V O L U M E 2 4 , N O . 4, A P R I L 1 9 5 2
67 1
RESULTS
The accuracy of the methods was verified by determining the sesaniol and sesamiri constants of sesaniin concentrates and sesame oils before and after the addition of known amounts of these substances. The results of the recovery tests with added sesamol as determined by the spectrophotonietric and photocolorimetric procedures are shon.11 iri Table I, while those for added sesamin as determined spectrophotonietrically are shown in Table 11. I t is seen from these tables that the recoveries of added sesaniol and sesaniin in sesamin concentrates and sesanie oil vary from 90 to 10lyo,which may be considered satisfactory. The spectrophotometric and photocolorimetric procedures were applied to a series of commercial sesamiri concentrates and t o sesame oils, as indicated in Table 111. The sesamin concentrates ranged from 0.04 to 0.23% free sesamol, 0.90 to 2.35% bound sesamol, and 5.13 to 19.23% sesaniin, respertivcly, n-hereas the oils ranged from 0.0003 to 0.08% free
seeaniol, 0.0002 to 0.13% bound seeamol, and 0.14 t o 1.06% sesaniin. LITERATURE CITED
( I ) Budowski, P., J . Am. Oil Chemists’ Soc.. 27, 264-7 (1950). (2) Budowski, P., and Markley, K. S., Chcm. Rei).,48, 125-51 (1951). (3) Budowski, P., Meneees, F. G. T., arid Dollear, F. G., J . Am. Oil Chemists’ Soc., 27, 377-80 (1950). (1) Budowski, P., O’Connor, R. T . , and Field, E. T., I h i d . , 27,307-10 (19501. --, ~~
(5) (6) (7) (8)
Ihid., 28, 51-4 (1951). Gravenhorst, C . O., I n d . Eng. Chon., 16, 47-8 (1924).
Honig, P., Chena. Weekhlad, 22, 509-12 (1925). 3Iooi.e. R.S . ,and Bickford, K. G., J . L 4 7 ~Oil ~ . Chemists’ Soc., 29, 1-4 (1952). 1%. S., and Mattill, H. .i..(.’hem Rev.,
(!>) Olcott,
29, 257-6s (1941).
R~~CF:IVE for : D review October I?, 1951. Accepted J a n u a r y 17, 1932. Presented a t the 25th fall meeting oi the American Oil Chemists’ Society, Chicago, Ill., October 8 t o 10. 1951. S o . 9 of a series of communications on sesame oil. Instriiments are named as part of t h e exact experimental conditions. This does not constitiite a recommendation oi t h e Department of A g r i c u l t i i r e of these instrunients over those uf a n y other manufacturer.
The Dead-Stop End Point I-ticreduction at the cathode. An>-titration w-hich can be arranged so that a change from an electrolytic redox couple to onlr one process or vice yersu orcurs at the equivalence point can be adapted to a dead-stop end point.
F
(4)origiiially proposed the dead-atop method using tivo platinum electrodes and a small applied potential. The end point in a titration !vas shown by either the disappearance or appearance of a current flowing between the electrodes. The phenomenon was explained by assuming t h a t hydrogen mid oxygrn x e r e adporbed on the surface of the cathodicali7. and anodically charged electrodes and the removal of one of t h r w changed the system. Bottger and Forche ( 1 ) suggested that the 1,Fniv. applied potential used by Foulk and Baxden was not sufficient for hydrogen formation on the cathode, and showed that when the iodine-thiosulfate reaction ivas rarried out in divided cells, the electrode potential varied tyith the concentration. Delahay ( 2 )reported that the “dead-stop” end point depend8 011 the change from a reversible redos couple to an irreversihlr redox couple or vice versa at the equivalence point. These observations suggest t h a t thc dead-stop phenomenon is not a polarization effect based on gas adsorption on the electrode surface, but at least in part an electrochemical phenomenon based on oxidation a t the anode and reduction at the cathode. If this is so, then hy applying suitable potentials and using appropriate electrodes any redox titration may be followed using the deadstop method. In order t o test this theory, dead-stop titrations were carried out with various sJ-stenis, and electrode potentials and currents !yere measured. OULK and Bawden
0.3
- 0 I(0
0,’s
O(5
0.;
IlO
~
ML. I, Figure 1. 0.1
Electrode Potential and Current between Electrodes N thiosulfate titrated with 0.1 N potassium iodideiodine with 50
mv.
applied
EXPERI;\I E S T A L
Apparatus. A systein such as that described by Kernimont and Hopkinson (8) is convenient for applying the potential t o the electrodes and measuring the current which flows. A Fisher
ANALYTICAL CHEMISTRY
672
E. I. du Font de Sernours ti Co. in dilute sulfuric acid and filtering to remove insolubles. S o n e of the solutions was standardized, since only a qualitative esaniination of the end-point phenomena was desired and the potentiometric break was beingohserved. Appropriate indicators-e.g., o-phenant'hroline ferrous Kith ceric sulfate, thiocyanate with ferric, and the color of iodine and permanganate-were found to change xithin a few hundredths of 1 mi. of the end point indicated by the dead-stop and potentiometric breaks. General Procedure. Twenty-fire milliliters of the solution to be titrated were placed in a 100-rn1.1)eaker and 150 nil. of 1 to 4 sulfuric wid (water only in the case of thjosulfate) irere added. The platinum wire electrodes, previously cleaned in dichromate-sulfuric acid cleaning solution, the saturated raloniel electrode and the glass paddle of a mechanical stirrer were placed in the solution. The previously determined potential was applied to the platinum ?levtrodes, the stirrer was started, and the reagent was added from a buret. The current flowing between the elertrodes and the potential of each electrode 1's. the saturated calomel electrode viere nieasured with the potential applied and the solution'sbeingstirred.
1.2 1.0
$08
V I '
v;
E 06 P
W
04
0.2 0.0 Figure 2.
Dead-Stop End Points with Ceric Ion A. FeSOd 100 m v . applied B . KAFe(CN)s, 50 m v . applied C. Tiz(SOd)n, 100 m v . applied
Elecdropode was used in this ~ o r k The elcctrodcl potentials were measured with the polarizing potential applied using a Beckman Model G p H meter and a Reckman saturated calomel electrode as the reference electrode. Reagents. Solutions of ferrous sulfate (S),potassiuni permanganate, iodine with potassium iodide, ceric sulfate, sodium thiosulfate, ferric ammonium suliate, and sodium metavanadate were prepared in the usual way. Titanous sulfate solution was prepared by dissolving metallic titanium sponge obtained from
80
75
w
60: u a
z
.a
0
455
z n
30
15
ML. 0.1 N KMNO,
5
10
I5
20
25
ML. CE'4 Figure 3.
Complete Current Curve
KdFe(CN)e titrated with ceric ion
0 I
Figure 4.
Dead-Stop End Points with Potassium Permanganate
+
A . FeSOd H&'O4,100 m v . applied B . K&Fe(CN)s,50 m v . applied C. Tia(SOda, 50 m v . applied D . HeOz, 50 m v . applied
'fable I .
Potentials t o Be ipplied
r t 1 111
I'otrntial, Mv.
__
Foulk and H a i d e n found that a potential of 15 mv. applied itcross the electrodes gave satisfactory results in the iodine-thiosulfate titration. I n view of t h e electrolytic nature of the end point t,he potential t o be applied in a particular case would be expected t o depend to some extent on the reversibility of the couple. K h e n a sgstzm jvhich is irreversible with no applied 1)oteiitial or in which the electrolytic processes are different a t the two electrodes is under Consideration, special care is necessary to predict the effect on the current
is the basis for the In order to teat the theory that electrol! tlcatl-stop end point, a wrics of titrations T V ~ Pcarried out using sevt~ralcommon oxidizing anti reducing agent5. Observations of the potential of each elcv-tiode and the current flowing in the vicinity of the end points are shown in Figures 2 and 3 using ceric sulfate as the oxidizing agent, Figure Iusing potassium pixrmanganatc as the osi:lizing agent, and Figure 5 using miscellttneous redox systems. C'omplete current olmrvations are sh0u.n for one case iit Figure :3. .ill these s!-stems behave i n the same manner.
Consider an oxidation using potassium permanganate. The manganese system in acid medium contains A h - + + and %In04-. These do not form a reversible redox couple, but oxidation of lh-L to manganese dioxide and reduction of M n 0 4 - t o RIn04-or manganese dioxide are accomplished by electrolysis and a current flows as seen in Figure 4. Hydrogen peroxide (Figure 4, D ) is another example of an 11 1 eversible system T\ hich forms an electrolytic couple, since it is oxidized a t the anode to oxygen and reduced a t the cathode to I i j dro\yl ion, with the result that a small current flows.
(lonsider the oxidation of ferrous sulfate by ceric sulfate. I3rbi'ore the end point the reversible ferrous-ferric couple is present. The current used in the electrolysis will be limit,ed by the amount of ferrous ion present and hence will decrease toward zero as the ceric sulfate is added. The consumption of ferrous ion by electrolytic oxidation will not affect the accuracy of the titration, because the same amount of ferrous ion will be produced by reduction of the ferric ion. The reversihle cerous-ceric couple is
Therefore the dead-stop end point may be used when a n c,lertrolytic process is possible even if a redox system is not l)rrsent, as nil1 be sholvn in another paper. Changing the pot,ential changes only the semitivity (change in current with volume of reagent) and not the location of the end point. Table I shows some potentials which have been found to he satisfactory for the couples listed. For combinations of
1
ANALYTICAL CHEMISTRY
674
Delahay, P.. Anal. Chirn. A c t a , 4, 635-40 (1950). ENG.CHEM.,-4s.k~.ED.,17, 530 “1945). Duke, F. R., IXD. Foulk. C. IT., and Bawden, 9. T., J . Am. Chem. S o c . , 48,204551 (1926). Gale, R. H., and Rlosher, E., As.41.. CHEY.,22, 942-4 (1930). Latimer, W. H., “Oxidation States of the Elements,” 1). 56. Sew Tork, Prentice-Hall, 1938. I b i d . , p . 67. Wernimont, G., and Hogkinson, F. J., ISD. Esc;. CHEX, -1s.4~. ED.,15, 272-4 (1943).
these couples such as are found in redox titrations, it is best to use the higher potential if there is a difference, in order to obtain a larger break a t the end point. CONC LU S I 0 3
The applications of the dead-stop end point are mole numerous than originally proposed by Foulk and Baivdcn and depend on the ingenuity of the analytical chemist to h a l e an electrol\ tic couple a t the equivalence point of the titration. LITERATLRE CITED
( 1 ) Hottger, W., and FoIche, 1%.E., Mtkrocliiurze, 30, 138 -53 (1942).
R K C E I V Efor D review M a y 9, 1951. Accepted January 14, 1952. Presented 1x.foi.c the Division of .4nalytical Chemistry a t the 119th hIeeting of the AMERICANCHEMICAL SOCIETY, Cleveland, Ohio. Abstracted in part irorii the thesis for the degree of master of science submitted by H . G. Scholten t o the Graduate Faculty of Michigan State College.
Titration of Enols and Imides in Nonaqueous Solvents J..inlES
S. FRITZ
Iowa S t a t e College, Ames, Iowa
l h e purpose of the present investigation is to define the scope and limitations for the titration of imides and enols as acids in nonaqueous media. Conipounds containing the .A-Clb-.\’ or .A-X€€-..i’ configuration can be titrated in dimethylformamide or eth?-lenedianiineif A and -4’ are carbonyl groups o r an)- of several other electron-withdrawing groups. Sodium methoxide is used as the titrant and the indicator is thl-molblue, azo violet, or o-nitroaniline, depending on the acidic strength of the compound being titrated. The method offers a fast, accurate method for determining enols and imides, if other acidic conipounds are absent. The titration of theobromine in the presence of caffeine is particularlJ-notelc-orthy.
ltI-ellon to orange red. SCOPE
Enols. Coinppunds of the type .\-CH2-.A’ H I ‘ P suffic~ieiltly acid to permit titration, provided A and -1’are groups possessing ‘suitahle elertron-\vithdra\\.inp propei.ties. If .I anti A ’ are 0 0 0 0 I
,I
-e--It, -c-11,
,i
--C--oR,
~
--C-SIIAY,
or -(‘S,kCCUI‘8tf’ titratioit in diniethylforriianiide is pos4lile ’using azo violet indica0 tor. The amide group, -C--SH,, hae weaker electron-withtlt:tn-ing properties. This is shown by the fact t h a t while malononitrile gives a good azo violet end point, cyanoacetamide gives a very poor end point and malonamide is not a t all acid toFq-ard azo violet. Cyanoacetamide does give a sat’isfactory end point in c.~thylr~nediamineusing 0-nitroaniline indicator; malonaniide is ,slightly acid to this indicator. The -.-C=l-group in conjunction n-ith the carbonyl group in l-phenyl-3-mrbet hoxy-5-pyrazolone accounts for the fact that 0
II
this winpound can be titrated as an acid. The -C-OSa group has very slight, if any, electron-rvithdrawing propertiea. The carboxyl groups in cyanoacetic acid and in malonic acid can lie sharply titrated in dimethylformamide, but these compounds apparently have no further acidic properties. Compounds of the type A4-CH,-CH,-*4’ were found riot tmo be arid t o azo violet even if A and .A‘ are strong electron-attract-
ing groups.
Compounds of the type -C-
C-CH2-
(c,\ample
