the solvent in bulk, such calculation underehtimates {. The potential at the surface of a charged sphere, .,\.here D aiid 7 are the solrent dielectric constant nhich is related to [ but always greater than it and ~iscosity,and ni, zi and t’i are the concentration, valence and mobility of the ions in solution. because the shear plane in electrophoresis is located If n e aisume that all of the ionic mobilities are a t zome unknown distance from the surface where equal, then according to Booth (see page 1966, the potential is lower than a t the burface, can be calculated from the charge and the radius of the equation 5.10 of Booth’s publicationg) virus particles by means of the Debye-Hdckel theory, or with greater accuracy by the method of Loeb, Wiersema and O ~ e r b e e k . ~ The ’ latter method is strictly applicable only to symmetrical aiid q3* recliiccs to electrolytes and thus is not completely valid for L__M ‘ 2: 11,%,2 phosphate buffers. T?k. For the yirus particle in 0.002 ionic strength, 7111 G 9, huffers { e / k T v a s fouiid to have a value of As caii 1)e seeii in Table 111, tlie \ d u e of tlie about 1 when { was calculated from electrophoretic portion of dz which is iiidependeiit of conceiitramobility. When surface potentials computed from tion of SBll\‘ charge were used as estimates of {, f e / k T for the q3* V3*(6 ) kame conditions had values of 3.2 and 2 3, dependn-as fouiid by coniputatioii to be greater than tlie ing on nhether the computation was made for uniportion of d, n-hich is depeiideiit on virus coiiceii- uriivaleiit or di-divalent electrolytes. In view of the uniatisfactory state of electrokinetic theory, tra tion oiie can say no more than that { e l k T is of the order q *iSlpaS,*2(b ) b -2 of magnitude of 2 for the virus in 0.002 ionic strength phosphate buffer a t pH 6.9. The values This result is iiiconsisteiit n ith the experimental a t other pH values n-ould be somenhat more or data show11 in Fig. 1. romewhat less, depending upon valence or nioThe zeta potential is employed in the Booth bility. charge effect e q ~ a t i o n . The ~ zeta potential can he Since the values of dz for the virus protein in calculated from the electrophoretic mobility by 0 002 M phosphate bilffer a t pH 6 9, obtained hy mean? of the Henry equation” or of the more adding the numbers in the trio right colunini of e l a h r a t e 1300th electrophoresis Theke Table 111, are about - 10 X low3,the values of and other equatioirs assert that the moMity is s’ computed from the Booth equation are only proportional to D{,lq. The mobilities recorded in ahout 4y0 1e.s than s. Boothg pointed out that the preient study fall in the range in n hich tlie his equation predict. charge effects that reduce s Booth equation reduces to the Ileiiry equation. by only a few per w ! t . ’T’huh, the Booth equation Because of the fact that 7 in the iicighhorliood of a underestimate. tlic o l w r \ ed charge effect. charged p:irticle is prohahly greater than that of (2G) F. BooLli Proc l t o b Ssoc ( L u n d u n l , A203,514 (1950).
(?7) A L.Loch T’ 11 \\ i r i ~ e i i i d dnd . J. T n G 0 t o in “Ckctrophorcsis ’ tdited hv 11 Bier 1059 p
KISETICS ,%SD M ~ ~ C I ~ I ~ ~ X 0 1 I7 STHE M A41R OXIl>,4TIOS OF THE DITHPOXITE I O S (S,O,q-) IS ;iQUEOUS SOLUTIOS BY ROBERTG. RISKER,THOMAS I?. GORDON, DAVID31. XASON,’ ROYR. SAKAIEA ASD
1\-ILLIAM
H.
CORCORAX f
‘ h i i i i r a l Etioiticcriiig Lahoi.atoy!j, (‘aliJornia Instilute o j l‘tchnnlog!j, l’trandciin, (’nlifornia Received Sepleniber 16, 1960
study of the :Lir oxitlatior1 of sodirini ditliioiiite \vas condiirted in aqrieoris solution which was 0.1 JI i n sodiniii l i ~ droside. T h o conccntr:rtion of t,hc. dit~hiciriitcn-as mc:isurcd :is i t funrtiori of time a t 30, 40, 50 and 60”. Initial conccntraUnder conditions in wliic-li diffusion of nir was not controlling it was found tioris v:tricd from 5 x 1 0 - 3 to 20 X 10-3 that t,he oxidat,ion was h:tlf ordor with rcspcct to ditliioriitc ~ 1 1 dfirst order n-it,li respect to molecular oxygen. The half ordw mechanism ~q-:ts:tttril>iit,;thlet o the prcsenrcl of the SOZ- radicd ion as ai1 intermediate.
574
It. G. RINKER,T. P. GORDON, D. AI. MASON,It. It. SAKAIDA ANI) W. 11. COIKORAN Vol. 04
order rate constants but found that they drifted over a wide range. Bassett and Durrant3 in 1927 studied the reaction between dithionite and molecular oxygen. Their explanation of t'he oxidation mechanism was based upon arguments that dithionite exists in three isomeric forms. Each isomer was believed to decompose in a may such that all of the observed reaction products could be explained by one or more of t'he mechanisms. Although Bassctt and Durrant's work is of considerable value as a compilation of experimental observations of dithionite reactions, subsequent work has invalidated their interpretations. Nicloux4 in 1933 was interested in determining the over-all stoichiometry of t'he react'ion between dithionite and oxidizing agents of varying strengt,hs. With a comparatively weak oxidizing agent such as silver ion, the dithionite was oxidized to sulfite. With molecular oxygen, an equimolar mixture of sulfite and sulfate was formed. IGnally, with a very strong oxidizing agent, sulfat'e was the only product. Xu experimental details were given. The results of Lynn5 suggested that the atmospheric oxidabion of Na2SzO4proceeded according to a first-order mechanism with respect to dithionite a i d t,hat the rate increased with temperature. He stated that the over-all stoichiometry \vas descrihed by the equat8ion Sa?S.O?r
+- + H,O 0 2
=
SaHSO4
+ XaHSO:+ ( 3 )
-411 examiiia,tioii of his data at 50" showed that the oxidation rate increased with an increa,se in air flow which meant that the diffusion rate of oxygen was a contribuhing factor. Without correcting for diff'usion his results for the actual oxygen-dithionite reaction were, therefore, only approximate. His experiments did show that in the presence of 0.1 M sodium bisulfite the oxidation rat,e was extremely high. In the presence of 0.1 ill sodium hydroxide the rat'e was inhibited, and the reaction proceeded smoothly. Experimental Apparatus.-The air oxidat,ion of the sodium dithionite was carried out in a glass reactor with a volume of approximately 600 cc. Air %-asbubbled a t a rate of 2500 to 3000 cc./min. into the reaction mixture near the bottom of the reactor through a medium-coarse glass frit which could be inserted or removed through a standard-taper opening on the reactor. During the reaction, the frit was always immersed in the liquid contents. Before the air entered the frit, it was passed first through a,ri illundum-stone trap to remove suspended solids and aerosols, nest through a molecular-sieve dryer, and finally through a teniperat,ure-conditioning coil. The air flow was measured by means of a wet,-gas meter placed on the downstream side of the re>tct,or. Mixing was accomplished by means of a glass stirrer having two impellers. The st,irrer was supported and sealed through a mercury-in-glass bearing. Stirring speeds could be varied up to 1100 revolutions per minute, which was the speed used throughout the experiments. The react,or and the temperature-conditioning coil were thermost:it,ed in a 28-liter wat,er-bath, which maintained the temperature within i 0.02". Heaters in the water-bath were energized by an electronic-relay circuit which in turn received its signal from a mercury-expansion switch immersed in the water. H. Aansett and R. 0.Drirritnt, J . Chena. Soc., 2, 1401 (1027). (4) 31. h ~ i c l o u xC, o m p t . F e n d . , 196, G l f i ( 5 ) 9. L m n , 1'11.U. T l w r i s , California 11
Samples were taken from the reactor through an opening in the top. A syringe needle or pipet was inserted into the reacting mixture, and the sample then was removed. Commercial-grade nitrogen was the source of supply for the oxygen-free atmosphere required in preparing the sample of fresh sodium dithionite to be charged to the reactor. The nitrogen contained approximately 0.01 yo oxygen by volume and this amount proved excessive. Hence, most of the oxygen was removed by passing the nitrogen through two ' chromous scrubbing towers in series. Each contained 0.2 h chloride solution in 1.0 -V hydrochlorir acid. Arid and water vapors were subsequently rrmoved h y passing the nitrogen through a 0.1 M sodium hydroxide solution and then through a calcium chloride dryer. Concentration of Dithionite as a Function of Time.-Chemically pure sodium dithionite (Baker Chemical Company, Batch S o . 3712, Lot No. JTB6113) \\-as the starting point for the source of dithionite ions. In preparing the dithionite for a typical run, approximatelv 25 g. of the powder was placed in an oxygen-free flask into which a stream ol oxygen-free nitrogen had been passed for several minutes. Then, approximately 100 cc. of distilled water, free of oxygen and carbon dioxide, was injected into the flask through ; rubber serum-bottle-cap. The mixture was heated to 60 with constant agitation until a saturated solution was obtained. With great care to avoid contact with osygcn, a 60-cc. sample of the saturated solution was vithdrawn with a syringe and injected into a side-armed test-tube already filled with nitrogen. The tube with its contents then was cooled to 0" in an ice-bath while maintaining the pressure of the nitrogen constant at slightly above one atmosphere. At the lower temperature the liquid became supersaturated with sodium dithionite. K i t h sufficient agitation, crystals of Na&04.2H20 were formed. Again, with great care, practically all the liquid was removed froin the tube leaving the white crystals settled a t the bottom. Approximately 4 cc. of distilled water a t 0" and free of carbon dioxide and oxygen was injected into the tube to wash the crystals. This washing process was repeated once again to obtain reasonably pure crystals. The final saturated solution with a volume of approximately 20 cc. was used to supply thr reactor with an initial concentration of dithionite. Before injection of the dithionite, the reactor was filled with 505.0 cc. of 0.1 iLl sodium hydroxide solution. The reactor was then immersed in the water-bath and allowed to reach steady conditions of temperature, stirring rate and air flow. The time of initial air flow was noted in order to account for evaporation losses from the reactor. At steady conditions, 5.00 to 15.00 cc. of the saturated dithionite solution, depending upon the initial concentration desired, waq injected into the reactor. Since the delivery times of largr synnges are of the order of several seconds, the time a t whicsh half of the sample was injected was talien :is the initial timc of the reaction. Wthout delay, 1.00 or 2.00 cc. samples were removtd from the reactor with a calibrated syringe and injected into 150-cc. flasks containing a mixture of 50 cc. of 0 1 31 potassium hydrovide and 15 cc. of methyl alcohol Also, the flasks contained a nitrogen atmosphere M hich was maintained during titration. The solutions in the titration flasks were stirred magnetically. The time at whivh half the sample from the reactor had been injectrd into :t titration flask was recorded as the injertioti time. The concentration of dithionite in thc. titt:ition flashs w:ts determined h r titration with a stand,trdmltt .boiiwiis solution of methyleni blue having a concentration i;i the r:tnge of 9.0 x 10-1 totiiosulfate,a 10 to 25-cc. s:tmple wits buffered with a solut,ion which was 0.5 ill in sodium acetat,(?and 0.4 Af in acetic acid. I'sually, a volnmrt of the buffer solution equal t,o the volumtl of the sample w : i b added to give a const.ant pH of tiI)oiit 5 . Using a starc.h end-point, tlic: solution was titrated \\it11 a standardized solution of approximately 0.01 A' triiodide. I n case of escess addition of triioditle, back-titration was done with a standard thiosulfate solution which was approximately 0.01 in thiosulfat,i,. In the iodometric analysis for only t,hc tliiosulfnt,ction, a 10 to 2>5-cc. sample was buffered t,o a pEI of abut 5. To this was added a volume of 37 wight, r; formaldtshj.tit: in watrr cqual to about half the volume of the s:tmpicL. The purpose of the formaldehyde \vas to form :i complex with tthe bisulfite according to t,lie reaction
ti C ' :-O
+ H S ~ ) , ~ X+ : ~ ~ ' H : ( ( ) H isoax:t
(-1 )
H r .
1lic mixture \vas stirrrtl for 1.5 miriu tt room teinprratiirc. to allow sufficient t.ime for t,he complexing to occur. The unreaeted thiosulfate then was determined with thr, standard triiodide solution. The use of baririni chloride in a titration riiclthocl t,o d e t w mine sulfate was suggested in the lit,eratrircb by Fritz and Free1and.Q This inet,hod depends upon :t color change in the indicator Alizarin Itrd-S whirh acts as an adsorption indicator in the presmce of a barium siilfat'e precipitate. I n solution the alizarin anion was yellow; hut, on the surface of SUYpended barium sulfate and in the presence of excess barium ion, the alizarin complexed with the barium ion to give a red color to the suspended solid. In using the proc,edure outlined by Fritz and Freeland, however, the method was limited to concentrat,ioris of sulfate greater than 20 X M. I t was necessary, then, t,o modify the procedure in applying it to the concentrations of sulfate down to 4 X M . After several tests, it was noted that the limit of detection of sulfat,e was determined by the :$mount of prec,ipitate available for adsorption of alizarin. Too small a quantity of precipit>ate,although colored red n-it,hcomplexed alizarin, was not, perceptible tshrough the yellow color of the dizarin in solution. Therefore, by adding a quantity of semi-colloidal ba,riiim sulfate suspended in met,hyl alcohol to the titration mixture, it was possible to detect the end-point with an accuracy of 1 or 2%. A further improvement in detecting the end-point was to pass an intense light beam through the suspension during the t,itration. This aided in bringing out color changes more sharply. To determine the sum of the sulfite and sulfate roncentraetion by the above met,hod, i t was first necessary to oxidize the sulfite to sulfate by means of the triiodide ion. The amount of t.riiodide needed was det,ermined by a starch endpoint; but there was some uncertainty as to the effects of the starch on the subsequent sulfate titration. Therefore, i t was decided to avoid using the starch by first running a blank t o detmnine the exact triiodide reqiiirement. and t,hrn adding ~~
~~~~
~
~~
~
~-
(!I) .J. S. k r i t z and AI. Q. I'ruslaiid, Ariol. C ' h ~ m . 26, , 1533 (1954).
All the standard solutions used 111 t h c ~forcgoirig an& 5 e ~ mere prepared and standardized according to the procedures outlined by Swift.lo
0.024 I
0.020
Results
c
\ m
3
2
-z z
9
RUN
30C
0
RUN
300
-0
RUM
30E
0.012
0
v
-g
0.008
.3
13 3
m
0.00 L 0
200 .300 400 500 GOO Time 6, see. Fig. l.-lAtliionite concentration us. time in air oxidation a t atmospheric pressure and 30". 100
(J
0
RUM 3 0 F
0 RUN
300
?
RUN 3 0 H
1 b
RUN 3 0 J
0 0126 0 0126 0 0193 00193
-. . 2
-
E" 0.012 i
g 0.010
."a, g 0.008 L3
.I
5
IS
0.006 0.004
60 80 100 120 140 Time e, sec. li'ig. 2.--IXtshionite concentration us. time a t 30" using pure oxygen a t atmospheric pressure.
0
20
40
thc same aniount in the absence of starch to a fresh samplr buffered at a pFI of ahout 5 as specified above. Following the oxidat,ion of the sulfitc, tho buffered solution W:LS acidified to a pH of about 3.5 with 20 weight% aretic acid. The volume of ac.etic acid required was about eqiial to the voliims of the original sample. Then methyl alcohol conhining a semi-colloidal suspension of harium siilfat,e was added in an optimal amount equal to 38y0 by volume of the final miut,ure. With constant stirring, 90% of the estimated 0.1 AT tlariiim chloride requirement was added rapidly followed ty 3 to 5 drops of 0.020 weight % a.lizarin-red solution. The final titration was performed slowly rrith an interval 013 to 5 SN. I)et.wern drops of 1t:wiiirn t:hloritfe sohition. Finally, in determining the concentration of sulfat,e alone, t h e sulfite W : ~ Scomplcxed h y addition of forni:ildefiydc. Thc: siihscqucnt barium chlorido titration wts riot affcctctl nilvcrscly by the prcscncc of thc formnlilchydc.
Rate Studies.-Itate data obtained a t 30" are shoI+n in Fig 1, in which the concentration of dithioiiite in moles per liter is plotted against time in seconds. A tabulation of these data along with results at 40, 50 and 60" is available." Initial concentrations of dithionite ranged from 5 X to 20 X mole/liter. The rate of change of concentration mas large in the initial stages but decreased continuously and rapidly as the concentration decreased. I n the determination of the order of the reaction with respect to dithionite, the oxygen concentration was held constant for a given teniperature. Only for those experiments in mhich the order with respect to oxygen was to be determined was the oxygen concentration varied Still maintaining atmospheric pressure, pure oxygen was used as a replacement for air to gix-e noni1n;illy a fivefold increase in oxygen concentration. For an assumption of Henry's law, the concentration ratio of dissolved oxygen in equilibrium with tho pure gas at atmospheric pressure to dissolved oxygen in equilibrium TI ith air a t atmospheric pressure would be 4 8 This saim iutio wa5 23sumed for the w r y dilute dithionitr wlutionq in M hich the total gas pressure was at the atmoipheric level. For the equipment used, this method of varying the partial pressure of the oxygen in the gas phase was preferred to a method in nhich the total pressure on the system mould have becn varied. It was found experimentally that the initial rate of the reaction was increased by a factor of 5 when pure oxygen was substituted foi air, all other conditions rt maining the same. Therefore, the reaction was believed to be first order nith respect to molecular oxygen. Figure 2 shons rate data obtained a t 30" when pur(' oxygen was bubbled through the y s t e m with tlic total pressure being atmospheric. Those results have also been tabulated." The figure m!y be compared with Fig. 1 for the case in which air was used. Two procedures mere available for analyzing the data in the determination of iiiitirl rates. An initial concentration could have becn cstablished by use of solubility data obtained a i cxplained in Experimentul Procedure. Also n i l initial coiicentration could have heen obtniiied by exti:ipol*1' t 1011 of the cnnceiitratioa-tinie curve luck to aei o time Inasmuch as qamplinq of the reactor was ubuallj hcgiiii only 30 t o 50 v c after zero tinic, the uw of the extrapolation technique \T as qatiqfnctorv. Tlic o\trapolation method was prcfcrred h e c a u i e of the need for exact control of equilibrium conditions in the event that the solubility data 11 w e usrd. (10) E. H Swift, "Systrmztlc Qliantitatl\e 4nai>.1- '' PranticcHall, New Tork, h ' Y , 1939 (11) R G Rinker T. P Gordon, D. N. AIawn, R R SiEfiida. a n d TT €1. Corroran, Ani Doc Inst., Washington 25, I ) C 1)ocriiilsnt 4 L ( , , , \ O t t i l l . I f O L i l l l l f I l t I 1 b J Lr . i ( < ~ l i C < b\ l 11t11,g t h t . t , l l l l (1954) < I ~ , i ~ n tnni it n h i r ~ n bdy reinitting 91 50 for idiotogJiint> or $1 75 f o r ij 1nn1 ~ n i c r o f i l u l \ i l \ a r i c r 1) t3 inrrit i i r1 q iircd Aldht c l i t r-.I or i c i o n e ~otdcrs 1) L> ~ l , l ( t o
GingLsu
Chef l'hutudi 1~11c.itiun6~
!11
>k5 in the above scheme. ?I supgection steady-state concentration of intermediates necd 8204" 2802not be postulated in this hypothetical analysis. kz Acknowledgment.--Funds for this research on SO?- + O? complex B --+ SO^ 0:dithionite chemistry were mainly provided through k, a general grant from the Shell Companies FoundaSOL - + 02- --+ SO.! Or' tion. In addition, fellowship support was given by X.4 the Union Carbide Corporation and the RIonsanto 02" -t H,O +HO?OHChemical Company. The help of d l is gratefully k6 acknowledged. HO?- + SO?" +604" OH(23)
+
*
e
+
+ + +
REACTIOKS OF POLYSOAPS WITH CHLORIDE ASD RROTtITDE IOSS' BY DIETRICH WOERMANN A N D FREDERICK T. WALL Soyes Chemical Lnboratoru, L'niversili/ of Illmois, Urbana, Illinois Recezved Oclobei 3, 1969
One polyelectrolyte and five polysoaps with chloride and bromide counterions have been prepared from poly-4-vinylpyridine by quaternizing 0, 4.9, 5.8, 8.1, 13.9, 31.8% of its nitrogen with n-dodecyl groups and the remainder with ethyl groups. The solution viscosities of these chloride and bromide polymers have been studied in the presence of 0.0223 N KC1 and 0.02233 N KBr solutions, respectively. Characteristic differences exist between the reactions of the chloride and bromide forms of the polysoaps and polyelectrolytes. The results can be explained by Strauss' concept of a "critical dodecyl group content" of the polysoaps, corresponding to a critical separation of the soap groups. When the critical content is exceeded, by increasing the dodecyl group content, the soap groups exhibit a tendency to form aggregates and the intrinsic viscosity drops considerably. The critical dodecyl group content is analogous to the critical micelle concentration of ordinary soaps. Corresponding to the differences between the critical micelle concentrations of ordinary chloride and bromide soaps, the critical dodecyl group content of the chloridp form of the polysoaps is greater than that of the bromide form.
Introduction The gradual transition from typical polyelectrolytes to polysoaps (ie., polymers to whose chain structure cationic or anionic soap moleculesare chemically attached) has been investigated extensively in recent years. Especially careful study has been given to systems of polyvinylpyridine derivatives prepared by quaternizing some of its nitrogen atoms with n-dodecyl bromide and the remainder with ethyl bromide.2s8 Viscosity measurements of aque(1) This work was supported by a grant from the National Science Foundation. (2) U. P. Strnriqs, S. I,. (:c1stikld a n d E 11. Crook. Trim .JOURNAL, 60,577 (i9.w).
ous solutions of these polymers show that there is tin analogy between the properties of polysoaps and ordinary soaps. The results can be explained by Strauss' concept of a "critical dodecyl group content" of the polysoaps analogous to the critical micelle concentration of ordinary soaps. The purpose of this study was to investigate by means of viscosity measurements the different reactions of the chloride and bromide forms of a typical polyelectrolyte and of five polysoaps, prepared by quaternizing some of the nitrogen of poly-4-vinylpyridine with n-dodecyl bromide (0, 4.9, 5.8, 8.1, (3) A . I. Medalia, H. H. FrPrdnian and S. Pinha, J. Poluncrr SI1 , 40, 15 (1959).
