clusions can be reached about the accuracy of the method.
Table IV.
Interference
Comparative Data from Other Laboratories. Flame photometric cal-
cium determinations nere run on a series of samples which had been pieviously analyzed in other laboratories by various other methods. T h e results given in Table I11 ale a summary of these data. The type of method is indicated where it is k n o n n . Samples 1 t o F were analyzed by all of the laboratories from the same bale arid lot (alternate sheetq). The remaining (lata represent an analysis of a shipment of pulp and are not necessarily on thc w i l e lot. Some degree of variation is t o be expected in vien of the lack of homogeneity of calcium levels in the pulp. The higher levels in Laboratory A were probably due to the fact that p H control 17 as not maintained sufficiently well to preyent the coprecipitation of magnrsium. Siu samples ivhich had bren analyzed in Laboratory A were analyzed using the flamp photometric method. Nagnesium cleterminations mere made on the samples and calculated as calcium. These rewlts are summarized in the first iiu samples in Table IT’. The sums of the calcium and magnesium data are equivalent to the data found by the volumetric
Sample so.
130.1 130B 131-1 131B 213A 213B 320.1 321-1 323.i
of Magnesium in Volumetric Determination of Calcium
Calcium, P.P.M. Flame Oxalate photometer (volumetric) 161 111 113 90 129 178 113 164 110 152 112 152 30.5 29 27.2 26 21 20.1
method. This indicates that magnesium coprecipitation is probably the reason for the higher results. A new series of samples (last three samples, Table IV) was analyzed, controlling the PH more closely. The results showed close agreement when compared with flame photometric data. The flame photometric method thus hai the advantage of eliminating one of the most serious errors in volumetric calcium determinations.
Magnesium, P .P.hl., Flame Photometer 31.9 32.2 35.9 35.5 19.8 20.7 17.2 20.5 35.1
+
Calcium hlagnesium, Calcd. as Calcium 164 148
188 172 143 146
many of the analyses, is gratefully acknowledged.
I
ACKNOWLEDGMENT
LITERATURE CITED
(1) Hobbs, T. R., Pulp Paper M a g . Can. 54, XO.1, 87-90 (1953). (2) lloo_smuller, E., Das Papier 6 , 523-4 (lY5Y). (3) Richter, A . E., Scarth, J. C., Bernhardt, .I. ‘I.,Pulp Paper Mag. Can. 51, SO.3, 153-5, 158 (1950). RECEIVED for review October 30, 1956. Accepted hlay 31, 1957. Division of Cellulose chemistry. SvmDosium on New
Ultraviolet Spectrophotometric Determination of Ce r iu m(III) H. L.
GREENHAUSI, A, M. FEIBUSH?, and LOUIS GORDON3
Department of Chemistry, Syracuse University, Syracuse 7 0, Cerium(lll) can b e determined by measurement of its ultraviolet absorption peak a t 253.6 mp in 1 N sulfuric acid. Beer’s law is obeyed from 0 to 400 mg. of cerium per liter; a Ringbom plot indicates that the optimum range is 40 to 200 mg. per liter. The interference effects of ten rare earths and several other ions have been studied. The method is rapid and precise and does not have some of the instability problems associated with solutions containing cerium(1V).
A
was desired for the determination of small amounts of
RAPID m m o D
Present address, General Electric Co , Syracuse, S . Y. * Present addiess, The Texas Co., Beacon, N . Y. Present address, Case Institute of Technology, Cleveland, Ohio.
N. Y.
cerium(II1). in particular, in binary niivtures with other rare earths. Ultraviolet spectrophotometry was given strongest consideration in view of the analysis requirements. Available methods for cerium fall into two broad groups, in one of which the absorbance of a cerium(1T’) perhydrovide complex in basic solution is measured (2, 7 , 9, I S ) . One difficulty of these methods arises from the formation of a turbidity due to the instability of the complex. The turbidity effect can be eliminated by excluding peroxide, because cerium(TV) in basic carbonate solution has the same absorption spectrum as similar mlutions containing perovidr (1). I n the othrr group, the ahsorbance of cerium(1T’) is measured in acidic medium (3, 6 ) . The instability of very dilute cerium(1V) solutions in acid solution presents one difficulty, another being the complete removal of excess
reagent used to oxidize ceriuni(II1) : qome of these reagents absorb in the ultraviolet. The spectra of cerium(II1) in various media have been reported ( 1 1 , 1 2 ) ; it appeared that a useful method based on the absorbance of cerium(II1) could be developed. The method described depends on measuring the absorbance of cerium(II1) in I S sulfuric acid at 253.6 mii. REAGENTS AND APPARATUS
ilmmonium hexanitratocerate [standard of reference purity, 99.98(5)Yc], from the G. F. Smith Chemical Co., ?vas used to prepare a sulfatocerate solution ( l a ) . A cerium(II1) sulfate stock solution was prepared by reduction of the cerium with 10% hydrogen peroxide solution made by diluting 307, hydrogen peroxide in 1N sulfuric acid. This solution was evaporated to strong VOL. 29, NO. 10, OCTOBER 1957
1531
fumes of sulfur trioxide to remove nitrate. Any cerium which became oxidized was reduced with hydrogen peroxide and the solution was evaporated to the first fumes of sulfur trioxide to remove the excess peroxide. The final solution was made up to be 1N in sulfuric acid. It was subsequently standardized both by titration with a standard ferrous sulfate solution, after oxidation of the cerium with persulfate, and by the gravimetric oxalate procedure. Working solutions were prepared by aliquot dilution of this stock solution. Hydrogen peroxide, 30% Superoxol (Merck). A 10% solution was prepared by dilution with 1N sulfuric acid. Europium oxide 99.8%, gadolinium oxide 99.8%, dysprosium oxide 98%, holmium oxide 98%, erbium oxide 98%, and vtterbium oxide 99.8% (Research . CheGicals, Inc.). Lanthanum oxalate, No. 518, about 98% (Lindsay Chemical Co.). Praseodymium oxide 99.9%, neodymium oxide 99.9%, and samarium oxide 99.9% (Institute for Atomic Energy, Iowa State College, Ames Iowa). All the rare earth oxides used in the interference studies were converted to the sulfates by repeated treatment with sulfuric acid and subsequent evaporation of the solution to dryness. The final residues were dissolved in 1N sulfuric acid. All other chemicals were of reagent grade or C.P. purity. Spectra were measured with a Cary Model 11M spectrophotometer with matched fused I-cm. silica cells. A Beckman quartz spectrophotometer, Model DU, with photomultiplier attachment, hydrogen discharge lamp, and calibrated 1-em. silica cells was used for measurement a t constant wave length. Although the cells were matched in the visible region of the spectrum, there were considerable differences in their light transmittance in the ultraviolet region. Corrections were determined and applied t o all readings. I -
were measured and are shown in Table I a t 253.6 mp. The absorbances of solutions 0.05 to 3.5N remained essentially constant; a working concentration of 1N was chosen for the present work. A reagent blank was not needed, because a 1N sulfuric acid solution was transparent to ultraviolet light a t 253.6
m!J *
Table I. Effect of Concentration Sulfuric Acid on Absorbance a t
253.6 Mpa (1N sulfuric acid solution used as arbitrary reference) NorNor% mality of % ’ mality of DifferAcid Difference Acid ence 0 4.5 0.75 0 0.5 -0.25 5.0 2.0 1.5 0.25 5.8 2.2 0.25 7.2 1.0 2.8 3.5 0 8.0 0 ‘Each solution contained 80 mg. of cerium(II1) per liter.
RESULTS
Spectra in Sulfuric Acid. A series of sulfuric acid solutions was prepared, ranging from 0.05 to 8LVswhich contained 80 mg. of cerium(II1) per liter. The absorbances of these solutions 1532
ANALYTICAL CHEMISTRY
been reported to be sulfate dependent (6). The peak a t 253.6 mp has the greatest molar extinction coefficient685. This peak was chosen for the method not only because it provided the greatest sensitivity but also because it showed the least interference from most of the rare earths. Spectra of praseodymium. gadolinium dysprosium, neodymium, samarium, holmium, europium, erbium, and yttrium were obtained (4) in the approximately 210- to 370-mp range. The solutions on which measurements were made contained 1 mg. of rare earth ion per milliliter and were 1N in sulfuric acid. Adherence to Beer’s Law. Solutions of varying concentrations of cerium(II1) were prepared to determine adherence to Beer’s law and to prepare a working calibration curve.
0.500 -
-
220 250 280 310 WAVE LENGTH IN MILLIMICRONS Spectrum of cerium(ll1) sulfate in sulfuric acid
Figure 1.
100 mg. per liter
RECOMMENDED PROCEDURE
Dissolve a sample, containing 1 to 5 mg. of cerium, in sulfuric acid, and evaporate to strong fumes of sulfur trioxide to remove nitrate and chloride ions. Reduce any cerium(1V) present by carefully adding to the warm sample about 3 ml. of 10% hydrogen peroxide in dilute sulfuric acid. After initial bubbling stops, heat only to first fumes of sulfur trioxide to remove the excess hydrogen peroxide. Dissolve the residue in 1N sulfuric acid and transfer to a 25-ml. volumetric flask. Dilute to the mark with 1 N sulfuric acid. Transfer the sample to a silica cell and measure the absorbance a t 253.6 mp against a water blank.
of
The spectrum of cerium(II1) ion in 1N sulfuric acid (Figure 1) shows five peaks in the ultraviolet a t 212.1, 223.0, 240.4, 253.6, and 295 mp. Only the absorption in the region of 295 mp has
$*
U
2Q
,
40
70
I
100
200
I
I
400
,
I
700
CONCENTRATION ( MlLLlGRAMS OF CERIUM PER LITER) Figure 2.
Ringbom plots of cerium(ll1) solutions
Values at 295 mp not shown because spectrum i s sulfate-dependen
A plot of absorbance us. concentration gave a straight line from 0 t o 400 mg. of cerium per liter. A Ringbom plot (8) of per cent absorptance (100 7, transmittance) us. log concentration is shown in Figure 2. The optimum
Table II. Stability of Solutions Used in Preparation of Beer’s Law Plot
Absorbance a t Time of Preparation
Absorbance 1 Month Later
0.101 0.199 0.297 0.393 0.493 0.585 0.782 0.979
0.105 0,199 0.302 0.391 0.493 0.588 0.786 0,988
concentration range-i.e., the range in which the present method is most precise as indicated by the approximately straight-line portion of the curve-is 40 to 200 mg. of cerium per liter a t 253.6 m/J *
Cerium Taken,
Cerium Found after 1 Month,
Mg. 0.50 1 .oo 1.50 2.00 2.49 2.99 3.99 4.99
Mg. 0.53 1.01 1.53 1.98 2.49 2.97 3.97 4.99
Difference, Mg. 0.03 0.01
0.03 -0.02 0.00
-0.02 -0.02
0.00
Stability of Cerium(II1) Solutions.
The solutions of cerium(II1) which were prepared for the Beer’s law plot were remeasured 1 month after their preparation t o determine their stabilities and the amounts of cerium determined by use of the original calibration curve (Table 11). Reproducibility. Ten samples containing 2.49 mg. of eerium(II1) were prepared and diluted to 25 ml. in volumetric flasks. The cerium content was then determined spectrophotometrically. The solutions were read immediately after preparation (Table 111). To ascertain the efficiency of the reduction by peroxide, five samples were prepared containing 3.99 mg. of cerium(111) and 3.96 mg. of cerium(IV), giving a total of 7.95 mg. of cerium. The samples were treated with 5 ml. of 10% hydrogen peroxide in sulfuric acid. These solutions were evaporated to strong fumes of sulfur trioxide. A residue formed which was dissolved in 1N sulfuric acid. The solutions were heated to the $first fumes of sulfur trioxide; heating beyond first fumes may result in oxidation of cerium(II1) to cerium(1V). The cerium was then determined by the recommended procedure (Table IV). Effect of Diverse Ions. Several rare earth ions, and some other ions, were studied to determine their interference (Table V). Solutions used in the interference studies were prepared by adding the diverse ion t o a solution containing 2 mg. of cerium(111). The resulting solutions were treated with 10% hydrogen peroxide solution in dilute sulfuric acid, evaporated to first fumes of sulfur trioxide, transferred to 26-ml. volumetric flasks, and diluted to volume. Similar solutions containing cerium only were prepared t o serve as standards. The difference between absorbance readings of the sample and the cerium standard, after suitable cell corrections, was attributed to the absorbance of the diverse ion and interpreted as the effective error of that ion. The concentrations of diverse ion are given which will indicate a “cerium concentration found” 3% (an arbitrarily chosen value) greater than the 80 mg. per liter initially taken. Table V shows the order of magnitude of the errors due t o these diverse ions.
Table 111. Reproducibility of Determination of Cerium(II1) Solutions
Difference (Found Taken),
Cerium Found,“ Absorbance 0.500 0.490 0.488 0.495 0.497 0.495 0.493 0.497 0.494 0.487
Mg. 2.53 2.48 2.47 2.50 2.51 22 .. 45 90 2.51 2.50 2.46 Mean 2 . 5 0
Mg. 0.04 -0.01 -0.02 0.01 0.02 0.01 0.00 0.02
Av. diff. 0.017
0.020b
Absorbance
Mg. 8.03 7.87 7.93 7.91 0.390 7.89 . - . Mean 7 . 9 3
=
4”
(mean
Cerium(II1) solutions in 1N sulfuric acid are stable for a t least 1 month. Measurement of the ultraviolet absorption spectrum a t 253.6 mp affords an excellent method for the determination of cerium with the optimum range of 40 to 200 mg. of cerium per liter according to the Ringbom plot. Beer’s law is obeyed from 0 to 400 mg. of cerium per liter. With 80 mg. of cerium(II1) per liter, the method gives a 3y0 error with from approximately 100 to 200 mg. of praseodymium, neodymium, samarium, gadolinium, dysprosium, holmium, erbium, ytterbium, and yttrium present.
0.056
Total cerium taken, 7.95 mg.
Table V.
n
SUMMARY
Av. diff. 0.056 u
0
Interference of Diverse Ions Concn. of Diverse
- found)*
The values corresponding to a 3% error were read from graphs (4) of “error” us. “concentration of diverse ion.” I n the case of the rare earths, these curves are sufficiently linear, up to 3%, so that it is correct to assume proportionally smaller quantities of the diverse ion, as given in Table V, if errors less than 37, are desirable. However, there is some deviation from linearity for all the rare earths (except europium) above the 370 value; the direction of the deviation is such that relatively more rare earth can be accommodated as interferences above the 370 value if greater error is allowable.
Mg. 0.08 -0.08 -0.02 -0.04 -0.06
0.397 0.389 0.392 0.391
Cerium taken, 2.49 mg. ~
Difference (Found Taken),
Cerium Found,o
0.01
-0.03 u
0
Table IV. Efficiency of Reduction of Cerium by Hydrogen Peroxide
Diverse Ion La Pr Nd Sm
Yh
Ce(IV) Y Hf Zr Th Cu(I1) c1Na K NOS-
Ion Required to Give 3% Error in Cerium Concn., Mg./Liter 45 130 105 210 40 -. 145 200 120 130 200 1 140 40 35 25 10 18 ,pa0
uot ‘80 mg. of cerium(II1) per liter present. ++
* 200 mg. showed no interference. a
Interfere strongly.
Europium and lanthanum interfere more strongly than these rare earths, about 40 mg. per liter being allowable. Hafnium, zirconium, thorium, and copper interfere more strongly than the rare earths. Nitrate and uranyl ions cannot be tolerated. Sodium, potassium, and chloride do not interfere. Cerium(1V) interferes strongly but can easily be reduced to cerium(II1) with hydrogen peroxide. The method is precise and accurate, and lends itself well to rapid routine analyses of cerium. It does not VOL. 29, NO. 10, OCTOBER 1957
1533
have some of the instability problems associated m-ith the methods in which cerium(1V) or a cerium(1V) complex is the basis of the absorbance measurement. LITERATURE CITED
(1) Conca, S . , Pllerritt, C., Jr., A Y U . CHEJI.28, 1264 (1956). (2) Edwards, R. E., ilyers, A. S., Banks, C. V., Ames Laboratory Rept. ISC-165 (1951).
(3) Freedman, A. J., Hume, D. S., Ar.4~.CHEJI.22 , 932 (1950). (4) Greenhaus, H. L., M.S. thesis, Svracuse University, 1957. ( 5 ) RIedalia, A. I., Byrne, B. I., ANAL. CHEM.23, 453 (1951). (6) Sem-ton, T. IT., Arcand, G. RI., J . i l m . Chem. SOC.75,2449 (1953). ( 7 ) Plank, J., 2. anal. Chevz. 116, 312 (1939). (8) Ringbom, iZ., Ibid., 115, 332 (1939). (9) Sandell, E. B., “Colorimetric Determination of Trace Metals,” Interscience, Xew York, 1944. (10) Smith, G. F., “Cerate Oxidimetry,”
G. F. Smith Chemical Co., Colunibus, Ohio, 1942. (11) Stenger, V. A , Don- Chemical Co., Midland, JIich., private communication. (12) Stewart, D. C., University of California Radiation Laboratory. Rept. AECD-2389 (declassified 1948). (13) Telep, G., Boltz, D. F., ASAL. CHEX 25, 971 (1953). RECEIVEDfor reviepv RIarch 27, 1957, LVork supported in part by the U. S. Atomic Energy Commission under Contract dT(30-1)-1213.
Stabilization of Ferric Thiocyanate Color in Aqueous Solution Spectrophotometric Method Using Methyl Ethyl Ketone PAUL BAILY Transvaal and Orange Free State Chamber of Mines Research Laboratory, Richmond, Johannesburg, South Africa
b The ferric thiocyanate color can be stabilized in aqueous solutions by the addition of methyl ethyl ketoneacetone. The color is stable for at least 1 hour or more and is not affected b y exposure to light during this period. The optimum quantity of ferric ion is 0.02 mg. per 100 ml. of solution, using a 5-cm. glass cell at 490 mp.
T
for determining iron is one of the most convenient and is generally accepted as official in analytical practice. The distinctive color reaction between ferric and thiocyanate ions was made use of by Berzelius in 1826 ( 2 ) and was again proposed by Ossian in 1837 (6). However, without the aid of a stabilizer, the color tends to fade rapidly on exposure to light. Stokes and Cain (IO) suggested thiocyanic acid reagent stabilized IT ith mercuric thiocyanate. Potassium persulfate m-as added to oxidize the iron. Hydrogen peroxide and potassium permanganate have been used for the same purpose (7’). The extraction of ferric thiocyanate by a solvent immiscible nith water was suggested by Bernhard and Drekter ( 1 ) . They used a mixture of the monobutyl ether of ethylene glycol arid ethyl ether. The extract had a more intense color and did not fade for 24 hours. Narriott and Kolf ( 5 ) showed that the addition of acetone increased the sensitivity of the reagent. Again, although acetone increased the stability of the ferric thiocyanate color, it faded rapidly after 10 minutes. Winsor ( I d ) , n 110 usecl 2-niethosyethanol for stabilizing the color, found that this chemical HE THIOCYANATE XETHOD
1534
ANALYTICAL CHEMISTRY
was subject to photochemical change and formed a yellow color on exposure to light. Rakestraw, Nahncke, and Beach ( 8 ) used small amounts of ethylene glycol monobutyl ether to stabilize the color reaction and retard fading long enough to permit the colorimetric determination in aqueous solution. Lately, Lister and Rivington (4) carried out a spectrophotometric study of the ferric thiocyanate system. They found that benzyl alcohol was able to inhibit fading in many solutions even at 45” C. However, it was found that under the conditions used in this laboratory, 30 minutes elapsed before the color became stable. Walden (If) and Winsor (f2)studied the dielectric constants of organic liquids and their effect on inhibiting the fading of ferric thiocyanate color. As a result of their studies, the method presented here was developed. It uses a combination of acetone and methyl ethyl ketone, which proved suitable in preventing the fading of the ferric thiocyanate color so that spectrophotometric determinations could be carried out on several samples a t once. Maximum color development was practically instantaneous. REAGENTS
Ferrous ammonium sulfate solution. Prepare according to Snell and Snell (9) by dissolving 0.i022 gram of ferrous ammonium sulfate in 100 ml. of distilled water. Add 10 nil. of 1 to 1 sulfuric acid. Warm the solution and oxidize with approxiniately 0.1% .POtassium permanganate until the iron solution remains faintly pink. Cool and dilute to 1 liter. Carefully dilute a n aliquot of this solution 1 to 10 and
use as the standard solution; 1 nil. equals 0.01 mg. of ferric iron. Potassium thiocyanate solution. Add 40 grams of potassium thiocyanate (Merck I%Co., Inc.) to 100 ml. of distilled n-ater. Hydrochloric acid, reagent grade. 2N. British Drug Houses reagent was used. Acetone, British Drug Houses Analar reagent. Methyl ethyl ketone, British Drug Houses laboratory reagent. APPARATUS
All spectrophotometric measurements were made with a Beckman Model B spectrophotometer. Glass cells of 1.0 cm. and 50-cm. light paths fitted with glass stoppers M-ere used. A Beckman Xodel G p H meter was used for p H measurements. PROCEDURE
Measure a n aliquot of the iron solution (1 to 20 ml. containing 10 t o 200 p.p.m. of iron) into a 100-ml. volumetric flask. Dilute to almost 30 ml. with distilled water. Add 2 ml. of 2N hydrochloric acid, so that the p H of the resulting solution is approximately 1.5. Measure 20 ml. of acetone, folion-ed by 40 ml. of methyl ethyl ketone. into the flask and mix the content;. Finally, add 5 ml. of the potassium thiocyanate solution from a pipet. Make up to volume with distilled nater and mix the contents of the flash welC Measure the transmittance a t 490 m u . DISCUSSION
Optimum Wave Length. Figure 1 s h o w that the optimum wave length for this solution is a t 490 nip. A satisfactory u-orkable concentration range
