76
D. R . I3RIGGs
R e c e i w d . t u V i i s t I d , 1047
Certain deviations from the equivalent conductivity-concentration relationships Tvhich are characteristic of ordinary electrolytes in solution have been observed in the case of soaps (4)and dyestuffs (51, in that, as the concentration of such an electrolyte is increased, the equivalent conductivity passes through a minimum value follon-ed by a sometinics sharp rise and finally by a slow decrease again. electrolyte 11-hich shoivs this phenomenon is characterized by a tendency to form aggregates in solution at some fairly definite minimum concentration, belon- which it exists in a state of ordinary ionic solution but above 11-hich it exhibits to an increasing extent the properties of a colloid. This concentration is designated as its critical concentration, and appears t,o coincide closely 11-iththe concentration of the initial minimum in the equivalent conductivity curve. The increase in equivalent conductivity n-hich accompanies the aggregation of the micelle-forming ion, as was pointed oiit by NcRain many years ago, is clue to an increased mobility of the micelle over that of the unaggregated ion a i d results from the decreased v i m m resistance encountered, per unit of charge, by the aggregate. Mobility ineawrements on the micellar ions have confirmed this hypothesis. The final decrease in equivalent conductivity occurs as the effects of increasing ionic strength ultimately overshadow the effects of the aggregation process. Substances which exhibit this property are hetelopo!ar in composition, containing one residue in the molecule ivhich has a Presented a t t i w T\\-nir>--tirat Saticinai C'olloici S p i p ( ~ s i u : n\vliic.h . \vas held under t h e auspices of t h e Division of Col!oirl Chemistry of tlic Anieric:xti C'liemic,al Society at Palo Alto. Ca!ifornia, J u n e 1s-20, 1947. I P a p r r So. 2352, Scientific J o u x a l Series, llinnesotn A\gricult~i!,:ilExFeriment S t a t i o n .
Ion- affinity for water and which is rcsponsible for the tendency of the molecules (or the i m s containing the non-polar residues) t o aggregate. These substances act as colloid electrolytes only hy virtue of their tendency t o form aggregates from simpler ions in solution, and it may ne11 he that their conductivity properties, as colloid electrolytes, are overshadon-ed hy the changes irhich accompany tlie proce.Gs of aggregn tion through which they liecome colloid electrolytes. I n order t o stud>- he conductivity properties of colloid electrolytes, as such, it ~voiild appear t u he desirable to invcstigate nisiterials \i.hich are colloid electrolyte,q, not cgatioii in solution, hut ivhich exist as colloid electrolytes e1.m at the lon-est possible concentrations t o be employcd and xvhich do not, have any obvious tendency to change in their degree of dispersion or aggregat,ion n-ith change in concentration. The potassium, sodium, and lithium salts of purified and electrodialyzed gum arabic (1) appear to be examples of such colloid electrolytes. Khile the molecules oi the gum are not entirely homogeneous as to size or composition, they are nevertheless of colloidal dimensions, having a molecular n-eight of about 300,000 =t50,000 '6) and an equivalent n-eight of about 1200 f 30 ( I ) . The gum appears to be predominantly hytlrophilic in character, heing soluble in all proportions in n-nter;it is not surface active at a water-air interface (although it collects at a quartz-n-ater interixe). Its solutions in miter show true viscous flow up to high concentrations, and it exhibits no tendency to form gels. On the basis of these evidences of a strong oi-er-all affinity for ivater, it \vould seem probable that gum arabic in water is molecular and that any tendency for it to aggregate in w t e r solution ~vouldhe negligihle. This pnper reports studies on the conductivity and mobility properties of these ealTs of gum arabic. Figure 1 i;hon.s the equivalent conductivities (corrected for the conducti1-ity of the m t e r i of the potassium, sodium, and lithium salts of gum arabic in water at 25'C'. plotted against the equivalent concentration of the solutions. The conductometric data shomi in graph form in this figure v-ere obtained with the usual TT-heatstone bridge arrangement, using a 1000-cycle I.C. source. The data are a repetition of conductometric data previously reported (I) for these colloid electrolytes. hut correction has been made in the present instance for the conductivity due to the solvent. I t will be noted that, through the concentration range studied (0.00053 .I-to 0.042 S,corresponding toa \\-eight concentration range of 0.625 g. to 50.0 g. per 1000 g. of water), no iiiinimum in the curve is detected h i t there is an irici.ins< in tlie equivalent conductix-ity to a maximum follon-et1by a slon- decrease as the concentration increases. That the observed cowse of the equivalent conductivity-concentration curve cannot be explained by an increase in the mohility of the arabate ion is indicated by direct determinations of the niohility of the colloid. Table 1 gives mobility data at 25'C'. obtained on sodium arahate at pH T.0 a t various concentrations of the eolloici electrolyte n-ithout and with the ncldition of various amounts of sodium chloride. These data nere obtained by the microelectrophoresis tech-
n q u e (2), using qu:ii t z pniticles upon uliich t h , ; ~ i : i L , ' t i ~11~1' . ailsorbcti. Flguri. 2 shows a graph of the-e data in n hicli the mobility i. plotted agnins! ttlc logarithm of ionic. strengtil. Fire additional points are included in figure 2 which n-ere ohtained 11 ith the 'fiseliiis macroelectrophore-i; mcthod at 0.5%. 111 itcetate
buffer at pl-1 ti.2 and at higher ionic strengths than those for most of tlir microelectrophoresis datu. (The values for thcsc mobilities ;IS shou n in figure 2 have been recalculated to 25°C'. upon the asslimption that the mobility difference at thrse two temperatiiw-: 11 oulcl I)c n function only of thc temprJrature differences
in the x i ~ c o ~ i of t y the solIwit. '1'hese data shon a imooth continuation of the data obtained by the mi(.roelectrophoresis method and sei IY to confirm the correctness of the latter. even though in this casc the gum n-as adsorbed on quartz particles. I t is of interest to recognize that the effect of the ionic strength up()l1 the mobility of sodium arabate i*, nithin esperimental error, the bame whether it derives from added sodium chloride or from the sodium arabate itself, where it is assumed that the arabate ion, even though carrying a number of charges (200 or $0 per molecule), act< as a corresponding number of monovalent ions insofar ns the contribution to the ionic strength is concerned. Apparently the point charges are so remote from each other in this colloid that they can act as mononlent charges in thi? re-pect. This is proba1)l~not the case with all
F I G . 2 1Iobility (cm volt ~ I Y . of qodiuni arabatc at v a i i o u i ionic strengths ('1-1 Sniall circles iefer t o d a t a o1)talnrd b) iiiicioelectropliorcsis niethud at pH 7.0 on solutions of sodlui:i arabate without and n i t l i added sodium chloride Double circles refer t o data obtainet] by TIsplius cleciropliorcsi- m e t h o d o n wlutione of sodium aIabate in acetate buffers of 1311 6 2
colloid electrolytes. Such a situation, hon ever, is quite fortunate in the present instance, since it makes it more easily possible t o observe relationships which are unconiplicated by the strong ionic strength effects of polyvalent ions. X-hile these mobility observations have been made on the sodium salt, i t has heen found that there esi3.t only insignificant differences lietn-een these values and those of the potassium or lithium salts at equivalent concentrations of colloid electrolyte. I n the calculation which follon s. it is assumed that the mobilities of the colloid ion are independent of the cation for these three cations, I t is evident from the data in figure 2 that the mobility of the colloid ion decreases continuously with increase in concentration of the colloid electrolyte and that the rise in equivalcnt conductivity illustrated in figure 1 cannot be explained on this basis.
The obserwd ecluivnlent conthirtivity oi the tlralxrte .a!t in solution. -iobsd. X c 1 3 \\.hew cy is the coii(lrictomctric~ac.tivity of the araliate salt. iAr is the equivalent conductance of thc aratmte ion, ~undA,. is the equiv:dent conductance of the cation (lithium, pota;isium, or sodium in these esperinients). Values for hAr can lie cdculatetl from the measured niohility of the colloitl, i.e., XAr = nzF, Ivlzere m is the niohility in cm.','volt sec. amd I.' = 96,500 coulombs. Values for A, are not obtainable by dirert nieasiircment on the arabate salt liut, since the contribution t o the ionic strength of the cation oi these salts is denionstrated to be so near t o that contributed by any added chloride salt of the cation, it may be reasonably assumed that X, is equal t o that of the cation in the corresponding chloride solution of an eqiiivalent concentration. Such values of X, have been calculated from the equivalent conductivities, at 25'C., of potassium = ct(XAir
+
T.II31,E 2 . l l r i i d i t y of a i a h u i e i o l i , c q j i i i x l e r i l c.o~itirictii~ilies of c o n i p o i i c i ~ t ions a n d of [ h e po:assiurn, sodium, and lithi(riti a r a b n t e s . n t r d l h c ~ ~ i t i ~ I i i c i / ~ uclil,iiies t ~ i ~ t r i ~of these salts i:i nqiteoris scilutioiz trt i,rr!,ioits ccincetiirniions ut 26'C'.
cliloride, sodiiim chloride, ant1 lithium ( ~ h l o r i das~ given in tile 1,itcl.ocitioiial Critical Tables (31, using thc relationship, for potassium chloride for example,
11-here AI. R
J . P ~ >C'hein s 38, h67 (1934). , .411al Ed 12, 703 (1940 C i L t l t n i 7clhlts 101 \ I ;\Ic(;iau-Efill lhml'
1 3 ~ 1 ~D ~ 3R. : Irltl Cng Chrrll Iniei,intluiiu2
('onlprlJ,
1923) ( 4 ) l\rcl3$1\ J 1V Tiitris F:tretiav bo( 9 , 9J (l(Jl3) ( 5 ) ~ I O I I I I F ..JT I, ( O I I I J I$ R O R I \ - O T , < ~ \ I ) I ~ ~ I I I (I , I I 31, 120 119x5) I I 13 . Ti,irli I ' a l i d n ~ hoc 31, 1.36 (1083) f
.
I J M, SCW 1 (irk
Tian< 1 ' ~ i : d a y S w
