VOLUME
462
19, NO. 7
(6) Gortner, R.A., andMcNair' J. J.,Ind. Eng. Chem., 25, 505 (1933). (7) Hagglund, E., and Bratt, Southern Overcup L. C., Papier-Fabr., 34, Red O a k , Oak, Quercua Quercus Douglas Western Loblolly Black Western 100-3 (1936). Zyrata Species Fir Hemlock Pine Spruce Red Cedar rubra (8) Jayme, G., CelZuZosechem., Series A 20,NO.2,43-9 (1942). (9) Kurth, E. F., and Ritter, 0 41 0 28 0.36 0 27 0 16 0 29 0 49 Ash. ?& 2.81 1.10 1.12 0.53 3 30 1 21 0 61 Acetyl; % G. J., J . Am. Chem. Sac., 25.2 29.50 32.6 27.6 30.40 28 40 28.0 Lignin, yo 56, 2720-3 (1934). 14.65 24.8 23 3 15.5 14 IO 17.4 15.3 Hemicellulose, 7' (10) Mitchell,R.L.,andRitter,G. 65 0. 211.6 45.7 42.4, 52.7 a-Cellulose, % 57.20 51.6_ __. J.,Ibid.,62,1958-9 (1940). 98.02 1 m 99 20 101.18 97.66 Total 100 60 (11) Norman, -4.G., "CarboSeries B hydrates Normally As0.28 0 36 0 27 0 16 0.41 0.29 0.49 Ash % sociated with Cellulose in Xylbn (corrected for uronic Nature," in Ott's "Cellu10.50 8.1 20 0 18.2 7.30 10.10 6.20 anhydride), % lose and Cellulose Deriva4.52 4.1 4.2 5.8 5.0 3.8 2.80 Uronic anhydride, % 8.0 5.1 4.7 4.1 5.4 Mannan, % tives," p. 431,New York, 3 . 3 1 . 1 2 0 . 5 3 1 . 2 1 1 . 1 2 . 8 1 a 0 61 Acetyl, % Interscience Publishing a-Cellulose (corrected for Co., 1943. mannan, xylan, and uronic 44.5 46.6 45.6 47.5 43.7 40.6 48 3 anhydride), To (12) Ritter, G. J., and Kurth, 25.2 30.4 29.50 32.5 27.6 28.0 28 40 Lignin, % E. F., Ind. Eng. Chem., CHp (calculated from M e 0 25, 1250-3 (1933). 0.6 0.20 0.22 0.22 0.5 0 05 0.16 not i n lignin), % __ __ (13) Schmidt, Erich, Meinel, J.. 95,92 -9 97.48 96 28 98,42 92 05 93.16 Total Jandebeur, W., and Simson, W., Cellulosechem., 13, S o . 8/9, 12939 (1932). ACKNOWLEDGMENT (14) Schmidt, Erich, Tang, Y.C . , and Jandebeur, W., Ibid., 12,No. 7,201-12(1934). Thanks are due especially t o the analytical group of The Insti(15) Thomas, B. B.. Paper I n d . Paper W o r l d , 26, No. 10, 1281-4 tute of Paper Chemistry for certain of the analyses reported in (Jan. 1945); 27, No. 3, 374-8,382 (June 1945). (16) Van Beckum, W. G., and Ritter, G. J., Paper Trade J . , 104,No. Table 111, and t o F. E. Brauns who made the lignin determina19, 49-50 ( M a y 13, 1937): 105, No. 18, 127-30 (Oct. 28, tions given in Table 111. 1937); 108,No. 7,27-9 (Feb. 16, 1939): 109,Yo. 22, 107-9 (Nov. 30, 1939); Tech. Assoc. Papers, 21,431-4 (1938): 22, 619-22 (1939); 23,652-4(1940). LITERATURE CITED (17) Wise, L. E.. Murphy. M.,and D'Addieco. A. d.,Paper Trade ( 1 ) Anderson, E., J . B i d . Chem., 165,233 (1946). J., 122. No.2,35-43 (Jan. 10,1946). (2) Assoc. Official Agr. Chem., "Official and Tentative Methods of PRESENTED before the Division of Sugar Chemistry and Technology, SymAnalysis," 6th ed., p. 412,1945. posium on Current Progress in Carbohydrate Chemistry, under the title (3) Burkart, B., Baur, L.,and Link, K. P., J . Bioi. Chem., 104, " S e w Approaches t o the Analytical Study of the Hemicelluloses of Fibrous 171-81 (1934). Materials" a t the 110th l f e e t i n g of t h e AMERICANCHEMICAL SOCIETY, (4) Cundy, P . F., private communication, 1946. Chicago. Ill. (5) Freudenberg, K., and Harder, M., Ana., 433,230 (1923).
Table 111. Analysis of Composite Wood Samples (Extractive-free,oven-dry basis)
~~
Some Formal Oxidation-Reduction Potentials HOBART H. WILLARD AND GLORIA D. RIANALO' D e p a r t m e n t of C h e m i s t r y , Unicersity of M i c h i g a n , A n n A r b o r , 'Mich. T h e formal oxidation-reduction potentials of a number of oxidizing agents have been determined in alkaline solution and in some cases also i n acid solution.
T
HE electrochemical behavior of an oxidation-reduction system is expressed by an equation which includes a constant characteristic of each system. This constant, called molal potential, is the value in volts when all the reactants are present in unit activity-that is, when gases are a t 1 atmosphere and the ions or soluble compounds a t 1 molal activity. Swift (6) pointed out the impracticability of molal potential values and asserted t h a t they may be misleading in predicting the behavior of oxidizing and reducing agents in relatively concentrated salt and acid solutions which are frequently encountered in analytical experiments. Smith (4) in his study of the effect of the hydrogen-ion concentration on some oxidation-reduction systems included some measurements of formal potentials, which he defined as the potential values for certain reactions when the concentrations of the substances involved are 1 formal. A formal solution contains 1 gram formula weight of the stated compound per liter and gives only the total amount of the substance present, not the specific molal or ionic species in solution. Formal po1
Present address, Pateros, Rizal. Philippine Islands.
tentials ignore correction due to activity relationship and do not include hydrolysis effects for which molal potentials are corrected. Formal potentials are measured under the actual conditions of the experiment and in the presence c,f all substances involved in a reaction. Formal versus molal oxidation potentials differ over such a wide range under the normal variations in titrational environments that special Consideration of this situation is of prime importance. Smith ( 4 ) found t h a t t h e formal potential for the reaction
Cr207--
+ 14H" + 6e c*
2 Cr++'
+ 7H20
is below its molal potential by 20wc and the system Fe'++
+ e 73 F e + +
by 10%. These values were determined by potentiometric titration of the ferrous salt with potassium dichromate. The formal potentials of various systems have been determined in acid solution. Smith and Richter (5, 6 ) have measured the oxidation potential for 1,lO-phenanthroline ferrous su1fat.e
TULY
463
1947 ‘Fable I.
Miilal Potential ( 3 ) 1701ts
F
...
-CO 36
. , .
F
oi 1 3
1Iethod 1
1
2
Method 2 (with arsenite)
...
N e t h o d 2 (with chromic salt)
...
hfethod2 ( u i t h antimonite) CrrO; --5 Cr c +
- 1 . 3 6‘(acidic) - 0 . 1 2 (alkaline)
1Iethod 1 3Iethod 2 (with ferricyanide) BrO- eRrMethod 2 (with arsenite)
t O 76 (alkaline)
l l e t h o d 2 ( a i t h cyanide)
...
l r e t h o d 2 ( w i t h thiosulfate)
, . .
-LO
5.59 (acidic)
...
3Iethod 1
-0
3 5
1 3 5 1
a
5 3
...
. . ...
... . .
‘i’ 3 5 1 3 5 1 3 1 3 1 3 5
... ...
. . . .
. . . .
71 (alkaline) 1 3 5
...
-1.00
Nethod 2
accurate than *0.01 volt. The measurement of formal oxidaFormal Potential tion-reduction potentials was Other Reports made in solutions 1 F to 5 F i n 31edi u in O b s e n ed I’olts sodium hydroxide. The derivatives of, diphenylamine as +o 57 f O 48 0 01 F HCI indicators are destroyed in + O 56 0 1 F H C I + O 69 1 0 71 1 F H C I t o 77 strong alkaline solutions. No T O 72 1 F €Id301 f O 72 1 F HCIOI measurement was, therefore, , .Su7ft , . . ... I/? , taken a t concentrations higher + O 46 0 01 b KaOH +o 47 Swift (6) + O 49 than 5 F sodium hydroxide. + O 50 + O 52 Results are shown in Table I. +o 47 The values for the two methods + O 50 + O 52 and for different reagents in + O 46 + O 50 method 2 checked well xyith +o 5 2 each other. + O 50 +‘I 0 2 1F Hbl’ The results of the ahove ex+ 1 , l o 4 F HCI +1 10 2 F H ~ S O I perimeit conform xvith the find+ I 14 4 F I r m a ings of Smith (9) that the +1 34 8 F H2SOr Smith (5) formal potentids of oxidation-0 12 reduction systems in solutions -0 09 (concentrated salt nnd acid) -0 08 -0 12 ,.. encountered in analytical chen-0 09 -0 08 istry arc much diferent from molal potentials. 4 0 69 + O 66 In alkaline medium, the + O 70 formal potential of the ferri+ O 67 ... + O 70 cyan id e-f er r oc y a n id e , +o 67 -LO 65 chromate-chromic, and hypoiO.58 1FHCl h r o m i t e-ti r o m i d e s p t e m s + 0 . 5 8 1 F HClOa changes only 0.05 volt \\-hen Swift (6) the concentration of sodium - 0 08 -0 09 hydroxide is increased from 1 -0 16 F to 5 F . The formal potential 1’1’0 2 1 F HCl’ of the arsenate-arsenite couple -C1.02 1 F H C I O I S wif t ( 6 ) changes by about 0.08 volt, but +1 02 1F H2S0, the deviation of the formal ++1 1 . 0 7 2 F TiQSOa 14 4 F HBO4 potential from its molal poten+ l , 80 8 P H ~ S O L Smith and Richter f 5 ) tial is very high, ahoiit 0.63 volt. -0 74 ... ... The formal potential for the -0 8 5 -0 86 vanadate-vanadyl system becomes more negative as the concentration of the sodium hydroxide is increased. The change is 0.22 volt when the concentration of allcali is increased 1 F to 5 F . Vanadyl sulfate is thus an extremely strong reducing agent in alkaline solution. It is rapidly oxidized when exposed to air. This correlation of results shows further how important is the measurement of formal oxidation-reduction potentials in predicting the behavior of such systems. Attempts were made to measure the oxidation-reduction potential of diphenylamine sulfonic acid in alkaline solution. No potentiometric curve could be drawn, as the “jumps” mere inconsistent. Similarly, other derivatives of diphenylamine are not stable enough to allow a potentiometric titration in alkaline solution.
Formal Potentials of S o m e Oxidation-Reduction Systems
, .
3Igthod 1
... 1 3 5
(ferroin) and substituted phenanthroline complexes, vanadylvanadate system, 2,2’-hipyridyl, and methyl ferroin in various acid concentrations. The oxidation-reduction potentials for different systems involving the various states of oxidation of vanadium in acid solution have also been thoroughly studied ( I , 2). T h e formal potentials of systems in alkaline solutions vere measured in connection a i t h other work ( 7 ) on the use of indicators in such solutions and it seems desirable to record the results, even though they are not highly accurate, because of the lack of such data in the literature. EXPERIMENTAL
For the determination of formal oxidation-reduction potentials, systems were made either by mixing equivalent amounts of the oxidized and reduced forms in the desired acid or alkaline solutions (method l ) , or by partial oxidation or reduction (507,, determined by preliminary titration) as required by the system (method 2). The potential was measured by a potentiometer set up against a carefully checked saturated calomel electrode. The second method (par’tial oxidation or reduction) of making mixtures assumes t h a t the reaction follows the theoretical equation. Some mixtures viere heated as required but measurements were always taken a t room temperature, about 22 ’. The accuracy of the measurement was not sufficient to require careful temperature control or to take account of liquid junction potentials. The potentials measured were constant, but probably not more
LITERATURE CITED
(1) Jones, G., and Calvin, J. H., J . Am. Chem. Soc., 66, 1563, 1573 (1944). (2) Jones, G., and Ray, W. d.,I b i d . , 66, 1571 (1944).
(3) Latimer, W.M.,“Oxidation States of the Elements and Their Potentials in Aqueous Solutions,” pp. 294-301, New York,
Prentice Hall, 1938. The signs are reversed. (4) Smith, G. F., Trans. Illinois State A c a d . Sci., 36, 132 (1943). ( 5 ) Smith, G. F., and Richter, F. P., ISD.ENB.CHEM.,ANAL.ED., 16,580 (1944).
(6) Swift, E. H., “A System of Chemical Analysis,” pp. 49, 540-3,
New York, Prentice Hall, 1939, (7) Willard, H . H., and Manalo. G . D., ANAL.CHEM.,19,167 (1947). FROM a thesis presented by Gloria D. Manalo t o the Graduate School of t h e University of Michigan in partial fulfillment of t h e requirements for t h e degree of doctor of philosophy.
