ANALYTICAL CHEMISTRY
418 case, S - B is applicable; in the second, B - S . Generally, a t 25" C. and 760 mm. of mercury, the following expression applies: titer difference X N X 24.5 X 1000 (5) VC where N = normality of standard titrant with respect to contaminant 24.5 = liters per equivalent a t 25" C. and 760 mm. of mercury Vc = corrected volume of gas sample, liters P.p.m. by volume =
Slide rules may be constructed specifically for solving this equation for parts per million by volume of the individual air contaminant under analysis. The equation is most conveniently rearranged t o
vc 24.5 X titer difference
-
1000 X N p.p.m. by volume
0,001. This will satisfy the requirement that individual scales in a scale pair have the same linear unit. Volumes may be plotted in the range 10 to 1000 liters while titer difference (T.D.) is plotted as T.D. X 24.5 but marked as T.D. In this manner, differences in thP range 0.4 to 37 ml. may be expressed on the scale. Volumes and titer differences are plotted on the inside, sliding scales. Slide rules fabricated of cardboard and log paper in this laboratory have an accuracy of within 1%. As most industrial air analyses are performed in the range 1 to 100 p.p.m. by volume with an expected accuracy of =k2 p.p.m. using wet-test meters whose accuracy may be as variant as +5%, use of this rapid slide rule method for volume corrections and concentration calculations seems justifiable.
(6)
If two-cycle semilog paper is used, parts per million by volume in the range 1 to 100 and normality of titrant in the range 0.001 to 0.100 N may be plotted on the outside, stationary scales. The normality is plotted as 1000hi,but marked as N-i.e., a N of 0.001 would be plotted as 1000 X 0.001 = 1but marked on the scale as
LITERATURE CITED
(1) Barr, G.,J. soc. Chem. Ind., 49,21-T (1930). (2) Eshbach, 0. W., "Handbook of Engineering Fundamentals," pp. 2-121,New York, John Wiley & Sons, 1936. (3) Patton, T.C., Chem. & Met. Eng., 41,488 (1934). (4) Roof, J. G., I n d . Eng. Chem., 32,998 (1940). RECEIVED May 14,1951.
Ascorbic Acid in Analytical Chemistry Determination of Ferric Ions LA'SZLO ERDEY AND ENDRE BODOR Institute f or General Chemistry, Technical University, Budapest, Hungary ANALYTICAL chemistry ascorbic acid has usually been IhasNemployed only in an indirect way; the product reduced by it been determined or the excess of ascorbic acid used in the redox reaction has been estimated iodometrically and the amount consumed by the oxidizing agent found by difference. A method for direct application of ascorbic acid has been developed by Ptitsyn and Kozlov ( 6 ) ,who employed freshly prepared ascorbic acid solution for potentiometric determination of ferric(II1) ions. The authors have been able to reproduce their results. Information was received after this work was completed that who the reverse procedure has been applied by do Xascimento (4), determined ascorbic acid with ferric chloride. The great reducing power of ascorbic acid suggested the idea of trying its use as a volumetric solution. The redox potential of a solution containing ascorbic and dehydroascorbic acids has been compared to a saturated calomel electrode and found to be 0.433 volt a t 21" C., which, referred to a normal hydrogen electrode, is equal to +0.185 volt. This value approaches the numerous but rather divergent data found in the literature (6). Ascorbic acid reacts with weak oxidizing materials, producing dehydroascorbic acid. Strong oxidizing reagents, such as potassium permanganate, may also oyidize dehydroascorbic acid. Dehydroascorbic acid is more easily oxidized in alkaline than in acidic medium. The oxygen of the atmosphere acts catalytically to destroy ascorbic acid, and it is especially affected by traces of heavy metals. ASCORBIC ACID AS MEASURING SOLUTION
As the gram equivalent of ascorbic acid in the reduction-oxidation process is 88.03, approximately 8.9 grams of ascorbic acid were used for preparing 1liter of 0.1 N volumetric solution. The water was freshly distilled from a glass apparatus. The normality was checked with 0.1 A- iodine solution as well as with 0.1 N potassium iodate solution. When iodine solution was used, 5 ml. of 2 N hydrochloric acid were added to 20 ml. of ascorbic acid solution and the cold solution was titrated without any indicator. I n the iodate titration, 1 gram of potassium iodide was added to 20 ml. of 0.1 A' potassium iodate, acidified with 5 ml. of
2 N hydrochloric acid, and titrated with ascorbic acid until the yellow color of the iodine disappeared. The use of starch indicator was eliminated, as it decreased the rate of reaction, and a distinct end point may be observed without indicator. As the values obtained with the two different methods were in good agreement within the limits of experimental error, only potassium iodate was used in further checking; an exact and constant 0.1 A- solution is more easily prepared from it. I t may be concluded that the stability of the ascorbic acid solution depends on whether it is kept in an inert atmosphere or in
4 Figure 1
V O L U M E 2 4 , NO. 2, F E B R U A R Y 1 9 5 2
419
made containing equivalent amounts of dehydroascorbic acid as well as Days 1 2 3 4 5 6 7 10 15 20 60 ferrous and ferric ions. The solution Stored in air was kept for different periods and a t Factor 0.999 0 . 9 9 6 0 . 9 9 3 0 . 9 8 9 0 . 9 8 4 0 . 9 7 8 0 . 9 7 0 0.960 0.949 0 . 9 3 0 0 . 7 1 7 various temperatures in neutral and Diminution, % . . . - 0 . 3 -0 6 - 1 . 0 - 1 . 5 - 2 . 1 - 2 . 9 - 3 . 9 - 5 . 0 - 6 . 9 -28.2 slightly acidic media, and the amount Stored in COS Factor 0.999 0.996 0 . 9 9 3 0 . 9 9 0 0.990 0.989 0.989 0.989 0 . 9 8 7 0 . 9 7 9 0 . 9 4 3 of ferric ion remaining was determined Diminution, % . . . - 0 . 3 - 0 . 6 - 0 . 9 -0.9 -1.0 -1.0 -1.0 - 1 . 2 - 2 . 0 - 5 . 6 with ascorbic acid in the presence of potassium sulfocyanate. It was established that increase of temperature above 60" C. is not advisable, as air. The normality of a volumetric solution kept without any a t higher temperatures a considerable amount of ferric ion was particular care showed an average decrease of 0.3 to 0.4% per day reduced by dehydroascorbic acid. The acid concentration should during 2 weeks (Table I). This deterioration could not be be a t least 0.1 N , as the reducing effect of dehydroascorbic acid eliminated by boiling the distilled water, by keeping the solution is smaller in an acidic medium. At temperatures below 60" C. in a dark bottle, nor by storing it in the cold. Accordingly, if and in slightly acidic media only a negligible amount of the solution is kept a t room temperature in a bottle with a glass ferric ion was reduced by dehydroascorbic acid during a 5-minute stopper, the normality should be checked daily or even twice a titration. day, if it is used frequently. Experiments were also carried out to establish the effect of the concentration of acid, indicator, and dilution on the end point. No change in end point was noted when the concentration of acid was less than 0.5 hi. The change of indicator concentration 'O0 between 0.005 and 0.0005 N had no influence on the end point. It was established that highest accuracy could be obtained with -point ferric ion concentrations between 25 and 250 mg. per 100 ml. The effect of atmospheric oxygen was studied in parallel tests made in a carbon dioxide atmosphere. It appeared that atmospheric oxygen does not affect the results to any extent. This fact provides a great advantage compared to titanometric (3) 50. and chromometric ( 7 ) determinations of ferric ion. I/ Ir For establishing the accuracy of the method, comparisons were made with other methods applicable to the determination of 20 30.35 ferric ions. The same solution containing ferric chloride in a ml: O h uscor& acid concentration about 0.1 hi was analyzed gravimetrically, by the Figure 2 methods of Zimmermann-Reinhardt and Mohr, titanometrically, and by titration with ascorbic acid. The experimental data are When the volumetric solution was stored in a carbon dioxide presented in Table 11. Deviations are smallest in the estimaatmosphere in a Winkler-type tank buret connected with a cartions made with ascorbic acid and the data agree completely with bon dioxide generator (see Figure 1)its normality decreased only those obtained titanometrically. in the first 2 days and became constant for 2 weeks thereafter To determine the accuracy of the method, the mean error of (Table I). This indicates that the deterioration was caused the averages of twenty parallel assays was determined and found chiefly by atmospheric oxygen. to be 0.004% (Table 111). The change of efficiency depends also upon the quality of the According to the authors' experience, ferric ions can best be solid ascorbic acid used for preparing the solution. determined as follows: Table I.
Changes in Efficiency of 0.1 N Ascorbic Acid Solution
/
I ! ro
TITRATION O F FERRIC IONS WITH 0.1 N ASCORBIC ACID SOLUTION
The reaction of ferric ions with ascorbic acid is slow a t room temperature, especially a t the end of the titration. With heating, the rate of reaction increases and the end point can be indicated by potassium sulfocyanate.
+ 2Fef++
C6H806
C6H606
+ 2Fe++ + 2 H r
The reproducibility of estimation was first investigated, then compared xvith other gravimetric and volumetric determinations of ferric ions. Experiments have been carried out for establishing whether the end point indicated by potassium sulfocyanate changes with temperature. According to the authors' experience, the solution containing ferric ions in slightly acidic medium consumed the same amount of ascorbic acid a t 60" C. as in the cold. The same results were obtained if the titration was performed in the reverse direction-i.e., ascorbic acid was determined with ferric ions. Accordingly, the error of the indicator is negligible when 0.1 h' solutions are titrated in the presence of potassium sulfocyanate. The dehydroascorbic acid may also react with ferric ions under certain circumstances. This process naturally would interfere with the estimation. To answer this question, a solution was
Table 11. Comparison of SIethods
Gravimetric Estimation as Fez02 5.933 5.975 5.978 5.954
A v . 5 960
( M g . of Fe per mi.) Volumetrical Estimation According t o Iodometri-With ascorTitanomet- Zimmermann- cally accordbic acid trically ( 9 ) Reinhardt ing t o Mohr 5.969 5.964 5.964 6.005 5.958 5.958 5.952 5.994 5.964 5.964 5.975 5.977 5.969 5.969 5.986 5.972
5.965
5,964
5.969
5.992 ~~
Table 111. Accuracy of Method (29.00 ml. of 0.1 N FeCls solutlon titrated with 0 1 S ascorbic acid solution) Ascorbic Acid Solution Consumed, 111. 29 00 29 00 29 02 29,Ol 29 00 29.00 29 00 29,Ol 28 98 28.98 28 99 28.98 29 03 28.99 29 03 29.01 29 03 29.02 29 02 29.00 A v . , 29.00 Mean error, f0.017 Mean error of average, 'Z, 0.004
ANALYTICAL CHEMISTRY
420
The nearly neutral solution containing 25 to 250 my. of ferric ion should be acidified with 5 ml. of 2 Nhydrochloric acid in a 250ml. titration flask, then heated cautiously SO that the temperature does not exceed 60’ C. Then 1 ml. of 0.5 N potassium sulfocyanate is added to the warm solution, which resulta in a red color. The titration is continued and in the vicinity of the end point the volumetric solution is added drop by drop until discoloration takes place within a few seconds. More than 5 minutes should not elapse between heating and the end of the titration. One ml. of 0.1 N ascorbic acid solution is equivalent to 3.584 mg. of iron, 7.184 mg. of FeO, 7.984 mg. of Fe2O3,or 16.221 mg. of FeCls. The mean error of the average values is &0,004% in the presence of 25 to 250 mg. of iron.
Table IV. Ore Tested Magnetite Pyrite Bauxite
Determination of Iron
Ascorbic Acid Method, % 79.87 79.92 79.87 42.06 42.10 42.06 21.78 21.71 21.73
Zimmermann-Remhardt Method, cZ 80.60
42 51 21.75
INFLUENCE OF FOREIGN IONS
Strong oxidizing materials like iodine, iodate, bromate, permanganate, chromate, nitrite, hydrogen peroxide, vanadium (V), cerium(IV), etc., react with ascorbic acid and make the estimation inaccurate. Colorless and difficultly reducible materials do not disturb the titration. Nitric acid in a third molar proportion, and nitrate ions even in multiple molar proportions do not interfere. The reaction is retarded by sulfate and phosphate ions, but even in tenfold molar proportion has no significant influence on the end point. Small amounts of fluoride ions do not interfere with the estimation and only when present in a molar ratio of 1 to 1 did the color of the ferric sulfocyanate fade to such an extent that its change could hardly be observed. In the presence of colored ions potassium sulfocyanate indicator could not be used. In this case the titration may be carried out electrometrically . DEAD- STOP END-POIhT TITRATION
As potassium sulfocyanate can be used only in colorless or slightly colored solutions to indicate the end point, observation of the end point was tried with Foulk and Bawden’s ( 1 ) “deadstop” method. The potential of an accumulator was decreased to 0.14 volt through a potentiometer. Two platinum electrodes 2 cm. long and 0.4 mm. in diameter were immersed in the solution to be analyzed and the circuit was closed through the electrodes. The intensity of the current was measured by a microammeter. Then 0.1 12; ascorbic acid solution was gradually added to the solution, which was heated to about 60” C. in the vicinity of the end point as in titrations made with potassium sulfocyanate. The intensity of the current is shown in Figure 2 as a function of the ascorbic acid used. The intensity increased suddenly when the solution was heated, owing to the decrease of internal resistance, then a sharp diminishing could be observed a t the equivalence point, after which only the intensity of the residual current could be marked. Accordingly, the end point is indicated by a sharp decrease in intensity. Ions of nickel, cobalt, and chromium do not interfere. The “depolarimetric” estimation was carried out also by the method of Guzman and Rancaiio ( 2 ) , when the primary potential was provided by a platinum-platinum-rhodium thermocouple. ESTIMATION WlTH 0.01 N ASCORBIC ACID
The estimation of ferric ions can also be performed with 0.01 N ascorbic acid. The volumetric solution was freshly prepared on every occasion by dilution of 0.1 A‘ ascorbic acid. On the basis of experimental data it can be stated that in the case of 3 mg. of iron the accuracy of the method is satisfactory and is equal to 0.5%. The concentration range 0.1 to 0.5 N for hydrochloric acid and 1 ml. of 0.5 N potassium sulfocyanate per 100 ml. was found the most suitable. TECHNICAL POSSlBILITIES OF IRON DETERMIN4TION WITH ASCORBIC ACID
There is a great advantage for technical analyses in the fact that the determination can be made in a hydrochloric acid medium. Ammonium salts, sulfate and nitric ions, and even small amounts of free nitric acid do not interfere with the end point.
An iron salt solution therefore which has undergone the most varied separating operations may be used for the assay without difficulty. The authors have tested among others the iron content of iron ores, sulfide ores, and bauxite, with the aid of this method (Table IV). In the case of many technical control tests it has been very difficult to effect the iron determinations in the presence of large amounts of phosphate or fluoride ion. With the present method even small quantities of iron are easily determinable in the presence of large amounts of phosphate and small amounts of fluoride ions. SUMMARY
Almost all osidimetric and reductometric methods used previously for determining ferric ions may be replaced by the new ascorbic acid procedure. Over oxidinietric methods it has the great advantage that iron need not be reduced before determination. Ferrous ions can be oxidized by boiling with a few drops of hydrogen peroxide. If nitric acid is present in the solution, it suffices to neutralize it. The method is preferable to the titanometric method, as the volumetric solution is easily prepared, it is much more easily checked, and the influence of atmospheric osygen is negligible. Compared to the chromometric titration, the preparation of the ascorbic acid reagent does not demand cumbersome procedures, no complicated mechanisms are necessary for its storage, and the estimation in most cases may be carried out without electrometric indication of the end point. In the presence of small amounts of nitric acid and fluoride ions and even large amounts of nitrate and phosphate ions, no interference could be noted. At present extensive studies are being made of the indirect application of the ascorbic acid determination of ferric ions, as well as of the use of redox indicators. Methods have been developed for the ascorbic acid determination of chlorate, bromate, iodate, and chromate as well as vanadic and ceric ions. Preliminary studies are under way on reductonietric determination of other oxidizing materials as well as of organic nitro and nitroso radicals by this new procedure. A colorinietric method has recently been developed, which is based on the strong change of color caused by ascorbic acid. The reactions between ascorbic acid and molybdic acid and tungstic acid, respectively, have been applied for establishing traces of metals catalytically. The results will be published shortly. The name “ascorbinometry” is proposed for this method. ACKNOWLEDGMENT
The authors are indebted to Pva B h y a i , Rlrs. L. Buz&s, Zsuzsanna Faber, and Gyorgy RAdy for their valuable assistance. LITERATURE CITED
(1) Foulk and Bawden, J . Ani. Chem. Soc.. 48, 2035 (1926). (2) Gurman and Rancaiio, Anal. SOC. espaA. fis. y qulm., 32, 590, 899 (1934); Z. anal. Chem., 99,284 (1934). (3) Knecht, E., Ber., 36, 1551 (1903). (4) Nascimento, Ruben do Rev. SOC. b r a d puim.. 16, 165 (1947). (5) Ptitsyn and Koelov, Zhur. A n d . Khim., 4, 35 (1939). (6) Sivadjian, J., “La Chimie des T-itamines,” Paris, Gauthier-Villars, 1949.
(7) Thornton and Ladusk. 1x1). E ~ G&EM., . RECEIVED December
1 , 1950.
h a L .
ED..4,240 (1 982).
