Behavior of Acid-Base Indicators in Acetic Acid System TAKERU HIGUCHI, JOSEPH A. FELDMAN', and CARL R. REHM School o f Pharmacy, University ofWisconsin, Madison, W i s .
This cquilibriuin coiistaiit ciiii be determined by determination of the ionization characteristic of the indicator and of the acid
Results of spectrophotometric and potentionletric studies on 13 indicators in anhqdrous acetic acid suggest that the color changes which occur during titration are governed largelq by the relative concentration of the base present and its conjugate acid. For stronglq basic indicators such as ethyl red, acid colors can be elicited, furthermore, in acetic acid by neutral salts such as sodiuni perchlorate and sodium sulfate. The following indicators were studied: ethyl red, pinacj anol, 4diniethylamino-4'-nitrostilbene, 2 -nitro- 9 - (4'- di methylaminobenzal)fluorene, 4 dimethylan~ino 4' sulfamylazobenzene, quinaldine red, 4'-dimethqIaniinobenzalrhodanine, rn-nitro-N,,V-dimethylanilinc, Rrilliant Cresq 1 Blue, 1-naphtholbenzein, Nile Blue i, Sudan 111, and Sudan IV. Their response was determined to the following buffer systems: ephedrine acetate-ephedrine perchlorate, sodium acetate-sodium perchlorate, sodium acetate-sodium sulfate, antip) rine acetate-antipyrine perchlorate, urea acetatenre3 perchlorate, and perchloric acid, all in acetic acid.
-
individually. I n solvents of low dielectric coilstwilt, w c h as acetic acid, the (.omparable equations are somen!ittt more comples hecause of thr w r y strong tendencies of the ionic. coniponeiits to form ion pairs :tiid mole highly associated spwirs. The strongest acids :tnd strongest bases in these systems diwoci:itP relativc~lylittle, even ;tt very low concentrations. The present studies sugglst t h t the bcth:tvioy of witl-imrc~inilicators in acetic acid is I w t w p r w ~ n t c dliy the gi.iiiaral re:iction
- -
SH;\
13)
where 5 is the solvriit or sonic othcr protopliilic. base present; SH-4 is the interaction prodiict of a strong acid, 11.4, and S; IHhc is t,hc acetate form of th(. indicator, eshibiting its hast, c-olor; and IH-1 is the arid forin of the indicstor resultiiig from interaction vith 11.4. Tht,refoi,e =
D
ESPITE t8hevo!riminoiis number of publiciitioiis relittcd t o the use of acetic acid and othcr nonaqueous solvents its titration media in analytical determinations, investigation of indicator behavior in these syst,ems has been relatively limited. Experimental data are given here pertaining to the behavior of 13 indicators in acetic acid in t,he presence of a number of different csations and anions.
+ 1H.4~2 1H.i 4-SH.ic
This follo\r.$, i,ariier ( 7 ) :
TI
here K a
=
iii
(1H-i) (SHAc) (SHA) (1H.1~)
(4)
p a r t > 1i)gic~allyfi.oin relationships prestwted
dissociation coustant of the acid
Kg = dissociation constant of the base K H . =~ autoprotolytic ~ constant of acetic acid K,b = dissociation constant of the resulting salt
THEORY
.llthough both the base form of the indicator (1H.l~)and the acid form (IHA) are written as products of acid-base reaction, the extent of proton transfer must be sufficiently different qualitatively to produce t,he observed color difference. Perchloric acid appears to be able to effect a substantial or a nearly coniplete transfer, whereas acetic acid probably forms only a hydrogen bond with the protophilic center. The proton transfer froin a strong acid (HA) to the solvent or other protophilic base (S) to form an interaction product (SHA) is probably limited if S is ail acetic acid molecule, but the transfer is nearly complete if S is, for example, an aliphatic amine. I n the latter instance the over-all reaction can be written more simply as
Because the physical chemistry of color changes of indicatoi 9 in nonaqueous solvents such as acetic acid is fundamentally different from that in the more fainiliar aqueous systems, it may be well to review very briefly the basic differences between the two systems. In water, because of its high dielectric constant, all true salts c:tn be considered to be both totally ionized and dissociated. For this reason the color of an indicator solution is determined primarilj by the hydrogen ion concentration or the p H of the ,vstem. Thus, for indicators such as methyl orange, these relationships exist:
+
( H + ) ~ , o I Ft I H
+
+
IT4.%- IHAc ri 1I'Ac-
and log
I = - log €1+H*o - log KI IH = p H - PKI
anti
K = (11'4c -) (IHA) (11+.4-)(IHBc)
+ -
(1)
\\here I represents the base form of the indicator; I H + , the acid form; and K I , the dissociation constant of acid I H + . For aqueous solutions containing a weak acid and its salte g benzoic acid-the following relationships also exist. ~
HBz
+ I e I H + + Bz-
The equilibrium constant for this reaction is
K = (IH+) (Bz-) = -Kr (HBz) (I) Ka J
Present address, Duquesne University, Pittsburgh, Pa.
+ IHA (5)
m-liere M + A - is a salt of a strong acid and a strong base, and i\l+Ac- is the corresponding acetate salt. A protonated ephedrine cation, for example, can play t,he role of lI+ in the above equation. Because cations of this type are not basically different from elemental cations such as S a + , I
