(15) Ibid., p. 718. (16) Plotnikova, 0. M., Lysenko, V. I., S b . T r u d y Vses. ~9’auch-Issled.; Gorno Metallurg. Inst. Tsvet. M e t . pp. 339-42, 1962; A n a l . ‘4bstr. 10, 4632 (1963). (17) Pollock, E. X., Ziopatti, L. P., i l n a l . C h i m . Acta 28, 68 (1963). (18) Polyakova, V. V., Fedorova, V. V., Spektr. ilnaliz. u . Tsvetnoi M e t . Sbornik, pp. 172-5, Metallurgizdat, Moscow, 1960; A n a l . Abstr. 8, 3695 (1961). (19) Ringbom, A., Z . A n a l . C h e m . 115, 332 (1939). (20) RPliiEka, J., Star:?, J., Talanta 9, 617 (1962).
(21) Sandell, E. B., “Colorimetric Determination of Traces of Metals,” 3rd ed., p. 443, Interscience, New York, 1959. (22) Silverman, L., “Handbook of Analytical Chemistry,” pp. 13-30, L. Meites, ed., McGraw-Hill, Sew York, 1963. (23) Smith, G. F., McCurdy, 1%‘. H., ANAL.CHEM.24, 371 (1952). (24) Turkington, R. W., Tracy, F. M . , Ibid., 30, 1699 (1958). (25) White, J. C., C. S . A t . Energy Comm. Rept. CF-59-4-106,April 21, 1959.
(26) Yuasa, T., J a p a n Analyst 11, 359 (1962); A n a l . Abstr. 11, 119 (1964). (27) Yudelevich, I. G., Shokarev, hl. &I., Sosnovskaya, T. I., Stanevich, V. V., Alontseva, S . T., Sb. A-ekotoryye V o p r . Emission. i Molekulyarn. Spektroskopii, pp. 126-33, Krasnoyarsk, 1960; Selenium-Tellurium Abstr. 4, 1356 (1963). RECEIVED for review February 27, 1964. Accepted May 26, 1964. Work was supported by the Defence Research Board, Department of National Defence, Ottawa, Canada, under Project E-19 of the Defence Industrial Research Program, 1963.
Titrimetric: Determination of Semimicro Amounts of Sulfate in Presence of Phosphate BRUNO JASELSKIS and STANISLAUS F. VAS Department o f Chemistry, loyola University, Chicago, 111.
b Application of barium iodate as a reagent for the determination of sulfate in the presence of phosphate is investigated as a function of pH, solvent composition, and equilibration conditions. Sulfate solutions are equilibrated with solid barium iodate in a 40% acetone and 0 . 1 M acetic acid aqueous mixture. l h e liberated iodate is determined iodometrically. Amounts as low as 0.9 mg. of sulfur in the presence of 6 . 0 mg. of phosphorus can b e determined with a relative standard deviation of less than 1%.
I (4)
Macdonald has extensively reviewed various titrimetric methods for the determination of sulfate. The estimation of sulfate by redox tjt’ration methods, after precipitation as barium sulfate with slightly soluble salts like barium iodate or barium chromate, was suggested by hndrews ( I ) and Sojbelman (j), and has been dt:,scribed in greater detail by Erdey and I?myai ( 2 ) . In the presence of appreciable amounts of phosphate, however, the determination of sulfate becomes complicated because of the formation of slightly soluble barium phosphate. Thus, a preliminary separation of phosphate from sulfate is essential and can be effected either by ion exchange or by the reduction of sulfate to hydrogen sulfide. The interference of phosphate in the precipitation of sulfate as barium sulfate also can be rninimized by increasing the hydrogen ion concentration and keel)ing the barium ion concentration low. These conditions are obtained when a slightly soluble barium salt, such as barium iodate, reacts with an acidic solution of isulfate and phosN A RECEXT PUBLICATION
phate. The iodate ion thus liberated and determined iodometrically has the advantage of providing a twelvefold amplification factor over direct titration methods. The present paper attemptq to establish not only the optimum conditions for the determination of sulfate in the presence of phosphate JT ith bolid barium iodate, but also the effects of certain experimental variables on the final result. EXPERIMENTAL
Apparatus. The samples were agitated with a n Eberbach electric shaker. A Beckman Model G pH meter was used for the adjustment of pH. Materials. Reagent grade barium iodate was obtained from Matheson Coleman &: Bell. Baker and Adamson reagent grade anhydrous sodium sulfate was used for the preparation of t h e standard sulfate solutions. T h e other reagents used were also reagent grade, except the commercial dioxane which was purified by distillation over sodium and stored over sodium wire in a dark bottle. Procedure. DETERMINATIOK OF SULF.4TE I N P R E S E N C E O F P H O S P H A T E .
A series of sulfate solutions were prepared by taking known volumes of standard 0.01JP sodium sulfate in 25-ml. volumetric flasks. Five milliliters of 0 . 5 X acetic acid and 0.5M sodium nitrate (the latter to keep th.e ionic strengt,h approximately 0.1)) followed by 10.0 ml. of acetone, were added to each flask, and t’he volume was made u p to the mark with distilled water. .lfter being mixed, the contents of each flask were t r m s ferred (without, rinsing the volumetric flask) into a 50-ml. Erlenmeyer alas.stoppered flask, and spl)roaimately 0.2 gram of powdered barium iodate was added. The flasks were then agitated in an Eberbarh shaker. -4fter 2 hours
of agitation, an additional 50 mg. of barium iodate were added and the agitation was continued for another hour. The flasks were then placed in a constant temperature bath at 25’ C. for one hour. When the precipitate had settled, the supernatant solution was decanted and centrifuged for about 5 minutes. A 20-ml. aliquot, of the clear solution was then pipetted out. To prevent any solid barium sulfate or barium iodate from being sucked i n , a piece of Tygon tubing plugged with borosilicate glass wool was fitted t,o the tip of the pipet. The amount of iodate in the aliquot was then determined iodometrically . The possible interference of phosphate and other anions like chloride, bromide, perchlorate, and nitrate, in t’he presence and absence of sulfat,e,was checked by addition of known concentrations of these ions to the solution before equilibration with barium iodate. The effect of varying concentrations of acetone, ethyl alcohol, tert-butyl alcohol, and dioxane was studied by determining the solubility of barium iodate in solutions containing these solvents in the presence and absence of sulfate. The optimum equilibration time was obtained by estimating the amount of iodate liberated. a t different time int’ervals. The solubility of barium iodate in pure water and in acetonewater mixt,ures, in the presence and in the absence of phosphat,e, was determined a t various pH’$ keeping the ionic strength a t about 0.1. The temperature dependence was studied by changing the temperature a t various stages of equilibration. RESULTS AND DISCUSSION
Solubility of Barium Iodate Alone and in Presence of Phosphate as Function of pH. T h e solubility of
barium iodate a t constant temperature remains essentially constant in VOL. 36, NO. 10, SEPTEMBER 1964
1965
Total Iodate Concentration a s Function of Sulfate Present. The total concentration of iodate a t equilibrium in the qupernatant solution deliendnot only on the amount of barium sulfate precipitated but alio on the barium iodate 5olubility a t a given iodate concentration Totai iodate concentration, T , at equilibrium with the amount of sulfate, C90,-2, added can be expressed b j the equation
f I
!
I
I
I
I I
I !
I
I
I I
0 I
2Ki/Kz(H+,iKi
2C,,,-zT'
I
/
+ 1)2TJ+ T3 -
2R,(H+/K,
/
-
+ 1)*
=
0
(1)
where K I and K 2 are the apparent solubility products of barium sulfate and barium iodate and K , is the ionization constant of iodic acid. At pH's higher than 3 the term (H+,.'K, 1) is approximately equal to 1 and can be neglected. Furthermore, in the presence of sulfate the contribution of the term 2K1!62(Hi 'R, 1)2 T4 is negligible, and a simplified working equation is obtained :
+
1
!
0.0
2.0
1 4.0
1
I
I
60
80
10.0
P"
Figure 1.
Solubility of barium iodate
-0
Barium iodate alone Barium iodate in 0.02M 0- -0 Barium iodate in 0.02M and 40% acetone-water mixture
A- - - A
phosphate phosphate
the pH range 2.5 to 9.5. At low pH's t h e solubility of barium iodate increases as a result of the formation of the HIOB species (KHIoa = 1.9 X IO-'), and it also increases in basic solut,ions because of the formation of basic barium iodate salts. In the presence of 0.025f phosphat'e, the solubility of barium iodate increases with increasing pH. In the pH region 2.5 to 5.5 only, 0.02M phosphate does not interfere, The increase of iodate concentration is caused Sy the formation of insoluble barium phosphate. The effect of pH on the solubility of barium iodate alone and in the presence of phosphate is shown in Figure 1. Solubility of Barium Iodate as Function of Solvent Composition. Appreciable solubility of barium iodate in aqueous solution produces high b ~ a i i k s and renders the titrimetric method useless for determination of sulfate a t l o a concentrations. T h e s o l u h i l i t ~ can ~ be decreased by the introduction of organic solvents. hcetonc, tert-butyl alcohol, anti p-dioxane have ljeen srlecteci because they are commei~c~iallyavailahle and are not ositiized by icdic acid. The effect of the concentration of these solvent. on the solubility of bai.ium iodate is determined by the method of .!entoft and Robinson ( 3 ) . .ketone has been selectcd on the basis of availability and ease of Iimdiing. The optimum concentration of acetone is about 4070 acetone. Be>.ontl this point, further addition of acctonc cat1 change in the s0lul)i iodate. 1966
ANALYTICAL CHEMISTRY
+
T 3 - 2Cso,-zT2 - 2K2
=
0
(2)
If the concentration of sulfate in the equilibrating solution is higher than 2 X 10-3M and if the solubility of barium iodate is reduced by the addition of organic solvents or cooling, the term 2 K 2 becomes small with respect to 2CSo,-?:'T2 and can be neglected. Thus one obtains a linear equation : T =
(3)
The dependence of the total iodate liberated on the sulfate concentration added is shown in Figure 2. I n aqueous solutions the amount of liberated iodate is not directly proportional to the concentration of sulfate and can be calculated by the Equation 2. However, nonlinearity can be diminished, to a limited extent, by reducing the barium iodate solubility by organic solvents and by lowering temperature, as shown in Figure 2 , curves B, C: and D. In 40% acetone, 0.1-11 acetic acid, and 0.131 sodium nitrate niisture, the sodium iodate is al)prosirnately 7 5 times more soluble than the barium iodate, and thus with increasing sulfate concentration the interfei,ence of sodium ions become? irnlmrtant. This produces recoveries of sulfate lower than 1007c. Temperature Effect on Amount of Iodate Liberated. The temperature effects are rather pronounced on the re1)rotiucitiiiity of results. I3ariuni iodate solu hi lit y h an &e* all pr o simately by R factor of 3 barium sulfatc solubility ?hang(,> by 2 . 4 i in the tewpcrxture range of 0" to 30" C. The equilibrium iodate c-oncentratioii i l estnl)li.shetl in lei> than 3 hours. while that for hariuni (3
Figure 2. Total iodate concentration as function of sulfate present
.--e Theoretical curve a t p H 3 in aqueous solution A - - A Observed valves at pH 3 in aqueous solution 8 . 0- -0 In 30% fert-butyl alcohol, p H 3 and 27' C. C. 0-0 In 30% tert-butyl alcohol, p H 3 and ' C. cooled after equilibration to 0 D . 0-g In 40y0 acetone-water solution, pH 3 and 27' C. A.
sulfate is not achieved even in 2 days. This may lead primarily to two types of errors: those arising from higher initial and equilibration temperatures than those of the analysis, and those arising from higher analysis than equilibration temperatures. Higher initial and equilibration temperatures than t'hose of analysis yield a higher concentration of barium ion, and thus a more complete barium sulfate precipitation and a larger amount of iodate liberated. Analysis a t higher than equilibration temperature yields high blanks and, on the whole, poor reproducibilit'y, and should be avoided. Amount of Iodate Liberated as Function of Equilibration Time. The amount of iodate liberated in a given time depends on the particle size, the solubility of barium iodate in a given solvent, and the concentrat'ion of sulfate. Equilibration times become shorter the smaller the particle size and the greater the solubility of barium iodat,e. At pH 3 in aaueous solution, 90% of iodate is liberated in less than one hour, while in 40yo acetone-water solution, two hours are necessary. For low concentrations of sulfate in a 40yo acetone-warer solution, an equilibration time of 3 hours is sufficient. If the amount of sulfur in the equilibrating solution is higher than 1 mg. I)er 25 ml., approsinlately 50 mg. of barium iodate must he added a t the end of 2 hours and agitation must be continued for another hour.
Table I. ~~
Concentration in equilibrating s o l u t i o q Sulfate, Phosphate, m 11 J!l 0 000
a
01
0 000 0 360 0 720 1 440 1 430 2 160 2 880 4 320 4 320
0 0 0 0
5 760
0 01
01 01 01
0 01 0 01 0 01
Sulfate Analyses Results
Thiosulfate used,a i d . 0 58 i 0 O4* 0 60 i 0 04 1 66 i 0 05 2 81 i 0 05 5 49 f 0 06 5 47 f 0 06 8 03 i 0 06 10 35 i 0 06 15 14 i 0 07 15 15 i 0 07 20 10 i 0 07
Sulfate, mmole Taken Found 0 0 0 0 0 0 0 0 0
0090 0180 0360 0360 0540
720 1080 1080 1440
0 0 0 0 0 0 0 0 0 0 0
Recovery,
%
0040
0041 0115 0195 0380 0381 0556 0717 1049 1051 1393
127 70 108 33 105 55 105 83 102 96 99 59 97 13 97 32 96 74
Thiosulfate used for 20 00-nil aliquot taken out of 25 00-ml volume of the equi-
librating solution The reported values of milliliters are the average of five determinations Sormality of thiosulfate is 0 06654 * Std dev for five determinations
less than 1%. Calibration of the method against the known sulfate concentrations reduces errors caused by equilibration and handling procedures. Interferences. Chloride, bromide, perchlorate, and nitrate in concentrations of 0.351 or less do not interfere. Phosphate, if present in higher concentrations than 0.0231, interferes and yields somewhat higher values t h a n those theoretically predicted Cations, which form insoluble iodates or which can oxidize iodide to iodine, must be eliminated. This can be done by passing the solution through a cation exchange resin in a hydrogen cycle and the resulting solution containing sulfate and phosphate can be analyzed successfully. LITERATURE CITED
(1) Andrew, L. JT’
Evaluation of Results. Evaluation of results is based o n the analyses of five replicate samples a t w r i o u s concentrations of sulfate and phosphate in 40yc acetone-water solutions. T h e results for the determination of sulfate are summarized in Table I. Overall recovery of sulfate varies with concentration in the equilibrating solution. However, the volume of thiosulfate used in the titration of iodate
that is liberated is directly proportional to the sulfate concentration, and thus the unknown sulfate concentration can be determined from the calibration curve. By this method, if temperature is controlled to 0.5” C., amounts of sulfate as low as 0.01 mmole (0.32 mg. of sulfur) per 25 ml. of equilibrating solution can be determined within a relative standard deviation of 2%) and at higher concentrations than 0.02 mmole, within
J . A m . Chrm. 11, 567 (1890). ( 2 ) Erdey, L., Banyai, Eva, Z . ilnal. Chem. 161, 16 (1958). (3) Jentoft. R. E., Robinson. R. J.. ASAL. CHEM.26. 1156 11954). ( 4 ) lfacdonald, A,’ M. G., Ind. Chemist 36, 345 (1960). ( 5 ) Sojbelman, B. I., Zh. Analit. Khim. 3 , 258 (19483. ~
RECEIVEDfor revie)? Alarch 20, 1964. Accepted June 12, 1964. Q’ork supported by a Frederick Gardner Cottrell grant in aid from the Research Corp
Densimetric Method for Characterizing Asphalt 1. W. CORBETT Esso Research and Engineering Co., linden, N. J .
b A relatively rapid and simple method for characterizing asphalt i s described. It i s accomplished b y first separating asphalt into asphaltenes and petrolenes, then applying densimetric techniques to the latter fraction. This characterizes petrolenes, the component or fraction grossly responsible for an asphalt’s physical and chemical properties. The results are expressed in terms of the fraction of carbon atoms in aromatic rings, the number of aromatic and naphthene rings per molecule, and a characterization index related to the degree of ring condensation. The densimetric technique was extended from published works b y Van Krevelen and b y Williams, and involves a calculation based upon the relationship between molar volume and atomic H/C ratio. It requires only the measurement of per cent carbon, per cent hydrogen, and the molecular weight of the petrolene fraction. This work demonstrates its repeatability and its applicability to straight reduced stocks from 15 crude
sources and to more detailed fractions obtained b y chromatography and molecular distillation of asphalts. Because of the simplicity of this method, it should prove useful in those laboratories wishing to follow changes in asphalt during manufacture, and service, and possibly for fingerprinting purposes.
&!t
.my IUETHOIIS have
been proposed for composition analysis of asphalt, but most of these are quite laborious and detailed. I n many instances laboratories are not equipped or cannot devote the time necessary for lengthy procedures, and therefore composition analyses are not carried out even when they could be the solution to a problem For these reasons, a relatively simple method for characterizing asphalt has been worked out with the hoile that it will fill this need. =isphalt is a w r y complex mixture of high molecular weight hydrocarbons of variable qize and type together with
sulfur, oxygen, and nitrogen compounds. This complexity is generally reduced by first making soine kind of a major separation. At this point a decision must be made as to the time and effort that can be devoted to further sellarations. To keep it simple, the first and most logical step is to separate the asphalt into its two structiirally distinct components, asphaltenes and petrolenes. This can he accomplished by solvent precipitation (deasphaltening) using a solvent such as n-hexane. I t has been shown that the petrolenes ( I , 4, j), are more variable in character than asphaltenes and have a greater influence on the overall properties of an asphalt. Asphaltenes are the dispersed phase of the colloid ( 6 ) and are a minor proportion of the asphalt. Although both components are necessary for the makeU ] J of an asphalt, the characterization of the petrolenea contribute> a great deal more toward characterizing the nhole asphalt. This work s h o w that petrolenes may be characterized by applying an anal) VOL. 3 6 , NO. 10, SEPTEMBER 1964
1967
