ANALYTICAL CHEMISTRY
124 (18) Clark, W.,“Photography by Infrared,” 2nd ed., New York, John Wiley & Sons, 1946. (19) Dean, J. D., Fleming, C. 31., and O’Connor, R. T., TeztiIe Research J., 22, 609 (1952). (20) Flint, E. A., Biol. Reas. Cambridge P h i l . Soc., 25, 414 (1950). (21) Foster, 1,. V., and Thiel, E. hI., J . Opt. Soc. Amer., 38, 689 (1948). (22) Guthrie, J. D., Hoffpauir, C. L., Steiner, E. T., and Stansbury, 21. F., U. S. Dept. ilgr., Bur. Agr. and Ind. Chem., Pub. AIC6l (revised 1949). (23) Haas, H., Battenburg. E.. and Tenes, D., T a p p i , 35, 116 (1952). (24) Hermans, P. H., “Contribution to the Physics of Cellulose Fibers,” Amsterdam, Elsevier Publishing Co., Inc., 1946. (25) Hock, C . W., Textile Research J., 20, 141 (1950). (36) Hock, C. W., Ramsay, R. C.. and Harris, M., J . Research Satl. Bur. Standards, 26,93 (1941). (27) Kerr, T., Protoplasma, 27, 229 (1937). (28) Kerr, T., Textile Rosearch J . , 16,249 (1946). (29) Land, E. H., Blout, E. R., Grey, D. S.. Flower, AI. S., Husek, H., ,Jones, R. C., Matt, c‘. H., and Rlerrill, D. P., Science, 109,37 (1949). (30) Lucas, F. F., Proc. Nall. A c a d . Sci. li. S., 16, 599 (1930). (31) Mangenot, G., and Raison, AI., Bofnn. Rea., 17, 555 (1951). (32) Ilaresh, C. J.. Biol. Photog. Assoc., 21, 14 (1953). (33) llerton, T., Proc. Rou. Soc. ( L o n d o n ) , A189, 309 (1947). (34) Mitchell, R. L., I n d . E n g . Chem., 38, 843 (1946). (85) .\loore, A. T., S t a i n Technol., 28, 149 (1953). (36) ;\IuNethaler, K., Biochim. et Biophys. Acta, 3, 15 (1940).
(37) Patel, G. M., doctoral thesis 2013, Swiss Fed. Inst. Tech., Zurich, 1951. (38) Preston, J. M., “Fiber Science,” Manchester, England. Textile Institute, 1949. (39) Preston, J. M., “Modern Textile Microscopy,” London, Emmott & Co., 1951. (40) Ritter, G. J., Paper I d . , 33, 926 (1951). (41) Ritter, G. J., Paper Trade J . , 101, No. 18, 92 (1935). (42) Rollins, hl. L., Botan. Gaz., 108,495 (1947). (43) Rollins, M. L., Tertile Research J., 15, 65 (1945). (44) Royer, G. L., and Maresh, C., Calco Chem. Div., American Cyanamid Co., Bound Brook, K. J., Calco Tech. Bztli. 796 (1947). (45) Sisson, W. A., “Cellulose and Cellulose Derivatives,” E , O t t , ed., New York, Interscience Publishers, 1943. (46) Siu, R. H.. “Microbial Decomposition of Cellulose,” Sc\v I-ork, Reinhold Publishing Corp., 1951. (47) Tripp, V. W., and Rollins, M. L., ANAL.CHEM..24, 1721 (1952). (48) Uber, F. AT., ed., “Biological Research Methods,” ?;en l-ork. Interscience Publishers, 1950. (49) Van Iterson, G., Cheni. Weekblad, 24, 166 (1927). (50) Kakeham, H. R. R., and Spicer, N., Textile Research J . 21, 187 (1951). RECEIYEDfor review September 2 8 , 1953. Accepted February 3, 1!154. Presented in part before the joint symposium of the Division of Cellulose snd Polymer Chemistry at the 123rd Meeting of the AMERICAN CHEXICAL ~ C I E T Y Los , Angeles, Calif., 1953.
Polarography of Thiourea R.
L.
EDSBERG’
Central Research Laboratory, General Aniline
&
Film Corp., Faston, Pa.
T
HIOUREA is known to form complexes with a number of metals and salts (1). Although it is not reducible at the dropping iiiercury electrode, it was thought that a t an anodic mercury electrode aome reaction might take place with the mercurous ions. Experimentally it was found that an insoluble mercurous complex was formed. AI’PARATUS
.in automatic and recording polarograph described earlier (2) \vas used in this work. The dropping mercury electrode when disconnected from the polarograph had a drop rate of T = 6.90 seconds in water. Under these conditions the mass of mercury delivered was tn = 0.948 mg. per second. All polarograms were obtained using the cell described earlier ( 2 ) and at a temperature of 29.5’ C. .4 saturated calomel half cell with a saturated potassium rhloride bridge (a) was used as a reference electrode in these experiments.
for the residual current found on the polarogram of the bufl’er alone. DISCUSSION
I n buffer solution of p H 4 a definite polarographic w a w is e\hibited (Figure 1). IVhen is plotted against conccntration, a linear relationship results (Figure 2). A procedure for the quailtitative determination of thiourea is now being used in which rz measure of the diffusion current represents thiourea concentration The reaction of thiourea a t the anodic dropping mercury elttctrode could be an oxidation or a reaction with mercurous or mercuric ions to form insoluble or very stable complex products. In an attempt to identify this reaction the following polarographic evidence was obtained. Polarograms of the same thiourea concentration run a t plI 4 and 6 have the same half-wive potentials. This eliminates a n y reaction that might involve hydrogen ions.
REAGENTS
Thiourea was recrystallized a t room temperature from a solution in ethyl alcohol by adding petroleum ether. Analysis of this product gave C = 15.87% and H = 5.34%; theory is C = 15.78y0and H = 5.30%. Buffer solutions, p H 4 and 6, were Clark and Lub buffer mixtures containing sodium hydroxide and potassium biphthalate. PROCEDURE
A cell containing 25 ml. of buffer solution was deoxygenated by x stream of nitrogen and then polarographed from -0.2 to +0.3 volt. An aliquot of a standard aqueous thiourea solution was then added, and the polarogram was repeated. T h e half-wave potential becomes more negative with increasing concentration, so measurements of limiting current cannot be made between set potentials. Each limiting current measured was corrected
*.2’
+ I‘
u)
20.
>
I
-.2 0
0 . O O U V Y THIOUREA IW BUFFER pH 4
1
2
3
4
5
6
7
V O L U M E 26, NO. 4, A P R I L 1 9 5 4
725
The half-wave potentials of polarograms run on various concentrations of thiourea were found to shift to more negative values a i t h increasing concentration. In order to measure these potentials as accurately a8 possible, the I R drop across the cell must be considered. The resistance of the system, from the mercury reservoir through the cell of buffer pH 4, potassium chloritlr bridge, and saturated calomel electrode, was measured as 2300 ohms with a General Radio Co. impedance bridge Type 650-.I. Psing the espression E = ZR and the id values measured, the voltage drop a t each concentration was calculated. These corrections were subtracted from the measured half-wave potentials t o get the true half-wave potentials.
Assuming an insoluble reaction product
(5)
Substituting this concentration i n Equation 3 describes the polarographic wave in terms of thiourea concentrations.
.4t any point Ede on the polarographic wave in the concentration 2). of thiourea at the electrode would be proportional to (id
-
5
I
At the half-wave potential
Ei/z
=
jd
E"'
= 22
- 0.0:3 ah 10::
(8)
(id)
If Equation 8 is a true espression of the reaction, a plot of log ( i d ) versus must produce a straight line with a slope of 0.03 -_ a. Figure 3 represents the correct half-wave potentials plotted against the log i d .
2
4 6
6
IO 12 14 I6 I6 20 THIOUREA, x I O - 4 y
Figure 2.
22
i4 26
-4slope of 0.0516 was found, indicating
a ratio of a_ = 1.72. The experimental value of n = 1.72 b cherks fairly well with the theoretical value of 2.
&
Linearity Check
The shifting EllI and pH independence suggests a complex or insoluble reaction product with mercury rather than an oxidation If the zwitterion form of thiourea is accepted (4),the reaction may he represented as follons
4! 2
\
Since the mercurous ions are furnished by an anodic oxidation from the dropping mercury electrode, the electrode reaction will he
2Hg
-+
Hg?"
+ 2e
(2)
T h e equation for the polarogritpliic wave will then be Ede
=
E'red
+ 0.06 -yIOg [HgrVT]a t 30'
C.
(31
if activities and activity coefficients are neglected. The free mercurous ions a t the electrode will be present in a concentration governed by the solubility product or equilibrium c.onstarit of the product from Reaction 1.
L
*.02
4.03
Figure 3.
*.04
4.05 1.06 EK in VOLTS
+.07
t.08
t.09
.E'/* Change with Concentration
A plot of log (id - z ) & versus E,I,:IS in Equation 7 should prodyce a straight line if the reaction is reversible. Diffusion currents a t various potentials on the wave were carefully measured from the polarogram illustrated in Figure 1 . Each potential was corrected for the ZR drop as described earlier. These values when plotted as in Figure 4 show that no straight line results, indicating an irreversible reaction.
726
ANALYTICAL CHEMISTRY
!I
SH,
SH*
\
2 C-S-
6
//
+ Hg,++
+NH*
+
>'-S-Hg-Hg-S-(:
9
juH? \\
+SH,
ISH*
Sartori and Liberti ( 3 ) have also described the reaction of thiourea at the dropping inercury electrode as being a complex or insoluble compound formation rather than an oxidation. They did not, however, give any details of t'he reaction or enough esperiment,al evideiice to allow a derFvation of possible reactions. I
0
*.02
4.04 €4..
Figure 4.
t.06
+.08
ACKYOWLEDGMEIVT
+.IO
in VOLTS
Reversibility Check
The only other possibility that seems to fit the experimental data is that a soluble conipIe\: is formed having the composition
The author gratefully acl;rio\\-ledgcs t'he help and advice of Sidney Siggia, H. J. Ptolten, and I,. J. Frauenfelder. Thanks are also expressed t o I,. T. Hallett, under whose direction this work \\-as performed. LITERATURE CITED
(1) Degering, E. F., "Organic S i t r o g e n Compounds," PD. 1367-87, Ypsilanti, Llich., L-niversity Lithoprinters, 1943. (2) Edsberg, R. L., Eichlin, Dale, and Garis, J. J., ANAL.CHEM.,25,
Judgirig from known complexes and reactions of thiourea, the first assumption is more reasonable, and the reaction is probably as indicated herewith.
798 (1953). (3) Sartori. G., and Liberti, .I.,Proc. Intern. Polarog. Congr. Prague, I d COILQT.,1951, Pt. I, PP, 260-4. (4) Sidgwick, X. V.,"The Organic ('hemistry of S i t r o g e n , " pp. 27592, Oxford, England, Oxford Press, 1942. K E C L . I Y Efor D review October 8, 1953.
Accepted Sovember 2 1 , 1953.
Comparison of pK; Values Determined by Electrometric Titration and Ultraviolet Absorption Methods 1. M. VANDENBELT, CAROLA HENRICH, and SUZANNE G. VANDEN BERG The Research Laboratories, Parke, Davis & Co., Detroit 32, M i c h .
T
HE dissociation constants of organic compounds are receiving
increased attention as the possibilities of practical application t o problems of structure and identity and to the determination of absorption spectra, are more generally appreciated. In view of this importance, it is of interest to compare pKA values obtained by two methods of determination, electrometric titration and spectrum absorption, as a partial evaluation of the relative quality of data and the merits of the tlvo methods in routine laboratory operations. DISSOCIATION C O h S T A h T \ ALUES
Methods. I n connection with extensive observations of the absorption spectra of purified benzenoid compounds (6, 7 ) , the pK: values of those derivatives with ionizable groups were determined by the absorption technique of Sager, Schooley, Carr, and Acree ( 1 4 ) . T h e potentiometric values of these compounds were obtained subsequently, using the method of Parke and Davis (IS) and equipment identical x i t h theirs. Results. Table I summarizes p K ' values obtained in routine application of the two methods to substituted phenols, anilines, and benzoic acids, at a room temperature of 25" C. The absorption figures given are the means of values from all the intermediate curves obtained for each ionization; the titration figures are t o the nearest 0.05 unit read off the plotted curve through the transparent plastic overlay ( I S ) . Agreement with Literature Values. It is of course importafit that values obtained by both absorption and titration methods, in addition to agreeing with each other, be of acceptable accuracy. A comparison with representative constants available in the literature shows that either method gives satisfactory values.
The pK; of phenol, for example. iq reported as 9.94 (8) and 9.95 ( 9 ) ; aniline as 4.60 ( 1 1 ) and 4.636 ( 1 0 ) ; benzoic acid as 4.18 ( 2 ) ,4.21 (4),4.30 ( f ) , and 4.198 ( I O ) . Values for disubstituted compounds are equally succesPfu1: p-Hydrox!-benzoic acid, first anion, is given as 4.5:3 ( 4 ) and 4.54 (2); second anion, as 9.53 ( 1 4 ) and 9.39 (at 20" C . ) ( 1 2 ) ; p-t'oluidine, as 5.12 ( I I ) , 5 . l i (3),and 5.098 ( 1 0 ) : tu-toluidine as 4.77 (10); o-toluidine, as 4.46 ( 1 0 ) ; p-anisic acid, as 4.43 (41, 4.44 (2) and 4.47 ( , 5 ) . A number of the pK', values of the compounds in Table I have not been reported heretofore; some are given in mixed rolvents or a t tempei'atures tiiffe!,ent from that used in this study. In gerieral, the values by ab-orption and titration are in good :tgreement with literat,ure value? obtained under comparable contiitionr. DISCUSSION
I t is evident from inspection of the table that values obtained by the two methods are in good agreement. CompariPon would indicate that either is satisfactoyy, and that, in r e d t s obtained, there is little advantage i n one over the other. However, titration is usually a bit more rapid than the absorption method if the equipment is conveniently set up and manipulative expei,ience has been acquired. This arises from the automatic nature of the procedure; it is not necessary to knoiv or discover from conr;ecutive curves the solution p H of the limiting species, nor to learn by trial the intermediate pH a t which ionization takes plarc. The titration method has a great advantage when there are two or more ionizable groups in a molecule. Usually time is prohibitive, and execution is fairly difficult', to run a sufficient number of curves t o eqtab1i.h adequately the limiting species when
