v i t y c o e f f i c i e n t s . The a c t i v i t y c o e f f i c i e n t of an i o n i s gre a t e r i n solvents of low than of high d i e l e c t r i c constants, as demonstrated by the extremely low s o l u b i l i t y of t r u e s a l t s i n such media. This i s p r e d i c t e d by the i n t e r i o n i c a t t r a c t i o n theory of e l e c t r o l y t e s . The changes i n a c t i v i t y c o e f f i c i e n t s of uncharged molecular species from so l v e n t to solvent are much smaller than the changes of the i o n i c a c t i v i t y coef-f i c i e n t s . Therefore, i t f o l l o w s that r a t i o s such as ..'HX/..X~ o r . S/\ SH w i l l decrease with decreasing d i e l e c t r i c constant. Now, as an admittedly rough approximation, assume t h a t the change i n t h i s r a t i o between two d i f f e r e n t s o l v e n t s i s independent of the a c i d or base concerned. There i s some reason f o r b e l i e v i n g t h a t s p e c i f i c e f f e c t s upon the r a t i o w i l l be considerably smaller than on the i n d i v i d u a l a c t i v i t i e s , corresponding t o the impression that a c i d s and bases very s o l u b l e i n organic l i q u i d s w i l l have s a l t s which are s o l u b l e i n organic l i q u i d s . ^8H + ^HX I t then f o l l o w s t h a t the r a t i o - — - i n equation 5 w i l l be approximately u n i t y f o r a l l s o l v e n t s ; t h a t the  hydrogen i o n a c t i v i t y of a weakly i o n i z e d a c i d i s independent of the d i e l e c t r i c constant of the medium. It follows from equation 6 that the hydrogen ion a c t i v i t y of a highly ion-ized acid w i l l be greater i n solvents of low than i n s o l - vents of high d i e l e c t r i c constant. For intermediate cases-, a decrease i n the d i e l e c t r i c constant w i l l r e s u l t i n a more or l e s s large increase i n the hydrogen-ion a c t i v i t y . • From the work of Conant and H a l l , and from the above discussion, i t appears that there are three p r i n c i p a l i n -fluences a f f e c t i n g the apparent strength of an acid i n a given solvent. The f i r s t i s the " i n t r i n s i c strength? of the acid, the general tendency of the molecule to l i b e r a t e a proton. The constant expressing t h i s tendency may be re-garded as f i x e d and equal f o r a l l solvents at any given temperature. Next, there i s the basic strength, or proto-p h i l i c tendency, of the solvent, which determines to what extent the true strength of the acid can manifest i t s e l f . F i n a l l y there i s the d i e l e c t r i c constant of the solvent which determines the extent to which the ions, products of the reaction of the acid with the solvent, can e x i s t i n -dependently of each other'. Water i s a basic solvent of high d i e l e c t r i c constant. Therefore, strong acids react com-p l e t e l y with i t , and the products exist as separate ions. HC10 4 + H 20 x H^0+ + ClO^"* The a c i d i t y of a d i l u t e solution i n water can never be very high, however, because the acid molecules a l l react to pro-duce hydronium ions, whose proton a c t i v i t y i s not very great. In a non-basic solvent,, such as acetic acid or benzene, the acid molecules w i l l combine with the solvent to a much le s s extent, so that, the proton a c t i v i t y of t h e i r i n d i v i d u a l molecules may give a high a c i d i t y to the solution, although, because of the low d i e l e c t r i c constant there w i l l be few (unassociated) ions. Although as yet no method exists to compare the a c i d i t i e s of solutes i n d i f f e r e n t solvents i n a completely rigorous manner, i t has become evident that these predictions are f u l f i l l e d . Thus i t may well be possible to prepare a series of solutions containing only moderate con-centrations of the strongest acids i n suitable solvents whose true a c i d i t i e s range well into the m i l l i o n f o l d values of those obtainable i n water solu t i o n at the same concen-t r a t i o n . As indicators can be found which are s e n s i t i v e i n the range of high absolute a c i d i t y , i t should be possible to develop a "pH technique" which w i l l extend to high negative values. One must always remember, however, that the true measure of a c i d i t y i s to be taken as the hydrogen ion a c t i v i t y as measured with a suitable electrode, and not other l e s s precise e f f e c t s . These conclusions may be amplified by an i l l u s t r a t i v e example. Many very weak anhydro bases, which do not form s a l t s i n aqueous solutions, are neutralized by perchloric or sulphuric a c i d i n g l a c i a l acetic acid solution. This i s shown by the t i t r a t i o n curves obtained by Conant and H a l l (7, 2 7 ) . It should be noted that i n the formation of s a l t s from anhydro bases, no water i s eliminated, and so the water a c t i v i t y i s not involved. Their r e s u l t s confirm ex-perimentally that, whether acid strength i s defined i n terms of the r e l a t i v e hydrogen ion a c t i v i t i e s of t h e i r equimolal solutions, or i n terms of salt formation with anhydro bases, perchloric and sulphuric acids are stronger i n acetic a c i d than i n water. At present we have some information-about the behavior of acids i n a wide variety, of solvents, of which ammonia, water, and g l a c i a l a c e tic acid may be considered as representative examples. In each medium, evidence points to the proton being completely solvated. In comparing the three solvents, one i s led to consider that the essential difference between glacial„acetic acid and water i s the same as that between water and ammonia. The a c e t i c acid molecule has l e s s a f f i n -+ i t y f o r a proton than the water molecule; the CH^COOH^ ion gives up i t s extra proton more r e a d i l y than H^0+. The formation of a s a l t from an anhydro base i s the re-verse of the d i s s o c i a t i o n of the a c i d . The anhydro base (B), the anion of the a c i d (X~) oand the solvent molecule (5) may + a l l be regarded as competing f o r the proton H . Therefore the following equations may be written: X" + H + ,==1 HX (7) B + H + BH + . (g) .'S + H + SH + (9) 10. Probably i n most solvents the amount of unsolvated proton is- small, but we may s t i l l formulate the reaction of s a l t formation of B: SH + + B ;==! BH + + S i n terms of the competing reactions 8" and 9 on Page 9 . A very weak base B (corresponding to a very strong acid BH ) w i l l only form stable s a l t s i n a solvent which has con-siderably l e s s a f f i n i t y for the proton than B i t s e l f . Also i t i s obvious that s a l t formation can only occur with an anion (X ) which w i l l release the proton to B. The i n t e r -r e l a t i o n of these three competing reactions i s i l l u s t r a t e d i n F i g . 1 , due to Conant and H a l l ( 7 ) . s <* C 6H 50" NH3 £ RNH2 as o RCOO" C 6H 5NH 2 i -X u. O CCI3COO" H 20 >- NH2C0NH2 z u . CH3C0NH2 14. «c * C7H7SO3' HSO^ - AcOH C6H5NHCOCH3 « 11/ (C 6H 5) 3N AC z C I O 4 Acid Ion Solvent Base X~ S B FIG. r Relative Basic Strengths of Certain Acid Ions Solvents and Organic Bases The position of some common acids, anhydro bases and three solvents are indicated by the formulas. The scale i s ar b i t r a r y ; lack of data s t i l l makes i t impossible to evaluate exact r e l a t i o n s h i p . Acids are strong only i n the solvents l y i n g above them i n the diagram, e.g. perchloric i s strong i n a l l three s o l -vents; sulphuric and hydrochloric i n water and ammonia; and RCOOH only i n ammonia. Anhydro bases are neutralized by one equivalent of any acid which i s stronger than t h e i r nonium t t ions, provided that the action occurs i n a solvent l y i n g below the p o s i t i o n of thedons on the scale. - Thus RCONHg i s neutralized only i n ac e t i c acid and only by perchloric ac i d . I f any solvent i s found i n which triphenylamine may be t i t -rated, i t w i l l l i e beneath acetic acid on the diagram. Since perchloric acid i s very strong i n a c e t i c a c i d , and since the ac i d i c nature of the medium prevents solvoly-s i s of s a l t s of strong acids and weak bases, one may con-clude that acetic a c i d i s an excellent solvent f o r the analysis of substances weakly basic i n nature. The work of Conant and Ha l l provided the basis f o r a l l subsequent i n -vestigations, but i t has only been i n the past f i f t e e n years that s u f f i c i e n t , interest has arisen i n the method to y i e l d p r a c t i c a l dividends. In t h i s period, procedures have been published f o r the determination of pyridine and various pyramidines (4S); amines (1, 22, 23, 44, 61), amino acids and alcohols (3S, 44, 47, 5#, 59, 60); oxazolines (44); 12 quinine and nicotine (31, 35); alkylene oxides (1); sulphon-amides (61); basic nitrogen i n o i l s (71); and choline s a l t s of carboxylic acids (45). More clo s e l y related to the sub-je c t of t h i s paper.was•the determination of the s o l u b i l i t i e s of c e r t a i n inorganic s a l t s i n acetic acid by Davidson and h i s co-workers (9-17, 2 4 , 2 5 ) and by Kolthoff (40), followed by the successful analysis of metal s a l t s of carboxylic acids and of s t r i c t l y inorganic s a l t s by P i f e r and Wollish ( 5 0 , 51 53), Higuchi and Concha (3;2, 33), and Markunas and Riddick (44). Since Davidson has shown (9) that metathetical rea-ctions may proceed as e a s i l y i n a c e t i c acid as i n water, these l a s t r esults indicate that rapid and accurate quanti-, t a t i v e analyses could be c a r r i e d out on one or more metal -compounds provided they could be e a s i l y introduced into acetic acid solution and that methods of separation could be found. It was the object of t h i s i n v e s t i g a t i o n to begin de-veloping such methods. I I . EXPERIMENTAL A. Materials and Apparatus' For every t i t r a t i o n , the sample was dissolved i n an-hydrous acetic acid. The ordinary C P . grade of commercial acid was p u r i f i e d by the method of Kendall and Gross (37), that i s , by adding the calculated amount of acetic anhydride to react with the water present, then d i s t i l l i n g twice using a t h i r t y - i n c h h e l i x type column. The f r a c t i o n b o i l i n g at 117-118°C. was collected; samples melted at l6.5°C. ( l i t . 13 16.6° ( 3 7 ) ) . Two solvents were used for the ti t r a t i n g solution of perchloric acid: glacial acetic acid and 1, 4-dioxane. The , latt e r was reagent quality solvent, purified according to Vogel (63 ) , and d i s t i l l e d twice. The fraction boiling between 100.5°C. and.l01.5°C. was collected; samples had a melting point'of 11,8-11.9°'C. ( l i t . 12°C. ( 6 3 ) ) . The titrant used in a l l cases was an approximately tenth normal solution of perchloric acid in either glacial acetic acid or dioxane. These were prepared as follows: (a) Acetic acid solution. Approximately thirty-one ml. of a 60% aqueous solution of perchloric acid were mixed with glacial acetic acid (500 ml.), chilled acetic anhydride (33 ml.) was added, and the solution diluted to two l i t e r s . It was necessary to allow,the solution to stand overnight (or preferably two to three days) to allow the anhydride to react with the water present. This solution could be used with both indicator and potentiometer. (b) Dioxane solution. This solution was prepared in a,similar.manner, except that no acetic anhydride was added. Occasionally the .sol^ ution had a brown color:, but this did not influence the acid strength. The solution contained a small amount of water and so could only be used .for potentiometric ti t r a t i o n s . 14 Indicator A 0.5% solution of c r y s t a l v i o l e t i n g l a c i a l acetic acid was found to be.satisfactory. One drop was s u f f i c i e n t f o r a . f i f t y - f i v e ml. solution. The color change of c r y s t a l v i o l e t i s not sharp, passing from bluer-violet through various shades of blue and. green to yellow... The blue-green end point was e s p e c i a l l y d i f f i c u l t t o detect when the t i t r a t i o n produced a white p r e c i p i t a t e , as with potassium or ammonium s a l t s . The best procedure was to t i t r a t e a sample to the end-point pot-entiometrically, add the indicator and use t h i s solution as a standard. Potentiometer Two instruments were used: (a) A Beckman model G pH meter, s e r i a l number 1996, equipped with a glass saturated aqueous calomel electrode system. To obtain maximum s e n s i t i v i t y , the l a t t e r was of the sleeve type. The s o l u t i o n inside the electrode was changed frequently to prevent contamination by acetic acid d i f f u s i n g i n . The calomel electrode could not be used to t i t r a t e s o l -utions of lead or s i l v e r , since p r e c i p i t a t e s of the chloride froze the ground glass surface of the electrode. This d i f -f i c u l t y was overcome by using a salt bridge consisting of a U-tube, with ground glass stoppers at each end, f i l l e d with a supersaturated solution of l i t h i u m acetate i n acetic acid, with a l i t t l e added g e l a t i n to prevent c r y s t a l l i z a t i o n of 15 the s a l t . The U-tube connected the solution being analysed to a saturated aqueous solution of potassium chloride, into which the calomel electrode dipped. A l e s s cumbersome arrangement involved the use of a s i l v e r - s i l v e r chloride electrode as a reference electrode i n place of the calomel electrode. (b) A Fisher senior t i t r i m e t e r was used when a s e r i e s of analyses to the same end point was to be run. The same electrode systems as described above were employed. Standardization of the Titrant Sodium carbonate, or preferably potassium acid phthalate ( 5 7 ) , may be used as a primary standard. Approximately s i x tenths of a gram were dissolved i n g l a c i a l acetic acid (55 ml.) and t i t r a t e d either potentiometrically, or with c r y s t a l v i o l e t indicator to a blue-green end-point. Water acts as a weak base i n acetic a c i d . Its presence i n the solution when t i t r a t i n g potentiometrically can be tolerated up to 1.5% (44) but i t causes erroneous r e s u l t s when using an indicator. To minimize absorption from the at-mosphere, i t was convenient to use a l i p l e s s beaker closed by a t h i n rubber stopper through which a small hole was punched to admit the burette t i p . One must remember that organic l i q u i d s generally have a high c o e f f i c i e n t of cubical expansion and hence temperature changes w i l l have an appreciable effect on the normality of the solution. Allowance was made f o r t h i s by assuming that 16 the N/lO solution i n acetic acid or dioxane had the same co-e f f i c i e n t as the pure solvent. The metals and acetates were of the best obtainable purity, and were used without further p u r i f i c a t i o n . Other compounds were prepared as described. B. T i t r a t i o n of Metal Salts The metals studied were to be t i t r a t e d as acetates, so, to determine the s e n s i t i v i t y and accuracy of the method, .-. series of t i t r a t i o n s were c a r r i e d out on metal acetates as described below. To. check the accuracy each s a l t was an-alysed by the method indicated, according t o the in s t r u c t i o n s i n Vogel's "Quantitative Analysis" (64). The r e s u l t s of a l l the determinations are l i s t e d i n Table 1. Sodium Acetate The anhydrous salt was dried f o r one hour at 110°C. A convenient solution f o r t i t r a t i o n contained three m i l l i -equivalents of the sa l t i n f o r t y m i l l i l i t e r s of acetic a c i d . This concentration was used i n a l l subsequent analyses. It was necessary to warm the mixture to effect rapid solutionj the s o l u t i o n was then quickly c h i l l e d and t i t r a t e d immediately, eithe r potentiometrically or t o the blue-green c r y s t a l v i o l e t end-point. The gravimetric check was as sodium magnesium uranyl acetate. Acetates of Potassium, Lithium, Barium, Strontium, and  Ammonium These were analysed i n the same manner as sodium acetate. Since ammonium acetate i s so hygroscopic, i t had to be re-c r y s t a l l i z e d . This i s best done from methanol or acetic acid, followed by washing with ether and drying i n a vacuum desiccator at one centimeter pressure. A l l gave equally good r e s u l t s with indicator or potentiometer methods. The s a l t s were analysed gravimetrically as follows: potassium as perchlorate lithium as sulphate barium as sulphate strontium as sulphate ammonium (volumetrically) by t r e a t i n g with strong base, d i s t i l l i n g into excess standard acid and t i t -r a t i n g the excess. Magnesium Acetate Visual determinations with the indicator gave poor re-s u l t s , since the end-point occurs at an i l l - d e f i n e d yellow-green shade. The end point may be determined by comparing the sample being analysed with a solution of in d i c a t o r + magnesium perchlorate i n acetic acid i n a Coleman spectro-photometer. However, i t was easier and quicker to use the . potentiometric method. The gravimetric check was as mag-nesium ammonium phosphate. Acetates of Calcium, Manganese, and Cadmium These may be t i t r a t e d potentiometrically but not with the i n d i c a t o r . The gravimetric checks were as follows: calcium as the oxide, manganese as manganese ammonium phosphate, cadmium as the pyridine thiocyanate complex. Acetates of Cobalt and Nickel For reasons unexplained, no s a t i s f a c t o r y end-point 18 could be obtained when t i t r a t i n g with the acetic acid solution of perchloric acid. However, reproducible, 'but poorly-defined, end-points were obtained with a solution of perchloric acid i n dioxane. Both cobalt and n i c k e l were determined g r a v i -m e t r i c a l l y as the pyridine thiocyanate complex. Zinc Acetate zinc acetate i s quantitatively insoluble i n acetic acid, and so cannot be determined d i r e c t l y . It may be e a s i l y analysed by adding a measured excess of HCIO^ solution, and back t i t r a t i n g with standard sodium acetate solution. The gravimetric check was as the pyridine thiocyanate complex. Acetates of Lead and S i l v e r The calomel electrode could not be used i n either det-ermination f o r reasons described before, so both analyses were carried out using a s i l v e r - s i l v e r chloride electrode. In addition, s i l v e r acetate i s only p a r t l y soluble i n acetic acid, and had to be determined by the back t i t r a t i o n method described f o r zinc acetate. As a check, lead was determined as the molybdate and s i l v e r as the chloride. Acetates of Iron, Chromium and Aluminum Although reproducible end-points were obtained, they yielded anomalous r e s u l t s , so that the method could not be applied d i r e c t l y . ' The r e s u l t s are discussed further i n the next section. Acetates of Tin, Antimony, Copper, Bismuth and Uranium These acetates behave as i f they were e s s e n t i a l l y 19 undissociated i n acetic acid, and could not be t i t r a t e d by t h i s method. Since t i n (II) usually occurs as the chloride,' which in t e r f e r e s with the t i t r a t i o n , i t was removed by adding the required amount of a standard solution of lead acetate. T i t r a t i o n of Chloroacetates Several chloroacetates were prepared according to the method of Kastle and Keiser (57) . They were a l l soluble i n acetic acid, and, except for the ammonium s a l t , were analy-sed by the same method described previously f o r acetates. Each was also estimated gravimetrically by the same method as f o r the corresponding acetates, except f o r ammonium chloro-acetate. Its solution i n acetic acid and i n water changed strength continuously, possibly as a result of conversion into glycine, and could not be t i t r a t e d . T i t r a t i o n of Chlorides The presence of the chloride ion interfered with the t i t r a t i o n , a d i f f i c u l t y which was overcome by amethod due to P i f e r and Wollish (50, 5 2 ) . A s l i g h t excess of a solution of mercuric acetate i n acetic acid (6 gm. per 100 ml.) was added t o the solution to be t i t r a t e d . The halide was re-moved as HgX2, and the acetate ion lib e r a t e d could then be determined. Mercuric acetate i s e s s e n t i a l l y undissociated i n acetic a c i d (46) and the excess therefore did not i n t e r -f e r e . However, on standing, or on the addition of acetic anhydride, a complex mercuric compound i s formed (46) preventing the desired reaction. Hence the so l u t i o n should be f r e s h l y prepared, and.appreciable amounts of a c e t i c anhydride must not be present i n the solution to be t i t r a t e d . (The amount of acetic anhydride present may e a s i l y be de-termined (6, 41, 49).) C. Separation and Analysis of Metal Acetate Mixtures S o l u b i l i t i e s of Metal Salts i n Acetic Acid Zinc acetate was found to be quantitatively insoluble i n a c e t i c acid, and could be separated from other s a l t s which were completely soluble by di s s o l v i n g the mixture i n water, adding the r e q u i s i t e amount of acetic anhydride, and f i l t e r i n g . Various common reagents were tested as p r e c i p i t a t i n g agents with twenty-one representative metals' as outlined below. Whenever the metal acetate was only p a r t l y soluble i n acetic a c i d , the pr e c i p i t a t e was f i l t e r e d o f f and the resi d u a l solution tested. In a l l cases, a solution of one mil l i e q u i v a l e n t of metal acetate i n twenty m i l l i l i t e r s was used. Whenever a pr e c i p i t a t e was produced upon the addition of one of the test reagents, i t was removed by centrifuging and the l i q u i d tested f o r completeness of p r e c i p i t a t i o n as follows: The solution was evaporated almost to dryness twice with excess concentrated n i t r i c acid to remove a l l organic matter, and the residue taken up i n d i l u t e n i t r i c acid. The s o l u t i o n was then tested for the presence of the 21 appropriate ion according to the methods outlined i n Treadweil-Hall (62). Hydrochloric, Hydriodic, and Hydrosulphuric Acids In each case the dried gas was passed through the s o l -ution u n t i l the l a t t e r was saturated. Sulphuric Acid S u f f i c i e n t c h i l l e d acetic anhydride was added t o c h i l l e d reagent grade acid to react with a l l the water present. The mixture was allowed to stand out of contact with a i r f o r twenty four hours to allow the reaction to go to completion, r e s u l t i n g i n a yellow-tinged very viscous l i q u i d . Enough of t h i s was added t o each te s t solution to give a f i v e - f o l d excess. Those which did not react were allowed t o stand f o r two days, a f t e r which a portion of each was d i l u t e d with 5% water. N i t r i c Acid Reagent grade f i f t e e n molar acid was added to each tes t solution to give a f i v e - f o l d excess. Phosphoric Acid Reagent grade 85% acid was added to each solution to give about a f i v e - f o l d excess. Benzoic, Oxalic, S a l i c y l i c . T a r t a r i c Acids These are not very soluble i n acetic a c i d at room temperature, so a saturated s o l u t i o n at 100°C. was used, the test solutions also being maintained at t h i s temperature. About a f i v e - f o l d excess was added i n each case. The r e s u l t s of these two hundred ten t e s t s are summarized i n Table 2, Since the r e s u l t s with the sulphuric a c i d were at variance with those of Davidson (9) i n that only three metals gave pr e c i p i t a t e s , the following additional t e s t s were carried out. The sulphates of Mg, Mn(II), C r ( I I I ) , Fe(III), Ca, Co, Nr and Sn(II) were refluxed with g l a c i a l acetic acid f o r four hours, a f t e r which the l i q u i d was tested f o r the presence of the appropriate ion i n the manner described before. Then the sulphates were dissolved i n water, and s u f f i c i e n t a c etic anhydride added to react with a l l of i t , producing the same anhydrous medium as i n the f i r s t set of t e s t s . The r e s u l t s of these experiments are given i n Table I I I . Separations Effected On the basis of the foregoing t e s t s , the following a n a l y t i c a l separations were carr i e d out: Zinc and Magnesium Three m i l l i e q u i v a l e n t s each of the two metals were mixed and dissolved i n d i l u t e n i t r i c a c i d (5 ml.). The solution was d i l u t e d with a c e t i c acid (15 ml.) and heated to b o i l i n g . Acetic anhydride (33 ml.) was added cautiously, the mixture being boiled u n t i l the yellow color disappeared. The beaker was covered with aluminum f o i l , the suspension d i -gested on a steam bath f o r ten minutes, c h i l l e d i n i c e -water, and centrifuged. The magnesium solution was t i t r a t e d immediately. The zinc may be determined as described above, or by difference. 23 zinc and Copper The metals were separated by the same procedure as above. The zinc was determined as described before, and the copper obtained by difference. Barium and Calcium Three m i l l i e q u i v a l e n t s of each metal were dissolved i n n i t r i c acid (5 ml.), the solu t i o n d i l u t e d with acetic acid (15 ml.), heated to b o i l i n g , and acetic anhydride (33 ml.) cautiously added. The solution was boiled u n t i l 'the yellow color disappeared, then c h i l l e d i n i c e water. Concentrated n i t r i c acid (2 ml.) was added, and the solution vigorously s t i r r e d . The p r e c i p i t a t e of barium n i t r a t e was removed by centrifuging, and the calcium solution immediately t i t r a t e d . The barium may be determined by difference, or by decomposing the n i t r a t e with a c e t i c anhydride, and t i t r a t i n g i n the usual way. S i l v e r and Aluminum Three m i l l i e q u i v a l e n t s each of the two metals were brought into s o l u t i o n as above, except that concentrated n i t r i c acid was used. The solution was c h i l l e d and saturated with dried H2S gas. The s i l v e r sulphide was removed by c e n t r i -fuging, dissolved i n d i l u t e n i t r i c acid, and the solution boiled to expel the l i b e r a t e d HgS. S u f f i c i e n t a c e t i c an-hydride was added t o react with the water and decompose the n i t r a t e . T h i r t y - f i v e m i l l i l i t e r s of standard perchloric acid solution were added, and the excess back-titrated with standard sodium acetate solution, using s i l v e r - s i l v e r chlor-ide and glass electrodes. Lead-Tin The procedure was exactly the same as i n the preceding separation; great care had to be exercised not to allow any water to enter the system u n t i l the separation was com-plet e , since stannous sulphide would p r e c i p i t a t e i f water was present. The lead sulphide was removed by centrifuging, treated and analysed as the s i l v e r sulphide above. Magnesium and Antimony; Magnesium and Lead; Magnesium and  Copper The lead, antimony, and copper were removed with hydrogen sulphide; the procedure, being exactly the same as i n the preceding case, except that aqua regia was required to d i s -solve the antimony. S u f f i c i e n t mercuric acetate reagent (6 gm. per 100 ml.) was added to the magnesium solution, a f t e r removing the sulphide,, to react with a l l the chloride present. The magnesium may then be determined i n the usual manner. If accuracy greater than four percent i s not required, the magnesium may be t i t r a t e d d i r e c t l y i n the presence of the antimony. These r e s u l t s were generally of the order of k% too high. S i m i l a r l y , magnesium may be t i t r a t e d d i r e c t l y i n the presence of copper; the r e s u l t s were of the order of &% too high. Nickel-Chromium About three m i l l i e q u i v a l e n t s of each metal were mixed ' and dissolved i n d i l u t e hydrochloric (5 ml.) acid (since n i t r i c acid renders chromium passive). The solution was heated to b o i l i n g and a c e t i c anhydride {'j>0 ml.) cautiously added. A f i v e - f o l d excess of a hot solution of oxa l i c acid i n acetic acid was added, the beaker covered and c h i l l e d : i n i c e water. The pr e c i p i t a t e was removed by centrifuging, evaporated twice to dryness with excess concentrated n i t r i c acid, and the residue taken up i n d i l u t e n i t r i c a c i d . The solution was treated with acetic anhydride i n the usual manner, and the n i c k e l t i t r a t e d with perchloric acid i n dioxane.. Cadmium and Magnesium A solution of three m i l l i e q u i v a l e n t s of each metal was made up i n the usual way, c h i l l e d , and saturated with dry hydrogen iodide gas. The pr e c i p i t a t e was removed by c e n t r i -fuging, washed, and the r e s u l t i n g solution of magnesium t i t r a t e d . Tin and Magnesium A solution of three m i l l i e q u i v a l e n t s of each metal i n acetic acid, was prepared i n the usual way, but kept at 100°C. A . f i v e - f o l d excess of oxalic a c i d solution i n acetic acid was added, and the mixture c h i l l e d . The p r e c i p i t a t e of mag-nesium oxalate was.removed' by centrifuging, evaporated to . dryness twice with concentrated n i t r i c acid, and the residue taken up i n d i l u t e n i t r i c acid. Acetic anhydride was added to react with the water and decompose the n i t r i c acid, a f t e r which the magnesium was t i t r a t e d i n the usual way. • The r e s u l t s of a l l these determinations are summarized i n Table IV i n the next section. D. Introduction of Metals Into Acetic Acid Solution Acetic acid i t s e l f i s e f f e c t i v e i n dissolving only the most elec t r o p o s i t i v e of the metals. I f , however, the acid strength could be increased, more metals could be introduced d i r e c t l y into the organic phase. A possible method of accomplishing t h i s suggests i t s e l f from the following con-siderations: I f AB i s a very weak organic base and HX a weak acid, then a mixture of AB and HX w i l l remain a mixture (or s o l -ution) of one i n the other with no chemical reaction occur-r i n g . On the other hand, i f AB i s a strong organic base and HX a strong a c i d , a s a l t w i l l be produced on mixing: AB + HX -> ABH+ + X" With a f a i r l y weak aci d (say CH^COOH) and a f a i r l y strong base (piperidine f o r example) one would obtain a compound of the sort AB...H+...X~ with a hydrogen bond, and a greater or l e s s e r " a c t i v i t y " depending on the strength of the base and the temperature. I f the a c t i v i t y i s s u f f i c i e n t l y great, the p o s i t i v e ion may be displaced by metals of high enough p o t e n t i a l . To test t h i s theory, twenty f i v e representative metals 27 were treated with ammonium acetate, piperidinium acetate, and dicyclohexylammonium acetate, the l a t t e r t o give a higher b o i l i n g s a l t . Test tube experiments only were carried out, to cover the ground as rapidl y as possible. No attempt was made to determine the extent of,the reaction i f the metal did not dissolve completely. The res u l t s of the tests are summarized i n Table 17 and Table V. Preparation of Piperidinium Acetate Piperidine ( 4 9 . 4 ml.) and ace t i c acid ( 2 8 . 6 ml.) were dissolved i n separate portions of acetone. Mixing immedi-ately produced a w e l l - c r y s t a l l i z e d white p r e c i p i t a t e . The compound was not stable toward heat; drying at 5 5 ° C caused i t to turn yellow. Therefore, a f t e r f i l t e r i n g and washing with ether, the s a l t was stored i n vacuo u n t i l dry. The y i e l d was 7 1 . 5 gm. (98%) of a s o l i d melting at 1 0 1 - 1 0 4 ° C . with decomposition. Preparation of Decyclohexylammonium Acetate Dicyclohexylamine ( 1 0 0 . 7 ml.) and acetic acid ( 2 8 . 6 ml.) were dissolved i n separate portions of acetone. Mixing pro-duced a powdery white p r e c i p i t a t e i n 98% y i e l d , melting at 1 1 6 - 1 1 8 ° C . A second method of increasing the a c t i v i t y of an acid i s t o introduce electro-negative atoms into the molecule, as i n the chloroacetic acids. To determine the extent of t h i s increase i n a c t i v i t y a survey was made of the r e a c t i v i t y of the twenty-five metals mentioned toward pure monochloroacetic, 28 d i c h l o r o a c e t i c , ^  and t r i c h l o r o a c e t i c acids and equimolar s o l -utions of each i n acetic acid. Each test was c a r r i e d out from room temperature to the b o i l i n g point of the a c i d con-sidered. These r e s u l t s are also summarized i n Table IV and Table V. . ' . However, f o r reasons discussed i n the next secion, these acids were unsatisfactory f o r general use, and the f o l l o w i n g method was adopted. To introduce the metals into acetic acid solution, s u f f i c i e n t n i t r i c acid was used to bring the metal(s) into aqueous solution, followed by s u f f i c i e n t acetic anhyd-ride to react with a l l the water added and to decompose a l l the n i t r a t e present. In c e r t a i n cases, aqua regia was re-quired, but i n general t h i s method was fast and e f f i c i e n t . I I I . RESULTS AND DISCUSSION A. Determination of the Potentiometric End-Point The t i t r a t i o n curves obtained are i l l u s t r a t e d i n plates jr~XJ|i8. The f i r s t i n each set i s a plot of volume of t i t r a n t v. e.m.f. reading. The end point i s , of course, the i n f l e c t i o n point on the curve, but i n t i t r a t i o n s the p o t e n t i a l break was not sharp enough to determine the end point accurately from the graph. A more precise method i s to plot -^-| , the change i n potential per unit volume, v. V, the volume added near the end point, as was done i n the second graph i n each set. Since the point of i n f l e c t i o n i n the t i t r a t i o n curve has the condition "7 7^ = o, the dV^ ' corresponding point on the d i f f e r e n t i a l curve must be a maximum, i . e . (~) = o. I f the l a t t e r curve i s not symmetrical, the end-point may be found by extrapolation (39), assuming that i n a small region about the end-point dE varies l i n e a r l y with V. For a l l but the most basic of av the metal acetates, the s e n s i t i v i t y of the t i t r a t i o n did not j u s t i f y extrapolating the volume to f r a c t i o n s of a drop. The end points are indicated on a l l the graphs. An attempt was made to calculate the error involved i n the t i t r a t i o n s on a th e o r e t i c a l basis, by adapting the method of Roll e r (54-56) . This was found to be impossible, however, because of a lack of data on d i s s o c i a t i o n constants i n acetic acid. The s e n s i t i v i t y of the method may be increased by d r i l -l i n g a hole i n the side of the beaker and i n s e r t i n g one electrode h o r i z o n t a l l y so that the t i p i s very close to that of the other electrode, which i s i n the usual v e r t i c a l pos-i t i o n . The increase i n s e n s i t i v i t y i s caused by a decrease i n e l e c t r i c a l resistance. Since fluorpsulphonic acid i s too active to exist i n water, i t was thought i t might be stronger i n acetic acid than perchloric acid, and hence give better defined end-points with the weaker bases. However, t h i s was not the case; the a c i d i t y , as measured by the po t e n t i a l , was l e s s than that of perchloric acid. Moreover the .end-point -30 obtained when the fluorosulphonic acid was used as t i t r a n t was very poorly defined'. • B. T i t r a t i o n of Metal Salts The r e s u l t s obtained from the t i t r a t i o n s are summarized' i n Table I; the t i t r a t i o n curves are i l l u s t r a t e d i n Plates I T - XJU B . The t i t r a t i o n curves show the same' changes as one proceeds from strong to weak bases as i s found i n water. Acetates of the f i r s t and second group metals show ' a pronounced pote n t i a l "break* at the end-point, while those of metals usually considered "weak" i n water show a les s pronounced break. An exception was ammonium acetate, which i n acetic acid i s as strong as potassium acetate. In the case of ni c k e l cobalt and cadmium, the accuracy of the t i t -r a t i o n was not always as great as the res u l t s i n Table I might indicate, since the end-point was poorly defined. Contrary to the res u l t s published by P i f e r and Wollish (51) no end-point could be found f o r bismuth, antimony and t i n , and a very low value was obtained f o r the percentages of i r o n . Inspection of Plates iSJ r t N I >SI shows that the curves obtained f o r bismuth, antimony, copper, t i n , and uranium are of the same form as that f o r the solvent alone i . e . t h e i r acetates are bases too weak to be t i t r a t e d i n acetic acid with perchloric a c i d . The curves f o r chromium, iron and aluminum yielded end-points showing approximately 11-12% purity, while the gravimetric analysis gave r e s u l t s close to 100%. This i n spite of the fac t that the poten t i a l breaks are more TABLE I TITRATION OF.METAL SALTS Salt % Purity Found Gravimetric Determination NaOAc 99.59 99.55 99.50 99.58 99.52 KOAc 99.56 99.41 99.49 99.48 99.47 LiOAc 99.73 99.70 99.77 99.75 99.60 Ba(0Ac) 2 99.40 99.47 99.49 99.41 99.46 Sr(0Ac) 2 99.65 99.60 99.49 99.41 99.46 NH^OAc 97.58 98.08 97.58 97.73 97.95 Mg(0Ac) 2 99.82 99.69 99.71 99.85 99.74 Mn(0Ac) 2 99.69 99.31 99.88 99.62 99.73 Ca(0Ac) 2 98.95 99.06 99.00 98.96 99.05 Cd(0Ac) 2 98.72 98.25 98.55 9 8 . 6 0 - 98.72 Co(0Ac)2 91.35 91.0*5 91.17 91.06 91.00 Ni(0Ac) 2 93.61 93.30 93.21 93.29 93.42 Zn(0Ac) 2 97.73 97.50 97.26 97.40 97.30 Pb(0Ac) 2 98.36 98.49 98.57 98.30 98.33 Ag(0Ac) 2 „ 99.72 100.13 99.62 99.83 99.76 Fe(OAc) 3 l i 11.36 11.39 11.59 98.96 98.90 C r ( O A c ) 3 i 11.71 11.50 11.44 99.65 99.43 Al(OAc ) 3 S 11.06 11.23 11.15 99.03 99.15 ClCH 2 C00Na 97.96 98.13 98.07 98.10 98.00 C1CH2C00K 97.83 99.00 99.01 98.96 98.89 (GlCH 2 C00) 2Ba 97.36 97.30 . 97.46 97.50 97.35 (ClCH 2 C00) 2Mg 97.35 97.46 97.33 97.49 97.46 A F i r s t column i n M % purity found" had concentration of three m i l l i e q u i v a l e n t s of sample per 4 5 ml. of solvent, the second s i x m i l l i e q u i v a l e n t s and the t h i r d nine m i l i i e q u i v a l e n t s per 4 5 ml. Each value i n "% purity found' represents the mean of two determinations. " 3 2 c l e a r l y defined than those f o r cetain metals whose curves y i e l d reasonable r e s u l t s . A possible explanation f o r these low r e s u l t s may be i n the work of Weinland et. a l ( 6 5 - 7 0 ) , who found that solutions of. chromic, and f e r r i c s a l t s i n acetic acid give r i s e to numerous complexes many of which were quite stable. , One of these has the form . [cr^OH.)^ (AcO)^J AcO. T i t r a t i n g t h i s with perchloric acid solution would result i n the compound ^Cr-^OH^tAcO^.J CIO^,. that i s , only one molecule of HCIO^ has been consumed f o r three atoms of chromium instead of the expected three molecules of HCIO^ f o r one atom of chromium. , This would give r i s e to a r e s u l t 1 1 . 1 1 % of the theoretical,_not f a r from the values obtained. Moreover, the values obtained remained e s s e n t i a l l y constant over a three-fold increase i n concentration,, which would indicate that a d e f i n i t e compound or set of compounds i s involved, under the conditions at hand, and not an equilibrium mixture. Further work along t h i s , l i n e would promise to y i e l d very i n t e r e s t i n g r e s u l t s . When the end-point i s i n d i s t i n c t or poorly defined'in acetic acid alone, i t may be improved, as shown recently by P i f e r et a l . , ( 5 3 ) by the addition of a miscible organic l i q u i d of low d i e l e c t r i c constant. ..This was i l l u s t r a t e d by the improvement of the end-point f o r cobalt, n i c k e l , and cadmium acetates when a solution of.perchloric acid i n dioxane was used as a t i t r a n t . The reason f o r the increase i n s e n s i t i v i t y i s that perchloric acid i s l e s s ionized i n the solvent of lower d i e l e c t r i c constant, and hence, as shown i n the introduction, the solution has a higher a c i d i t y . C. S o l u b i l i t i e s i n Acetic Acid Zinc acetate i s the only one quantitatively insoluble i n a c e t i c acid; those of s i l v e r and thorium, aluminium and bismuth are only p a r t l y soluble. When an aqueous solution of arsenious a c i d i s treated with'racetic anhydride, a white p r e c i p i t a t e forms, but i t i s doubtful i f t h i s i s arsenic acetate. The re s u l t s of the tests with possible p r e c i p i t -ating reagents, as summarized i n Table I I , have several i n t e r e s t i n g features. In many cases the s o l u b i l i t i e s p a r a l -l e l l t h o s e i n the aqueous system,, but there are some out-standing exceptions. For instance, strontium and barium n i t r a t e s are precipitated immediately and quantit a t i v e l y upon the addition of concentrated n i t r i c acid or a solution of sodium n i t r a t e i n acetic acid. Lead i s also p r e c i p i -tated, but not quantitatively, and magnesium p r e c i p i t a t e s a f t e r standing overnight. Although barium chloride was insoluble i n a c e t i c acid, strontium chloride was very soluble, i n comparison with other chlorides. Cadmium chloride also p r e c i p i t a t e d r e a d i l y , but tended to dissolve i n excess HCI. P r e c i p i t a t i o n was not quantitative, but was when hydrogen iodide i s used. Hydrogen sulphide gave p r e c i p i t a t e s which were l e s s c o l l o i d a l than those obtained i n water i n certa i n cases, notably cadmium lead and s i l v e r . (In general, p r e c i p i t a t e s TABLE II RESULTS OF ADDING THE REAGENTS LISTED TO A OF 1 MILLIEQUIVALENT OF THE METAL ACETATE IN SOLUTION ACETIC ACID Reagent Ca Mg Th A l U Mn Cr Mo Fe Cd Co Na Sn Pb Sb Bi As Se Te Ag Sr Ba Anhydrous I^SO^ - - - x - - - - X X - H2 S 0 4 + 5% H 20 X x - - - 0 - - - - X X X 15N HNO3 a - - - - - - -Dry HCl gas - - _ 0 x - - - - 0 fa Dry HI gas - - x x - - - - X Dry H 2S gas O O X X X X X X X X 0 X 85% H 3P0 4 x fa fa 0 fa fa x 0 S a l i c y l i c acid 0 Benzoic Acid T a r t a r i c Acid X X X O X X X X X X X Oxalic acid X X X X X X X X LEGEND '•" x Quantitative p r e c i p i t a t i o n . 0 Non-quantitative p r e c i p i t a t i o n & Quantitative p r e c i p i t a t i o n on standing several hours. fa Non-quantitative p r e c i p i t a t i o n on standing several hours. Not tested. •F-formed i n a c e t i c a c i d are not well c r y s t a l l i z e d ) . A l l the metals found i n group II i n the usual a n a l y t i c a l scheme pre-c i p i t a t e d r e a d i l y except t i n , which remained i n solution as long as no water was present. (Addition of 1% water caused formation of a brown precipitate.) Nickel and cobalt s u l -phides, which do not p r e c i p i t a t e i n a c i d i c aqueous solutions, p r e c i p i t a t e d more slowly than the metals of group II , but p r e c i p i t a t i o n was quantitative a f t e r f i f t e e n minutes. There was also a reaction with chromium to produce a small amount of grey-green c o l l o i d a l p r e c i p i t a t e . This was somewhat sur-p r i s i n g , as there was no such reaction with i r o n or aluminum, and prevented the separation of other metals from chromium using hydrogen sulphide. Although no quantitative p r e c i p i t a t i o n s resulted from addition of s a l i c y l i c acid, color changes i n several cases showed complexes had been formed. Iron changed from red to dark brown, cobalt from red to rust brown, n i c k e l from green to yellow, and the solutions of aluminum and antimony from c o l o r l e s s to yellow. The molybdenum sample became deep red v i o l e t , with the formation of a very f i n e brown p r e c i p i t a t e . Other substituents on the benzene nucleus might r e s u l t i n the formation of insoluble complexes. Benzoic acid produced no p r e c i p i t a t e with any metals, but t a r t a r i c and oxalic acids were very e f f e c t i v e , almost a l l the p r e c i p i t a t i o n s being quantitative. Although p r e c i p i t a t i o n i s best done i n a hot solution with these reagents and s a l i c y l i c acid because of s o l u b i l i t y consider-ations, precipitates were better crystallized when brought down in the cold in acetic acid. However, almost a l l were colloidal in nature, and so a centrifuge was much preferred to a f i l t e r for separation. Perhaps the most unusual phenomena occurred with sul-phuric acid and the sulphates. The results of the experi-ments carried out are summarized in Table III.. According to Davidson (14) sulphates in general are insoluble in acetic acid, and metals precipitate on the. addition of anhydrous sulphuric acid. Table I I I shows that four hours refluxing failed to dissolve an appreciable amount of any of the sul-phates tested, but on the other hand, Table II shows that most metals failed to be precipitated upon the addition of sulphuric acid. Moreover, when the sulphates were dissolved in water and the solution treated with acetic anhydride, two, chromium and iron,, failed to precipitate at a l l , and the precipitation of several others was far from complete. These results show that one cannot rely on the behavior of the sulphates, and that sulphuric acid i s not a good precipe itating agent f o r this reason. D. Separations As mentioned before, precipitates formed in hot acetic acid solution are very finely divided, and often d i f f i c u l t to separate from the f i l t r a t e , even in a high-speed centri-fuge. Precipitates formed in the cold settle more readily, 37 TABLE I I I RESULTS OF SOLUBILITY TESTS WITH SULPHATES Sulphate Result of Refluxing with G l a c i a l Acetic Acid Result of d i s s o l v i n g i n water and adding acetic anhydride. . Mg Insoluble Quantitative p r e c i p i t a t i o n Mn(II) Insoluble P r e c i p i t a t i o n almost quantitative Cr Insoluble No p r e c i p i t a t i o n Fe(III) Insoluble P a r t i a l p r e c i p i t a t i o n on addition of l a s t 5%"of acetic anhydride, but pre c i p i t a t e dissolves on b o i l i n g Ca Insoluble Quantitative p r e c i p i t a t i o n Co Insoluble P a r t i a l p r e c i p i t a t i o n < Ni Insoluble P a r t i a l p r e c i p i t a t i o n Sn Insoluble P r e c i p i t a t i o n almost quantitative. but are not well c r y s t a l l i z e d and clog f i l t e r paper. Thus a centrifuge should be used whenever possible. The r e s u l t s obtained are l i s t e d i n Table IV. They show that good precision and accuracy are obtainable with t h i s method, often with appreciable saving of time. The complete analysis of an a l l o y of cadmium and magnesium, f o r example, requires approximately f o r t y - f i v e minutes, i f a standard per-c h l o r i c acid solution i s a v a i l a b l e . I f a rough estimation of a binary a l l o y of copper or antimony with a t i t r a t a b l e metal i s required, i t may be completed i n f i f t e e n minutes by t i t r a t i n g the metal i n the presence of the copper or antimony. E. Introduction of Metals Into Acetic Acid Solution The r e s u l t s of the tests carried out with the amine acetates, acetic acid, and the chloroacetic acids are sum-marized i n Tables V and VI. Ammonium acetate and p i p e r i d -inium acetate were completely i n e f f e c t i v e i n attacking metals other than calcium and magnesium, and these reactions were very slow. For t h i s reason they are not included i n t h e % t a b l e The r e s u l t s with chloroacetic acids show that even elements below hydrogen i n the electromotive series (such as copper and mercury) are attacked vigorously by these acids. According to Doughty (18) the reaction between copper and t r i c h l o r o a c e t i c acid i s not a simple one. A variety of com-pounds, depending on the temperature and concentration, are produced, such as copper chloroacetate, copper dichloroacetate chloroform and other products. 39 TABLE IV RESULTS OF TITRATING METALLIC MIXTURES Mixture T h e o r e t i c a l % of Metal Determined % Found Zn-Mg Mg determined 27.11 2 7 . 0 0 26.95 26.98 Zn-Cu Zn determined 5 0 . 7 0 50.59 5 0 . 5 0 50.65 Ba-Ca Ca determined 22.58 22.56 22.75 22.69 Ag-Al Ag determined 7 9 . 9 9 79.86 79.91 79.81 Pb-Sn Pb determined 6 3 . 5 8 63.42 63.49 63.36 Mg-Sb Mg determined I 6 . 6 5 1 6 . 4 8 16.53 16.57 Mg-Sb Mg determined i n presence of Sb 16.65 17.53 17.67 17.73 Mg-Cu Mg determined 27.67 27.38 27.46 2 7 . 5 5 Mg-Cu Mg determined i n presence of Cu 27.67 29.58 29.65 29.71 Mg-Pb Mg determined 10.50 10.57 1 0 . 4 0 10.48 Ni-Cr N i determined 5 3 . 03 52.85 52.89 52.93 Cd-Mg Mg determined 17.79 17.60 17.69 17 ".59 Sn-Mg Mg determined 1 7 .00 1 6 . 8 7 16.91 17.10 TABLE V. REACTIVITY OF AMINE ACETATES, ACETIC ACID CHLOROACETIC ACIDS TOWARDS METALS AND THE Metal Dicyclo- Pure Equimolar Equimolar Equimolar Pure Pure Pure hexyl- acetic soln. of soln. of soln. of chloro- dichloro- t r i c h l o r o -ammonium acid chloroacetic dichloroacetic t r i c h l o r o a c e t i c acetic acetic acetic acetate i n acetic i n acetic i n acetic acid . acid acid Ca ++ +++ +++ +++ +++ +++ +++ +++ Mg ++ +++ +++ +++ +++ +++ +++ +++ Th — + _ _ — — _ A l _ _ +++ _ + + + ++ U _ + +++ + ++ + ++ Mn + +++ +++ +++ +++ +++ +++ +++ Zn + ++ +++ +++ +++ +++ +++ +++ Cr _ ++ +++ +++ +++ ++ +++ Mo _ + + + ++ + + +++ Fe ++ * +++ • +++- ++.+ ++ +++ Cd mm +++ + + ++ — Co + + + ++ ++ + + ++ Ni _ _ + — + — Sn _ +++ +++ ++ +++ Pb +++ +++ + ++ +++ Sb + + + ++ +++ Bi - - — - — •. - + ++ As — _ - - - -Cu S o + ± - - ++ +++ u c Te _ _ — — - - -Ag _ - - - mm Hg — . . - - - -W • _ — — — - - -S i - - - - - — — LEGEND: +++ vigorous reaction + slow reaction - no apparent reaction ++ moderate reaction ± slight' reaction O 41 TABLE VI ACIDS TESTED WHICH ATTACK THE METALS CONSIDERED WITH SUFFICIENT RAPIDITY. TO BE OF PRACTICAL USE i Reagent Metals which dissolve r e a d i l y Dicyclohexyl-ammonium acetate None Piperidinium acetate :None : Ammonium acetate Ca Acetic Acid Ca, Mg, Mn, Fe Equimolar chloroacetic i n acetic Ca, Mg, A l , Mn, Zn //' Cr, Fe Equimolar dichloroacetic i n .acetic Ca, Mg, U, Mn, Zn., .Cr,, Fe, Cd, Sn, Pb // Co Equimolar t r i c h l o r o a c e t i c i n acetic Ca, Mg, U, Mn, Zn, Cr, Fe, Sn, Pb // Mo, Co, Hg Pure chloroacetic Ca, Mg, Zn, Mn, Cr, Fe // U Pure dichloroacetic Ca, Mg, Mn, Zn // Cr, Fe, Cd, 'Sh/ Pb, Sb * 1 : ' Pure t r i c h l o r o a c e t i c Ca, Mg, Mn, Zn, Cr, Mo, Fe, Sn, Pb, Sb, Cu // A l , U, Co, Bi Those elements ; a f t e r the double stroke iri the table react much les s rapidly than those before i t . 1+2 Beside possible undesirable side reactions, another disadvantage of these acids from the point of view of t i t -rations was the large quantity of acid often required to dissolve the metal completely. Because of the l i m i t e d s o l -u b i l i t y of chloroacetic and t r i c h l o r o a c e t i c acids i n acetic acid, the r e s u l t i n g mixture was often not completely s o l -uble i n acetic acid. Moreover, the table shows certain common a l l o y s would not react r e a d i l y with a s i n g l e a c i d . For these reasons the method of introducing metals into acetic acid solution using n i t r i c acid and acetic anhydride was preferred. Using t h i s method, a l l o y s corresponding to the metal acetate mixtures discussed previously may r e a d i l y be analysed i n a minimum of time and with good accuracy. IV. CONCLUSIONS Alloys, and mixtures of the l e s s soluble metal com-pounds, may be conveniently introduced into a c e t i c acid solution using n i t r i c acid (or aqua regia) and acetic an-hydride. Quantitative separations may often be effected i n t h i s medium using common reagents such as hydrogen s u l -phide and hydrogen iodide, n i t r i c , oxalic and t a r t a r i c acids. The constituents may then be t i t r a t e d r a p i d l y and accurately with a standard solution of perchloric a c i d i n acetic a c i d . In spite of the interest shown i n a n a l y t i c a l methods such as the one described, comparatively l i t t l e i s known of the chemistry of acetic acid solutions. Much work needs to be done on the determination of d i s s o c i a t i o n constants, on s o l u b i l i t y r e l a t i o n s h i p s , and on the founding of a more rigorous t h e o r e t i c a l approach to the question of a c i d i t y i n non-aqueous solutions. With a clear insight into these questions, t h i s type of analysis may prove to be of much wider a p p l i c a b i l i t y than i s now imagined. i i t 44 BIBLIOGRAPHY 1 . Blumrich, K. and Bandei, G. Agnew. 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Ber. 42 : 2997. 1909. 67. Weinland, R. and Gussman, E. Ber. 42 : 3881. 1909. 68. Weinland, R.-and Gussman, E. Z. anorg. Chem. 66 : 157; 1910. 69. Weinlandj R;, Kessler, K., and Bayerl,. A. Z. anorg. Chem.-132 : 209. 1924. 70; Weinland, R; and 0. Loebich. Z. anorg. allgem. Chem. 151 : 271. 1926. 71; Wittman,G.- Agnew. Chem. A60 : 330. "1948;-700 FIGURE I I A Po t e n t i o m e t r i c T i t r a t i o n of Ten M i l l i l i t e r s of Standard Sodium Acetate S o l u t i o n , With P e r c h l o r i c Acid In Ac e t i c Acid 600 500 400 _i i i i i i _ 8.00 9.00 10.00 11.00 ML. TITRANT'ADDED,(0.09927N.) 12.00 FIGURE II B S HH LL a LL o < O CO •-3 1 4 12 10 8 0 Precise Determination of the End Point f o r Figure II A END POINT . i • i i 1 1 1 1 1 30 40 50 60 70 VOLUME ADDED NEAR END POINT (1 DROP 10 20 80 90 0.0148 ML.) 100 110 120 FIGURE I I I A T i t r a t i o n (potentiometric) of Approximately Three M i l l i e q u i v a l e n t s of Barium Acetate With Perchloric Acid i n Acetic Acid END POINT 600--i • * • • • 23.00 24.00 25.00 26 .00 ML. TITRANT ADDED (0.09927N.) 27.00 i i t- 1 C3 t-< ?G M O < •n C CO i-5 cc 1 4 1 2 10 8 0 FIGURE I I I B Precise Determination of the End Point f o r Figure I I I A Gr 1 0 END POINT 20 3 0 40 50 6 0 7 0 80 9 0 100 1 1 0 VOLUME ADDED NEAR END POINT ( l DROP = 0.0148 ML.) FIGURE IV A 700 600 500 P o t e n t i o m e t r i c T i t r a t i o n o f Approximately Three M i l l i -e q u i v a l e n t s of L i t h i u m and Strontium A c e t a t e s i n A c e t i c A c i d With P e r c h l o r i c A c i d . Ammonium and Potassium Ac e t a t e s Each Y i e l d A S i m i l a r Curve. END POINT END POINT 8 460 _ i i i i i i 23.00 24.00 25.00 26.00 27.00 ML. TITRANT A D D E D (0.09927N,) FIGUREIV B 14 f- Precise Determination of the End Point f o r Figure U A 12 2ND POINT 10 8 0 1 0 10 20. 30 40 50 60 70 80 90 100 • 110 VOLUME ADDED NEAR END POINT (1 DROP = 0.014^ ML.) 120 END POINT 650 FIGURE V A Potentiometric T i t r a t i o n of Approximately Three M i l l i -equivalents of Magnesium Acetate i n Acetic Acid With Perchloric Acid $00 550-520 2^.00 25.00 26.00 27.00 2&.00 ML. TITRANT ADDED ( 0 . 0 9 9 2 7 N . ) I6d FIGURE V B Precise Determination of the End Point f o r Figure V A 4* 11A § 3 M l-i < o cc DJ3 ao 40 20 0 E N D POINT • • - i 1 i 0 10 20 30 40 50 60 70 80 90 100 V O L U M E A D D E D N E A R E N D P O I N T ( l D R O P = 0 . 0 1 4 8 M L . ) 110 120 FIGURE VI A Potentiometric T i t r a t i o n of • Approximately Three M i l l i - END POINT, equivalents of Calcium Acetate ^ 1 i n Acetic Acid With Perchloric ! Acid I 1 1 ' ' 1 ' i i 1 1 1 1 i i i i i i • i • 2 3 . 0 0 . . 2 4 . 0 0 2 5 . 0 0 ML. TITRANT ADDED (0.09927N.) I 8 21 11 FIGURE VI B P r e c i s e Determination of the End Point f o r Figure VI A END POINT' 0 i i i i i i i i i i , i l i s 1 lio o 10 20 30 40 50 60 70 80 90 100 VOLUME ADDED NEAR END POINT (1 DROP = 0.0148 ML.) • FIGURE VII A ML. TITRANT ADDED (0 .100SN.) FIGURE VII B 3.5 3 . 0 <|M 2 . 5 g P 2 . 0 ' ait-1 o 00 < o •-3 o: 1 . 5 1.0L 0 . 5 0 0 Precise Determination of the En<i Points f o r Figure VII A END POINT i i 10 2 0 3 0— J 40 ' 50 6 0 7 0 SO 90 VOLUME ADDED NE2|R END POINT ( 1 DROP = 0 . 0 1 0 9 M L . ) 1 0 0 -0 _i i i — . — i u_ 1 1 1 0 1 2 0 ML. TITRANT ADDED (0.1031 N.) END POINT i FIGURE VIII B '5:' 0 Precise Determination of the.End Point f o r Figure VIII A -flu 0 10 20 30 40 50 60 70 90 100 110 120 VOLUME ADDED NEAR END POINT ( 1 DROP = 0 . 0 1 3 8 ML.) END POINT FIGURE IX A i i 3 0 . 0 0 ML. TITRANT ADDED ( 0 . 1 1 1 6 N) F I G U R E IX B 700 FIGURE X A Potentiometric Determination of 2.00 > 3.00 4.00 ML. TITRANT ADDED (0.09927 N.) FIGURE X B M tr* a t r 1 o O Cr. tr* >-3 GO 0 Precise.Determination of the End Points f o r Figure X A 1 END POINT I END POINT i 0 1 0 20 3 0 40 50 60 7 0 8 0 9 0 1 0 0 VOLUME ADDED NEAR END POINT (1 DROP = 0.0148 ML.) FIGURE XI Curves Obtained from the Attempted Potentiometric Determination of Copper ( I I ) , Antimony ( I I I ) , and Tin (II) Acetates i n Acetic Acid With Perchloric Acid. The Curve Obtained Upon Adding Perchloric Acid to Acetic Acid Alone i s Included f o r Reference. 1.00 00 3.00 4.00 ML. TITRANT ADDED (0.09927 N.) 7 0 0 -6 0 0 5 0 0 -460L 0 F I G U R E Curves Obtained From the Attempted Potentiometric Determination of Bismuth (III) and Uranyl Acetates i n Acetic Acid With Perchloric Acid. The Curve Obtained Upon Adding Perchloric Acid to Acetic Acid Alone i s Included f o r Reference.. 1 . 0 0 2 . 0 0 3 . 0 0 l i L . TITRANT ADDED ( 0 . 0 9 9 2 7 N ) 4 . 0 0 750 FIGURE XIII A Potentiometric Determination of Approximately Three M i l l i e q u i v a l e n t s of Barium Chloroacetate i n Acetic Acid With Perchloric Acid, The Curves Obtained f o r the Other Chloroacetates Tested Were Similar to the Curves Obtained f o r the Corresponding Acetates :ND POINT I o S 7 0 0 H t-1 f M < O IS. 650 6U0 26.00 27.00 2&\00 ML. rp T rp X J . RANT ADDED (0.09927 N) <4K M tr 1 a M o <! •x; O CO t~< >-3 CO 14 12 10 2 0 0 FIGURE X I I I B P r e c i s e D e t e r m i n a t i o n of the End Poi n t for Fi g u r e 'XIII A 10 END POINT I I L . 20 30 -V&0 50 60 70 80 90 100 VOLUME ADDED NEAR END POINT (0.0148 ML. = 1 DROP) 110 

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