tions can be kept for several days nithout difficulty, but are very sensitive toward hydrolysis to heat and light (the hydrolysis is being investigated further). Hydrochloric acid in excess is undesirable because of formation of iron (111)-chloro complexes ( 1 ) Lvith consequent unfavorable equilihriuni effects. Use of perchloric acid instead of 11)-drochloric acid \vas investigated; the precipit’ationof trtraniethylammonium perchloratr. with subsequent extensive adsorption and only slow desorption of iodine vitiated this procedure. Similarly, tetramethylammoniuni iodide is precipitated if more than very small amounts of potassium iodide are added initially. .Is normally carried out. the titration requires sufficient time to allow the reaction to occur n.ithout
difficulty, completely desorbing any iodine into the solution prior to the end point, which is quite sharp and can be estimated within a drop visually or determined more accurately from a dead-stop plot of current 1’s. volume. The direct titration of the tetrachloroferrate(II1) tvith thiosulfate was found to be unsatisfactory due to the slon-iies3 of the reaction. The thiosulfato-iron(II1) coniples which forms first ( 5 , undergoes further reaction only very slon-ly. ACKNOWLEDGMENT
The author thanks Laurel Kilkening for preparing one of the compounds used in this study and R. S. Juvet for helliful suggeqtions concerning the manuscript.
LITERATURE CITED
J., Schwarzenbach, G., Sillen, L. G., eds. “Stability Constants of Metal-Ion Complexes, Part 11-Inorganic Ligands,” p. 97, Special Publiration So. 7, The Chemical Society (London), 1958. ( 2 ) Kolthoff, I. 11.;Pandell, E. B., “Textbook of Quantitative Inorganic rlnalgsis,!’ 3rd ed., Chap. XVII, blacmillan, Tork, 1952. (3) Kolthoff, I. Il.>Sandell, E. B., Zbid., (1) Bjerruni,
n. 579r
(4) Meek D. IT..Drago, R. S., J . Am. Chem. SOC.83, 4322 (1961). ( 5 ) Page, F. ll.,Trans. Faraday Snc. 50, 120 (1954 1 . L. J. SACKS Cheniistri- Departnient Reed Cdlege Portland 2 , ( )re. R E C E I ~ Ef oDr review December 18, 1962. Resubmitted qTilile 10, 1963. -4ccepted Jiine 21 l’Ih.3.
A Fluorometric Procedure for the Determination of Cerous Ion SIR During an inr e-tigation of the aqueou- ceric wlfate do-imeter ($), it became nece-ary to determine cerouion concentration. helon. 10-5 gram ion per liter in the presence of about 10-4 gram ion per liter of ceric ion. Ailthough no witable method could he found in the literature. the reported fluore-cence of
Ce(II1) solution. ( 2 ) suggested the development of a fluorometric technique. The technique del eloped, while unsuitable for dosimetry, does provide a method for the determination of a range of Iariou- ion concentrations. either in the a b e n c e or presence of ceric ion, not c o ered ~ by a n y other analytical technique EXPERIMENTAL
Ce2(S04)3.8H20. G. Frederick dniith “Reagent Grade” used. The cerium content was determined by the standard method of ovidation with ammonium persulfate followed b y titration with standardized ferrous ammonium sulfate. T h e purity of the cerous sulfate was found t o he 93.3%. This result was confirmed by measuring t h e abqorbance at 254 nip of a number of cerou. \ulfate-O.4.lf &So4 solutions and calculatiiig the cerous ion concentration u.ing the literature value of 685 for the niolar ah.orptivity at 254 mp ( I ) . Thi. purity correction was applied to all relevant data. The deviation from 100% purity was not due to water alone a. the weight loss meawred after heating 48 hours a t 125‘ C. indicated the lire-ence of only 1.8% absorbed water. (TH4)4Ce(S04)4. 2H20. G. Frederick Smith “Reagent Grade” material waused. sample was standardized against A.R. arsenious oxide and found to be 99.9% pure. Aifurther check on the relative purity of the cerous and ceric ammonium sulfates waz made by giving a 2 x l O - 5 M ceric ammonium sulfate solution in 0.421 H2SO*a dose of CON yradiation sufficient to reduce all the ceric ions and then comparing Reagents.
Table I. Cerous Ion Fluorescence as a Function of [Ce(lll)]
[Ce(III)j gram ions per liter X lo6 1.0 5.0 10.0 20.0 50.0 100 300
220
Table 11.
[Ce(Iy)I gram ions per liter
x
10‘
0 188 0.188 0.188 0.188 0.188 0.188 2.89 2 89
2.89 ?.S9
1300
Fluorescence intensity arbitrary units 1.0 5.0 9.9 20.0 48.0 95
Independence of Q on [Ce(III)j
[Ce(IIIIi gram ions per liter x 1oj 0.280 0.467 0.933 1.87 4.67 9.33 0.467 0.933 1.87 4.67
Q 1 19
i.iu
1.16 1.17
1.18 1.13 11.7
11.4 1.3
11.3
ANALYTICAL CHEMISTRY
the fluoreccence of this solution with that’ of a Ftandard cerous sulfate-0.4M H2P04 2solution. dssuming the ceric animoniuni wlfate to be lOOyo pure, the cerous sulfate \vas found to be 92.57, pure. a value in good agreement n-ith that obtained by conventional
-111 solutions were prepared using triply distilled water (one distillation from alkaline KRIn04) and were made 0.4111 in sulfuric acid. Procedure. Fluorescence measurements were made in 1-em. quartz cells using a n Iniinco-Bowman spectrophotofluorometer with a Xenon lamp. The excitation and fluorescence wavelengths employed were 254 and 350 nip. respectively. The Q factors were determined from comparisons of the fluorescence intensities of standard cerous sulfate solutions in the absence and presence of ceric ammonium sulfate ( I s A and IBP, respectively). Since cerous ions were always present to some extent in Ce(1V) solutions, the fluorescence intensities of ceric ammonium sulfate solutions were ako measured to obtain appropriate blank corrections (IB).Then Q is given by I ~ A / (-I IB). ~P A Beckman DU spectrophotometer \vas used for all absorbance measurements. The ceric ion concentration in each solution was determined from its absorption a t 320 mp using the value of 5580 for the molar absorptivity (3).
A \
RESULTS A N D DISCUSSION
The fluorescence of cerous sulfate0 . 4 X H2S04solutions increases linearly with increasing cerous ion concentration over t8he range 10” to 10-4 gram ion per lit,er (Table I). Within these limits
cerous ion, in tlie absence of a n y int,erfering species, may be determined with a n accuracy of *57', by comparing its fluorescence intensity with that of a cerous sulfate standard. ;It higher crrous ion cencentrations, ,self-quenching apparently becomes important. Ceric ions presumahly deerewe cerous ion fluorescence by 'ilmrhing some of the exciting and emitted light. Calculations of the factor Q, the ratio of Ce(II1) fluorescence n tlie al)\ence and presence of Ce(IV ' , hayet-l on the Ce(IT') molar absorptivitie- a t 234 and 360 mp agree fair])- r w l l n-it11 the experimental results. FOY2 constant' ceric ion concentration Q i, independent, within tlie ~~sl)erimental error, of cerous ion conw itration over the to lo-' gram ion per range 3 X liter (Table 11). - i t ~~JJIW C't,!,III! conrentrations there is . ~ i irviclcnce ~ to suggest t,hat Q iiicrw-e- 4qnificaiitly. Log Q increahes linearly lr-itli increaqing Ce(IT7) concentration m.er tliP range 0.2 to 3.0 X g:r:iin iiin ]lei lit'er (Table 111). The equation log Q = 3760 [Ce(IV)] niay be deriwd i'ronl the average of the log Q ' [ C ' F ( I T J j ratios. I X n g this equation nitliin the limit? tievrilied. rerow iori c~riicc~ntratirlnof
Regarding the application of thia method to the analysis of Ce(II1) In the preyenre of cpecies other than Ce(IV), it should be noted that a n y substance 14 hich abiorbs near 250 or 350 mp mill reduce Ce(I11) fluorescence. The effect of each iubbtance present sould hai e to be in1 e+gated before proceeding nith the anal? +.
Table 111.
Dependence of Q on [Ce(lV)]
[WIy)1
gram ions 0 0 0 0 0 0 0 0 0 0 0 0 0
145
0 0557 0 0696 0 0774
188
242 484
676 686 728 773
3h40 3 700
0.167 0.268
:3200 3450 3060 3320
0,260 0.301
:3580
0,228
790
0.323
977
1 02 .~
0.330 0.39" 0.374
1.07
0.417
1.15
0,446 0.433 0 . 720 0 . 731 0.703 0.924 1 .OS6
821 872 879
1.16 1 85 1.94
1.94 2.42 2.89 2.89
1 ) (;reenhaus, H.
I
