Table II.
Observed Values of the Constant C
Notation A AM
Material A1203
90% Also3 A M C 82% A1203
+
+ 10% h50Oa + 15% MoOa
3% coo A M N 82% A1203 j-15% ?vloOa
+ 3% N10
S SA1 SA2
CY0
Meterz/Cc.
SiOz SiO2-AI2O3(13% AlzOs)
650 460 490
values were compared with "B.E.T. surface areas" (1). The average ratios between surface area and air adsorption obtained with various materials are given in Table 11. After being calibrated, the pycnometer can be used to estimate surface areas. The formula for this is A = CVa/W
490
where
700 790
A
SA3 SiO2-A12O3(25% &Oa) 750 a Average ratio of specific area to specific air adsorption.
From these and similar observations with samples of different materials, it was concluded that no changes occurred in the adsorptive capacity of the studied materials as long as moisture was excluded from them. The minute changes observed while the samples were kept in sealed containers may be attributed in part to moisture leaking into the containers, and in part to exposure to regular room air while carrying out observations on weight and air adsorption. CORRELATION OF AIR ADSORPTION WITH SURFACE AREA
Physical adsorption in the Henry's law range is proportional to both surface area and the Henry's law constant which is a measure of attraction of the surface for air. As most inorganic oxides have rather similar surfaces, it appeared likely that the Henry's law constant would be about the same for all and that, therefore, the surface area would be proportional to the air adsorption. To test this possibility, measured air adsorption
=
C(l/D
-
T'ai,/W)
surface area per gram of sample estimated from air adsorption V , = change in air adsorption when pressure is increased from 1 t o 2 atm. at ambient temperature (23" =t4" C.) W = sample weight in grams D = skeletal density (gram per cc.) determined using helium VBir = pyncometer reading Kith air (in cc.) C = a constant characteristic of the substrate =
V ois the difference between pycnometer readings (helium reading minus air reading). From the reading with helium, the skeletal density, D, can be calculated, and when this is known, the second form of the formula can be used. Since the proportionality constant, C, varies from one material to another, the method should be calibrated for each material individually. However, if not too high an accuracy is required, one proportionality constant may be used for a number of closely related materials. Thus, for the three molybdena-containing materials the maximum error when using C = 460 would be less than 7%. When C and D are known, routine measurements on the same substance can be carried out quickly. Less than 5 minutes are needed to obtain the weight and the pycnometer reading with air, and from these the air adsorption "surface area" per gram can be calculated.
For samples with air adsorption in excess of skeletal volume, pycnometer readings are negative and cannot be obtained by the normal procedure. Handwheel B (see operating instructions) can be backed off from the stop position 1, 2, or 3 turns, as required, to bring the pointer on the scale. With our unit, one turn is equivalent to 2.9 cc. Figure 3 shows the correlation of surface areas, calculated from air adsorption by the use of the constants in Table 11, with B.E.T. surface areas (1). REPRODUCIBILITY
Many check measurements were made a t constant temperature. Standard deviation in pycnometer reading averaged 0.025 cc. In the case of 200 sq. meters per gram samples with an alumina base, this produces a standard deviation in surface area of about 0.5%; for 50 sq. meters per gram, about 2%, etc. The lack of high precision a t low surface areas, which in lesser degree is shared by low temperature nitrogen methods, is not normally important for work on catalysts. LITERATURE CITED
(1) Innes, W. B., Ax.4~.C H m . 23, 759 (1951). ( 2 ) Kalberer, W., Schuster, C., 2. physik. Chem. A141,274 (1929). ( 3 ) Krieger, K. A., J . Am. Chem. SOC. 63,2712 (1941). (4) Lambert, B., Heaven, H. S., Proc. Roy. SOC.(London)A153, 584 (1936). (5) Lambert, B., Peel, D. H. P., Ibid., A144, 205 (1934). ( 6 ) Pvlagnus, A., Grahling, K., 2. physik. Chem. A145,27 (1929). ( 7 ) Milligan, L. H., J . Phys. Chem. 26, 247 (1922).
(8) Tuul, J., DeBaun, R. M., ANAL. CHEM.3 4 , 8 1 4 (1962).
RECEIVEDfor review January 19, 1962. Accepted April 9, 1982.
FIuorosuIfonic Acid as Titrant in Acetic Acid RAM CHAND PAUL, SHAM KUMAR VASISHT, K. C. MALHOTRA, and SARVINDER SINGH PAHlL Deparfmenf of Chernisfry, Panjab University, Chandigarh-3, lndia Fluorosulfonic acid forms a solid compound HSOIF. CHKOOH with acetic acid. This compound has a high conductivity (2.4 X 10-2 ohm-' cm.-' at 60" C.). A comparison of the equivalent conductivities of various acids in acetic acid and comparative potentiometric titration studies of perchloric acid and fluorosulfonic acid in acetic acid indicate that in this medium fluorosulfonic acid is a slightly stronger acid than perchloric acid. Conductometric, potentiometric, and visual ti-
820
ANALYTICAL CHEMISTRY
trations using malachite green and crystal violet have established the suitability of fluorosulfonic acid as an acidic titrant in acetic acid.
A
is widely used as the medium for nonaqueous titrations. The acid-base reactions occurring in the solvent are well established ( 2 , 8, 9). The solvent system is represented as CETIC ACID
2CHaCOOH
CHaCOOH2+
+ CHaCOO-
In this medium, only strong protonic acids can act as acids. Organic bases and acetates behave like strong bases; even weakly basic compounds have their basicity enhanced (9). Perchloric acid has been considered the strongest acid in acetic acid. When perchloric acid is used in this medium, even weakly basic substances which cannot be estimated by direct titration in water, such as amines (6), sulfonamides ('i'),and alkaloids (IO,l'i'),have been successfully titrated. Mixtures of
E 0 2
4
Z
Figure 1 , acid
8
6
X
12
10
d
Equivalent conductivity of strong acids in acetic
primary, secondary, and tfrtiary amines have also been analyzed (3, 16, 18. 20). Sulfuric acid is a moderately strong acid in acetic acid. If a hydroxyl group were replaced by a fluorine atom, the acid strength of sulfuric acid would be expected to increase because of the strong electronegative character of fluorine. Woolf (21) studied the conductivity of fluorosulfonic acid in water and concluded that it was a strong acid. However, a comparison of the relative strengths of acids is not feasible in water because of the leveling-solvent properties of this medium. Therefore, to test the suitability and acid strength of fluorosulfonic acid on a comparative basis, solutions of this acid were examined conductometrically and potentiometrically in acetic acid.
On cautious addition of fluorosulfonic acid (15 ml.) to acetic acid (15 ml.) with constant stirring and cooling, a white solid separated out. The reaction mixture was cooled and filtered in a dry atmosphere. The compound was washed once with acetic acid (3 ml.) and then with dry carbon tetrachloride. It was placed in a vacuuni desiccator to remove the volatile carbon tetrachloride. The compound had a melting point of 52.5" C. The equivalent weight was determined by a sodium hydroxide titration using phenolphthalein indicator in water and was found to be 79.8. The equivalent
720
weight required for HS0,F. CH, COOH is 80.0. The composition of the compound has also been confirmed by plotting a composition us. conductivity curve which shows a break a t the fluorosulfonic acid/acetic acid molar ratio of 1 to 1 a t 60" C. Apparatus. Because of the low conductivity of very dilute solutions of fluorosulfonic acid in acetic acid, a conductivity cell having a cell constant of 0.032 cm.-l was used for the equivalent conductivity studies with fluorosulfonic acid. The conductivity of the molten compound HS03F.CH8COOH was determined in a cell having a cell constant 0.46 em.-'. A glass conductivity cell, equipped with a microburet, replaceable platinum electrodes, and calcium chloride drying tube was used for the conductometric titrations. A magnetic stirrer was introduced to facilitate mixing. The resistance of the solutions was measured a t 25' C. by dipping the cell in the thermostat maintained a t 25" f 0.05' C. for equivalent conductivity studies and by passing the water a t 28" 0.05" C. through the jacket of the titration cell for conductometric titrations. The specific conductivity of the molten compound HS03F. CHICOOH was determined a t 60" f 0.1" C. A precision measuring bridge type, W.B.R. No. 108, with a logarithmic indicator amplifier type TAV, IKC, KO. 034 (Wissenschaftlish Technische Werkstatten, W'ielheim/Oby/Germany), was employed. I n the potentiometric titrations, a fiber type saturated calomel electrode
*
1
EXPERIMENTAL
Preparation of Reagents. Acetic acid (B.D.H.), containing at the most O.lyo water, was twice frozen. Each time, the liquid acetic acid (about %Yo by volume) was discarded. The remaining material was then fractionally distilled at 116.5'117" C./740 mm. T h e absence of water was confirmed by a melting point determination. The purified acetic acid melted a t 16.5' C. and had a specific conductivity of less than 0.5 X 10-7 ohm-' cm.-l and the water content was less than 0.02% (Karl Fischer). Fluorosulfonic acid ( I d ) was prepared by heating fuming sulfuric acid and dry KHFz and collecting the distillate between 160" C. and 250" C. Fluorosulfonic acid was redistilled through a short column in dry nitrogen; the fraction distilling a t 161-2" C./748 mm. was collected (purity by equivalent weight determination, 99.9%). The bases were purified by a previously suggested method ( I S ) . All solutions were prepared in a dry box. Transfer of the solutions was carried out using dry nitrogen under pressure.
H $0, F H c1
360 0
0.~4
o l e 1:o
1.2
1.6 1.8 0
Molar ratio acidlbase Figure 2. I.
II.
0.4
0 . 8 1.0 1.2
o4
1.6 1 . 8
Molar ratio acid/base
Results of potentiometric titrations
Potassium acetate with HClOd and HSOaF Quinoline with HCIOI and HSOaF
VOL. 34, NO. 7, JUNE 1962
821
700
t"" 4 Figure 3. Potentiometric titrations of diethylaniline with HSO$ and HCIOd
-0-0-
-0-0-
was used as the reference electrode, and a quinhydrone electrode was employed as the indicating electrode (16). The potential was measured with a portable (Pye) potentiometer (Cat. No. 7569 P). The null point was detected with an external ampere-scale galvanometer. Potentiometric titrations were also carried out with a fiber type saturated calomel electrode and a chloranil electrode. To compare the strengths of the two titrants, perchloric and fluorosulfonic acids solutions of the same strength were used with the same solutions of the bases-Le., potassium acetate, quinoline, and diethylaniline. To study the effect of the presence of small amounts of water in acetic acid on fluorosulfonic acid, a known quantity of water was added to a known volume of acetic acid (standard water solution). Different quantities of standard water solution were added t o fluorosulfonic acid solution. Quantities of acetic acid were added also to keep the final strength of the fluorosulfonic acid solution the same. For potentiometric titrations the same amount of &-picoline was always used. Titrations were carried out with a fiber type saturated calomel electrode and a quinhydrone electrode. I n the visual titrations, the basic solution was placed in the titration flask, and a drop of indicator solution in acetic acid was added. The reaction mixture was stirred with a magnetic stirrer during the titration. A calcium chloride drying tube was also employed.
44
-
40
-
36 -
YO r-( x 0
32 -
28-
$
d u
4
3 d
9u
24-
20-
44
0
0.2
822
0
ANALYTICAL CHEMISTRY
0.4
0.6
0.8
1.0
1.2
1.4
M o l a r ratio acid/b,,
RESULTS
Fluorosulfonic acid forms a white crystalline compound, m.p. 52.5' C., having the formula HS03F. CH3COOH. I n the molten state at 60" C., this compound is strongly conducting (spe-
HS03F HCIOd
cific conductivity, 2.4 X 10-2 ohm-* cm.-I a t 60" C.), which indicates that the cornpound may be ionic in nature and could be represented as S03FCHJ2OOH2'. Figure 1 compares the equivalent conductivity of fluorosulfonic acid with the equivalent conductivities of other strong acids (4, 9). From the position of the various curves, i t is evident that fluorosulfonic acid may be slightly stronger than perchloric acid. Similar results with fluorosulfonic acid and perchloric acid have been reported recently for the sulfuric acid solvent system ( I ) . Plots of electromotive force us. acid/base molar ratio for potassium acetate and quinoline with perchloric acid and fluorosulfonic acid are shown in Figures 2 and 3. Since the same amount of base was used for both acids tested, the results definitely support the findings based on the equivalent conductivities of fluorosulfonic and perchloric acids, namely, that fluorosulfonic acid is the stronger acid. Figure 4 shows the conductometric
Figure 4.
Conductometric titrations of HS03F in acetic acid 1. II. 111. IV.
Qulnoline Pyridine Dimethylaniline a-Picoline
1.6
1.8
340
300
I 0
I
,
0.4
I
0 . e 1.0 1.2
,
I
I
o
1.6
‘4
-
640
-
Pyridine a-Picoline Dimethylaniline Quinoline
600-
Y)
d2
m
5
b
560-
-
sao -
t-t-
k,
%
0.0
w
400-
water
0 . 3 %water
--Ce
440-
d
0.1 %water
-
480-
0.5
% water
1.0
Z water
2.0 % r a t e r
5.0
%writer
360I
320
o
l 0.4
0.e
1.2
1.6
,
1
1.6
ratio acid/bass
Potentiometric titrations of HS03F I. 11. 111. IV.
680
,
0 . 8 1.0 1.2
M&P
Molar ratio acidlbase Figure 5.
I
0.4
2.0
Eolar ratio aoid/ba$e Figure 6. Effect of water on potentiometric titrations with ocetous fluorosulfonic acid
titration curves for the titrations of solutions of pyridine, quinoline, LYpicoline, and dimethylaniline, with fluorosulfonic acid in acetic acid, When fluorosulfonic acid is added to a solution of the base in acetic acid, there is a rise in the conductivity of the solution until a maximum is reached a t the molar ratio of 1 to 1, when there is a break in the curve. Further addition of acid results in a gradual decrease in the conductivity of the solutions possibly due to the dilution effect. The course of these titrations can easily be explained on the basis of the behavior of acids and bases already established in the acetic acid solvent system. I n the conductometric titration curves, molar ratio of acid to base has been plotted against relative conductance, which is equal to the reciprocal of resistance. The cell constant has not been taken into account because of the difficulty in placing the electrodes in exactly the same position in the cell. However, their position was kept the same in the individual titrations. Figure 5 gives the potentiometric titration curves for pyridine,. quinoline, a-picoline, and dimethylanilme, with fluorosulfonic acid using calomel and quinhydrone as reference and indicator electrodes, respectively. As in the normal course of potentiometric titrations in water, there is a small change in electromotive force in the initial stages of the titration. The change in electromotive force increases as the acid/base molar ratio approaches 1. At this ratio, the curve is steepest. On further addition of the acid solution, the change in electromotive force decreases, and the titration curves gradually become almost parallel to the base line. Similar results are obtained by plotting AEIAV us. the acid/base molar ratio. Figure 6 shows the effect of water in acetic acid on titrations with acetous fluorosulfonic acid. The addition of water to acetic acid in which fluorosulfonic acid solution is prepared has practically no effect until a concentration of 1% of water is reached. Beyond this concentration, the total electromotive force rise becomes smaller. Apart from this, the solution begins to attack glass. Crystal violet and malachite green have already been used for titrations in acetic acid (5,11,14,19). Crystalviolet which is blue in basic solutions changes from a bluish green to a green color on the addition of fluorosulfonic acid. With an excess of acid, the color changes to yellow-green and finally to yellow. Results of titrations using crystal violet as the indicator are given in Table I. The color change for malachite green a t the end point with fluorosulfonic acid is green to yellow. Table I1 gives the results for titrations of fluorosulfonic acid using malachite green as indicator. VOL. 34, NO. 7, JUNE 1962
823
LITERATURE CITED
Table 1.
Titrations of Bases Using Acetous Fluorosulfonic Acid with Crystal Violet
Base Pyridine Quinoline a-Picoline Dimethylaniline Diethylaniline Potassium acetate
Table II.
Gram Equivalent Titrated x lo4 Experi- Theomental retical 8.21 8.23 5.91 5.95 10.62 10.54 6.81 6.82 26.71 27.00 12.77 12.74
Error,
yo
-0.24 $0.67 -0.75 -0.14 -1.07 -0.24
Color of Indicator a t End Point Green Green Green Green Green Green
Titrations of Bases Using Acetous Fluorosulfonic Acid with Malachite Green
Base Pyridine Quinoline a-Picoline Dimethylaniline Diethylaniline Potassium acetate ~~~~
Titrant Concentration, Gram Equivalent Weight of per Liter Base, G r a m 0.18288 0,06492 0.13756 0.07679 0.31128 0.09594 0.17068 0.08250 0.21676 0.39800 0.21508 0.12498
Titrant Concentration, Gram Equivalent Weight of per Liter Base, Grams 0.06492 0.18288 0.13756 0.07674 0.31248 0.09594 0.17068 0.08250 0.21676 0.39800 0.12498 0.21508
For both indicators, the color changes at the end point were very sharp, and no precipitate formed in either case. Fluorosulfonic acid has been successfully used to estimate alkaloids and amino acids. The details regarding the estimation of acetates, bases, alkaloids, and amino acids are being published separately. I n addition to being a stronger acid,
Gram Equivalent Titrated X lo4 Experi- The+ mental retical 8.22 8.21 5.97 5.95 10.46 10.54 ~. 6.81 6.82 26.71 26.80 12.77 12.74
Error,
76
-0.12 -0.33
+, o- . x
. -
-0.14 -0.34 -0.24
Color of Indicator at End Point Yellow Yellow
ypiiow . ~ ~ ~
Yellow Yellow Yellow
fluorosulfonic acid has the advantage that it is available in a pure form and its solutions can be completely anhydrous. With perchloric acid, however, this is not true. Perchloric acid solutions have to be prepared indirectly, or water has to be removed with acetic anhydride and other such reagents. Pure perchloric acid is unstable, and only a 70% solution is available.
(1) Barr, J., Gillespie, R. J., Robinson, E. A., Can. J . Chem. 39, 1266 (1961). (2) Conant, J. B., Hall, N. F., J . Am. Chem. SOC.49, 3047 (1927); 49, 3062 (1927). (3) Dimorth, K., hleyer-Brunot, H. G., Biochem. 2.323, 338 (1952). (4) Emeleus, H. J., Haszeldine, R. N., Paul. R. C.. J . Chem. SOC.1955. 563. (5) F1 (E
pharm. franc.
(7) Herd, R. I Sa'. Ed. 31, 9 (8) Kolthoff, I J . Am. C (9) Koltho 56, 1007 (1Y34). (10) Levi, L., Oestreicher, P . M ., Framilo, C. G.. Bull. Narcotics. U . N. Devt. Social 'Agairs 5, 15 (1953). (11) hlarkunas, P. C., Riddick, J. A., ANAL.CHEM.23, 337 (1951). (12) Meyer, J., Schramm, G., Z. anorg. Chem. 24, 206 (1932). (13) Paul, R. C., Sandhu, S. S., Singh, J. S., Singh, G. S., J . Indian Chem. SOC. 35, 877 (1958). .(14) Seaman, W., Allen, E., Ibid., 23, 592 (1951). (15) Shkodin, A. W., Izmailov, N. A., J . Gen. Chem. U.S.S.R. 20,39 (1950). (16) Siggia, S., Hanna, J. G., Kervenski. I. R., ANAL.CHEX 22, 1295 (1950). (17) Splenger, C. H., Kaelin, H. A., Pharm. Acta Helv. 18, 542 (1943). (18) Tomicek, O., Collection Czechostou. Chem. Commun. 13, 116 (1948). (19) Tuthill, S. M., Xolling, 0. W.,Roberts, K. H., ANAL. CHERT.32, 1679 f 1960).
(20) Wagner, C. D., Brown, R. H., Peters, E. D., J . -4m. Chem. SOC.69, 2609 (1947). (21) Woolf, A. A., J . Chem. SOC.1954, 2840. RECEIVEDfor review August 23, 1961. Accepted February 27, 1962.
infrared Determination of Aldehydes An Improved Group Type Analysis E. L. SAIER, L. R. COUSINS, and M. R. BASILA Gulf Research & Development Co., Pittsburgh 30, Pa.
b
An improved infrared procedure is described for the group type analyses of aldehydes. These analyses are based upon the integrated absorptivity of the aldehydic C-H stretching vibration. The application to both aromatic and aliphatic aldehydes i s discussed.
I
an effort to extend and improve previous work (7) pertaining to a group type analysis for oxygenated materials, a study was made of the aldehydic C-H stretching vibration in N
824
ANALYTICAL CHEMISTRY
a series of aromatic and aliphatic aldehydes over the region from 2600 em.-' to 2900 cm.-l Generally, two bands characteristic of aldehydes are observed in this region. The assignment of these bands has been discussed by several authors (1, 2, 4-6). The most recent work supports the hypothesis that the characteristic doublet occurring in this region results from a Fermi resonance interaction between the aldehydic C-H stretching fundamental and the first overtone of the aldehydic C-H bending vibration (8). This paper describes group type anal-
yses for both aromatic and aliphatic aldehydes. The application to aromatic aldehydes is new and the application to aliphatic aldehydes results in considerable improvement over our original method ( 7 ) . This improvement is due largely to the use of integrated absorptivities rather than peak values. EXPERIMENTAL
All data were obtained using a Perkin-Elmer Model 21 spectrophotometer equipped with a LiF prism. The wavelength scale was expanded to 50 cm. per micron and the scanning speed
