planer spacings and intensities with that of the original A-pattern. As heat treatment a t high temperatures undoubtedly causes molecular rearrangement @), the results observed here demonstrate the limitations occasionally encountered when diffused and often difficultly reproducible x-ray diffraction patterns are used to interpret polymer thermograms. ACKNOWLEDGMENT
The author wishes to thank Fred Smith, W. A. P. Black, and B. J. D. Meeuse for their generous material assistance. LITERATURE CITED
\I/ , 1\/' 100
,
,
.
200
,
,
,
.
40@
300
TEMPERATURE, Figure 3. A. B. C.
D.
(1) Abdel-Akher, M., Hamilton, J. K., Montgomery, R., Smith, F., J. Am. Chem. SOC.74,4970 (1952). (2) ~, Abdel-Akher, M.. Smith F., Zbid., 73, 994 (1951). ' (3) Brimhall, B., Znd. Eng. Chem. 36, 72 (1944). (4)
.
, 500
6 3
"C.
Effect of moisture on rice starch thermograms Rice starch vacuum dried Rice starch humidified Rice starch preheated to 130°C. Rice starch preheated to 130'C. and humidifled
relative humidity. Both curves have endothermic peaks a t 130", 275", and 310" C. Humidification alters only the magnitude of the 130" C. endotherm. A sample preheated to 130" C. and then analyzed shows thermogram C with a peak a t 165" C. This peak is shifted t o 115" C. (D)if the sample is first preheated to 130" C. and then humidified. These results suggest that the original 130' C. endotherm is not entirely due to the loss of residual moisture and that the dehydration-
hydration process is not completely reversible. Analogous results have been recorded with other polysaccharides. It is clear from the studies outlined here t h a t differential thermal analysis provides a new means for studying the dehydration-hydration process. I n comparison with other polysaccharides studied (9, IO), the diffused x-ray powder diffraction patterns of the humidified preheated rice starches are nearly identical both in their inter-
(5) (6)
,- - _-, . ( 7 ) Manners, D. J., Ann. Repts. Progr. Chem. (Chem. SOC. London) 50, (1953). (8) Meeuse, B. J. D., Kreger, D. R., Biochim. Biophys. Acta 13, 593 (1954). (9) Morita, H., ANAL. CHEM. 28, 64 (1956). (10) Morita, H., J. Am. Chem. SOC.78, ' 1397 (1956). (11) Morita, H., Rice, H. M., ANAL CHEW27,336 (1955). (12) Northcote, D. H., Biochim. Biophys. Acta 1 1 , 471 (1953).
RECEIVEDfor review May 18, 1956. Accepted January 18, 1957. Presented in part a t the Symposium on Thermogravimetry and Differential Thermal Analysis, 129th Meeting, ,4CS, Dallas, Tex. -4pril 1956. Contribution 320, Chemistry Division Science Service.
Rapid Quantitative Determination of Hydroquinone CHARLES B. JORDAN Coating & Chemical Laboratory, Aberdeen Proving Ground, Md.
F A rapid and accurate method which utilizes only standard laboratory equipment and available inexpensive chemicals, has been developed for quantitatively determining hydroquinone. It is a simple titration procedure and is used for rapid analyses when large numbers of samples are being examined.
H
is used extensively in many fields of chemistry, as a photographic developer, dye intermediate, medicine, antioxidant, and in many other applications. Xumerous procedures for its determination have been developed, several in recent years (1-7). Most procedures involve ordinary or electrometric titrations using YDROQUINONE
titrating reagents such as iodine, bromine, alkaline periodates or dichromates, compounds of cerium, iridium, vanadium, and various organic compounds. No quantitative procedure could be found in the literature which used ferric chloride as a titrating reagent for the determination of hydroquinone, even though ferric chloVOL. 29, NO. 7, JULY 1957
1097
'
ride is the most widely used qualitative analytical reagent for hydroquinone detection. As a ferric chloride titration is simple and exhibits good accuracy, the subsequent procedure is suggested.
Table I. Determination of Hydroquinone in Aqueous Samples; Typical Results
PROCEDURE
Blank Titration. Accurately weigh about 1 gram of chemically pure hydroquinone into a 500-ml. volumetric flask and fill t h e flask with distilled water t o the graduation mark. After the hydroquinone has dissolved, pipet a 50-ml. aliquot into a 125-ml. Erlenmeyer flask and titrate with a 4% aqueous solution of ferric chloride hexahydrate, a t a rate of 10 t o 20 drops per minute, using a yellow background for t h e titration. Agitation of t h e flask after the addition of each drop is essential. The addition of ferric chloride causes t h e formation of a green color which immediately disappears. The end point is attained when t h e green color is no longer formed. Sample Titration. The p H of t h e aqueous sample containing t h e unknown quantitg of hydroquinone is adjusted t o between 3.8 and 5.0 by dilute hydrochloric acid or sodium hydroxide, then titrated with t h e ferric chloride solution according t o the titration procedure above. Samples larger t h a n 50 ml. should be aliquoted before titration. Calculation.
Grams of hydroquinone in sample ml. of FeCh used in sample titration ml. of FeC13 used in blank titration
-
=
wt. of
Known Hydroquinone, Gram 0.020 0.040
Hgdroquinone Received, Gram
0.080
0.oSi +o.ooi
0.100 0.150 0.200 0.250 0.300
0.102 0.150 0.205 0.255
0.400
0.500 0.600 0.700 0.800 0.900 1.000
0.020 0.040
0.310
0.405
0.510 0.595 0.710 0.802 0.905 0,990
Error,
70
HydrcGram quinone
Error,
0.000
0 no0
0
n
i.2 20 0.000 0 +O 005 2 5 $0.005 2.0 +0.010 3.3 $0.002
i0.00i i . 3 $0.010 -0.005 $0.010 t0.002 +0.005 -0.010
2.0 0.8 1.4 0.3 06 1.0
DISCUSSION
The chemical mechanism of the procedure involves first the formation of the green ferric chloride-hydroquinone complex, which subsequently oxidizes to the quinhydrone complex or quinone. Molecular ratio determinations show t h a t under the conditions of the test as written, 3 moles of hydroquinone react with exactly 2 moles of ferric chloride. This ratio is contrary to published ratios in other procedures in which the hydroquinone is completely oxidized to quinone, but can be readily explained by assuming the formation in the final mixture of 1 mole of quin-
hydrone while 2 moles of quinone are being formed. The equilibrium of the reaction makes it necessary to control the temperature and pH. Temperatures from 35" to 95" F. gave satisfactory results. Temperatures above 95" were not satisfactory. The only p H values which proved satisfactory were those between 3.8 and 5.0. A yellow background is necessary for the titration, because of the yellow color of the ferric chloride. The yellow background filters out the yellow mlor and makes the green color more visible and the end point more exact. Kormal results show a maximum error of 3.3 yo based on per cent of hydroquinone in samples containing from 0.02 to 1.0 gram of hydroquinone (Table I). Compounds which are oxidized by ferric chloride and compounds containing colors which mask the end point will interfere with the procedure.
LITERATURE CITED
Belcher, R., Stephan, W. J., Ana!yst 76. 45-9 11951'1. Bogdanov, S: G,'Sukhobokova, N. S., Zhur. Anal. Khim. 6 , 3467 (1951). Furman, N. H., Adams, R. N., ANAL. CHEX 25, 1564-5 (1953). Rao, G., Rao, V., Sostri, M., Current Sci. (India) 18. 381-2 11949). Singh, B . , Singh,'A., J . jndian Chem. SOC.30, 143-6 (1953). Takahashi, T., Kimoto, X., Kimoto, RI., J . Chem. SOC.Japan 55, 283-5 (1952). Tomicek, O., Valcha, J., Chem. Listy 44, 283-91 (1950). '
C.P. hydroquinone
10
RECEIVED for review October 12, 1956. -4ccepted January 14, 1957.
Reduction of Aqueous Iodine by Trace Impurities JOHN H. WOLFENDEN Department of Chemisfry, Darfmoufh College, Hanover, N. H.
The slow development of triiodide ions in aqueous solutions of iodine is due mainly to the presence of almost inevitable traces of dust and not to the slow hydrolysis to iodic acid. Unless this effect is taken into account, the use of ultraviolet absorbance to measure very small concentrations of triiodide ion may b e in error. Some radiochemical and electrometric measurements with very dilute solutions of iodine may also b e affected.
T
absorption peak of the triiodide ion in aqueous solution at 3530 A. is so high [molar extinction coefficient = HE
1098
ANALYTICAL CHEMISTRY
26,400 ( I $ ) ] t h a t it provides a very sensitive and convenient method of detecting and measuring very small concentrations of iodide ion in the presence of iodine or of iodine in the presence of iodide ion. It has found a nuniber of applications (1, 3, 6, 8 ) . This paper points out a source of error attributable to the sensitiveness of the method. A saturated solution of iodine in distilled water a t 25" C. shows a n initial triiodide concentration, computed from the absorbance a t 3530 A., normally somewhere between 5 and 8 micromolar. The lower of these values is close to that attributable to the (virtually instan-
taneous) hydrolytic equilibrium form hgpoiodous acid: 12
to
+ HzO = H f + I - + HOI
The upper limit is well below the value of 15 micromolar which can be calculated, using the free energy data of Latimer ( 9 ) ,as the ultimate equilibrium concentration of triiodide ions resulting from the simultaneous collaboration of the hydrolytic equilibrium to produce hypoiodous acid, that to produce iodic acid:
