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
Known Hydroquinone, Gram 0.020 0.040
PROCEDURE
Blank Titration. Accurately weigh about 1 gram of chemically pure hydroquinone into a 500-ml. volumetric flask and fill the 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 the titration. Agitation of the flask after the addition of each drop is essential. The addition of ferric chloride causes the formation of a green color which immediately disappears. The end point is attained when the green color is no longer formed. Sample Titration. The p H of the aqueous sample containing the unknown quantitg of hydroquinone is adjusted t o between 3.8 and 5.0 by dilute hydrochloric acid or sodium hydroxide, then titrated with the ferric chloride solution according t o the titration procedure above. Samples larger than 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
HgdroError, quinone 70 Received, Error, HydrcGram Gram quinone 0.020 0.040
0.000
0 no0
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
0.310
0.405
0.510 0.595 0.710 0.802 0.905 0,990
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 that 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
RECEIVED for review October 12, 1956. -4ccepted January 14, 1957.
10
Reduction of Aqueous Iodine by Trace Impurities JOHN H. WOLFENDEN Department of Chemisfry, Darfmoufh College, Hanover,
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
N. H.
26,400 ( I $ ) ] that 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:
312
4- 3Ha0
=
6H+
+ 51- + 1 0 3 -
arid the iodide-triiodide equilibrium. Such a saturated solution of iodine continues to develop triiodide ion and it is natural to infer that this is due to the continued slow hydrolysis to form iodic acid. Evidence is cited to show that this is not the only nor indeed the predominant reaction. Allen and Keefer (1) reported a slon rise of triiodide concentration, particularly in solutions with a pH over 4.8 and corrected for it in their work by extrapolation back to zero time; they tentatively ascribed the effect to reduction of the iodine by traces of reducing agent in their buffer solutions, together with hydrolysis t o produce iodate in the region of higher pH. The effect is not due to alkali leached from the glass, as it is observed in silica as well as borosilicate glass vessels; to iodide contamination of the iodine, us it persists after repeated purification of the iodine; or to ordin a y impurities in the water, as it is observed with water redistilled from alkaline permanganate as well as with dvionixed water. The evidence that the slow developIiient of triiodide ion, which persists a t a slowly diminishing rate for a year or more, is not due to the liydrolysis to iodic acid is as follows: I n acid solutions of saturated iodine the equilibrium concentration of triiodide ion due to both hydrolyses (and especially t h a t to form iodic acid) should be sharply reduced. However, when solutions of 0.001 and 0.01M perchloric acid are saturated with iodine, the concentrations of triiodide ion observed at the outset are already in e w e s of the calculated equilibrium values and continue to increase thereafter. The rate of formation of triiodide is also greater by several powers of 10 than that computedforthe hydrolysis to form iodic acid, whether by the equation of Bell and Gelles (4)or by th:it inferred from the “sixth-order law” for the Dushman reaction described by XJ-crs and Kennedy (11 ) . A more direct piece of evidence is afforded by the spectrophotometric examination of the approach to the iodic acid equilibrium from the other side. To avoid the difficulty of mixing iodide and iodate (for sulisequent acidification) in stoichiometric proportions of the required precision, iodine was disproportionated by boiling with alkali under reflux and was then treated with perchloric acid. The amounts of iodine, perchloric acid, and water were adjusted so t h a t the final mixture contained a little more iodine than was required to saturate it and also known concentrations of hydrion ranging from 0.001 t o 0.5M. As \vas to be anticipated, the triiodide ion concentration (and the absorbance a t 3530 A , ) mas very high immediately after acidification and fell very rapidly :is the Dusli-
man reaction proceeded. What nas quite unexpected was that after a period dependent on the hydrion concentration the absorbance a t 3530 -4.passed through a minimum; this minimum was reached after a few hours a t 0.2M hydrion and after about 6 days a t 0.02M hydrion. Acidification with sulfuric acid showed behavior closely similar to that with perchloric acid. The subsequent rise in absorbance was a t a rate somewhat but not very much less than that observed starting with iodine alone in solutions of the same acidity. These observations would seem to exclude completely the possibility that the slow formation of triiodide ion is mainly due to the gradual hydrolysis of iodine to form iodic acid. The only obvious alternative to hydrolysis as a source of the unexpected concentration of triiodide ions is a reduction of the iodine by impurities. Other workers have been aware of thesensitiveness of iodine solutions to impurities, but have not addressed themselves specifically to the problem of their source. Reducing agents in the water might not be expected to survive distillation from alkaline permanganate ; the perchloric acid, made by dilution of the i270 material, is an even less lihcly source for them. Against this must be set the fact that, assuming for the purposes of a rough estimate an equiwltwt xeight of 100 for the reducing agent, an amount of the order of 0.00001 to 0.0001% by weight could produce the observed effects. Certain further observations, however, tend to support the view that the reducing is due not to dissolved contaminants but to dust adhering to the surface of storage vessels and cuvettes and surviving the repeated rinsings and other customary procedures of laboratory cleanliness. l\lien a series of stoppered cuvettes is filled with the same iodine solution, an occasional cuvette (perhaps one filling in half a dozen) nil1 show abnormally high initial absorbance a t 3530 $. followed by an unusually rapid rise of absorbance a t that nave length, whereas the absorbance in the visible, mainly due to iodine, is indistinguishable from that of the others. If a saturated solution of iodine, with or without added acid, which has been kept in a storage vessel for several months (so that its absorbance a t 3530 A. is now rising very slowly) is used to fill a series of stoppered silica cuvettes, the absorbance a t 3530 A. will rise rapidly for some days; this effect is much more marked with 10-mm. cuvettes (the rise usually corresponding to the development of between 2 and 5 micromolar triiodide in the first 2 1 hours) than with 50- or 100-mni. cuvette
