with vniious concentrations of maleic and fumaric acids gave positive rcactions to the reagent indicating that not all of the abietic-type diene acids had reactd In preliminary studies, the nir,thod has also been used to follow the diiappearance of abietic-type diene acids during the preparation of rosin polymers. The method may be applied also to determine the abietic-type diene acid contcnt of gum, wood. and disproportionated rosins, as wcll as other matc.rials which might contain the abietic-type diene rosin acids. ‘The 1 1 of~ the Lieberniaiin reaction in testq for cholesterol and other sterols prompted tests on such compounds n ith the modified rcagent. Cholcderol and miwd phytosterols gave only very faint pink tests under the conditions outlined earlier in this paper. TT hich R ould eliminate them as srrious interfwing con5tituents. I n tests with oleic a i d Iinoloic acids. only linoleic acid
gives any appreciable color at low concentrations. It produces a reddish color with an absorption maximum which was shifted considerably from that of the rosin acids. However, the presence of large quantities of fatty acids with respect to the rosin acids, as in the case of tall oil fatty acids. does produce a sizable blank reading and therefore decreases the sensitivity of the determination. It is felt that the modified Liebermann method presented offers the advantage over many other tests, in that the specificity and sensitivity of the test for abietic-type rosin acids is increased considerably. I n addition, the stability of the color affords more dependence upon the actual color produced rather than on the fleeting glimpse of a color which may or may not hare been present. This remores some of the guesswork from the interpretation of the qualitative results and provides a
more reliable quantitative means of estimating the abietic-type rosin acids. LITERATURE CiTED
(1) B,urchard, H., Dissertation, Rostock
University, Germany (1889); Chem. Zentr. 25 (1890). ( 2 ) Conner, A. Z., “Chemical Analyses of Resin-Based Coating Materials,” C. P. A. Kappelmier, ed., p. 31, Interscience, New York, 1959. (3) LaLande, W. A., Jr., J . Am. Chem. Soc. 5 5 , 1536 (1933). (4) Liebermann, C., Ber. 17, 1884 (1884). (5) Storch, L., Ber. osterr. Ger. Z. Ford., Chem.-tech. Ind. 9, 93 (1887); Chem. Zentr. 1419 (1887). (6) Su-ann, M. H., ANAL.CHEM.2 3 , 885 (1951). RECEIVEDfor review August 28, 1961. Accepted Xovember 30,1961. The Naval Stores Laboratory is a laboratory of the Southern Utilization Research and Development Division, .4gricultural Research Service, U. S. Department of Agriculture. R. L. Stephens was a fellow of the Pulp Chemicals Association while this work was in progress.
Spectrophotometric Determination of Furfural in the Presence of Sulfur Dioxide JOHN F. HARRIS and LAWRENCE L. ZOCH Forest Products laboratory, Forest Service, U . S. Department of Agriculture, Madison, Wis.
b The absorbance of aqueous furfural solutions is depressed by the presence of sulfur dioxide, because of the formation of the aldehyde-bisulfite complex. The equilibrium constant for the complexing reaction was measured. The strong absorptivity of furfural, together with its property of complexing with bisulfite ion, suggests its use as an indicator for measuring bisulfite ion concentra tion.
A
basis for making a quantitative determination of furfural is provided by the fact that aqueous furfural solutions strongly absorb ultraviolet light. The spectrophotometric methoci has been thoroughly studied by Root (9,S ) , who investigated the effect of storage conditions upon the stability of dilute furfural solutions and developed a n analytical procedure for determining micro quantitirs (less than 1.0 fig.). The method was extremely accurate n lien eniployctl for the determination of pure aqueous solutions of furfural. W i r n crude solutions are to be analyzed, the furfural may be readily distilled from acidic or neutral medium without loss. If no other volatile materials are present, Root’s method is applirable,
but in the event that other substances are present in the distillate, it is necessary to determine their effect on the absorption spectra. The presence of sulfur dioxide in aqueous solutions of furfural would be expected to have a pronounced effect on the absorptive 07
I
I
270
280
FURFURAL No MOLAL R A V ~
li EXCELLEST
-0 -5 I
$
04
8 8
3
03
02
01
1 0250
I
260
-
I 290
300
WAYELENGTH i -mp
Figure 1. Ultraviolet spectral curves for furfural-sodium bisulfite solutions Constant furfural mg. per liter
level
approximately
3.5
capacity of the solution because of the addition complex formed (4, 5 ) . EXPERIMENTAL
Reagents. The furfural n a s t h e middle-third fraction obtained from t h e distillation of material obtained by double distillation of comrnercially available material. It had been sealed in a glass ampoule and stored in a freezer for a 4-year period. When originally stored it was water-white; when used it was clear, with a very pale yellow color. Sodium bisulfite, analytical reagent grade (Baker). Apparatus. A Beckman Model DK2 recording spectrophotometer was used for curve plotting (Figure I ) and a Beckman Model DU was used t o measure t h e absorbance of solutions at 276 mp. Solutions. Fifteen solutions containing the same total concentration of furfural b u t 1Vith varying amounts of sodium bisulfite were prepared by mixing stock solutions of furfural and sodium bisulfite. The molal ratio of furfural t o bisulfite ion in these solutions ranged from 1 : O t o 1: 100. RESULTS
The curves of Figure 1 illustrate the effect of increasing the bisulfite ion conVOL. 34, NO. 2, FEBRUARY 1962
* 201
centration in the solution. All the solutions have the same total furfural concentration, and the presence of sodium bisulfite results in a marked reduction in absorption. There is no shift in the position of the maximum absorption peak and no change in the general shape of the spectral curve. The absorbances at 276 mp (A,, for furfural) for the 15 solutions are listed in Table I. All measurements were made at the ambient temperature of 22O & 1 ° C . CORRELATION METHOD
Equations 1 to 3 describe equilibrium in the system. C-C
+
0
I1
b-4
C
\ - /
i
+
SOs-' H20 OHHAOII
u+
C-C /I
C-C
ll
%
OH I
OH
HZSOs
H+
+ HSOI-
(3)
Equations 1 and 2 involve formation and ionization of the complex. I n analogous analyses of similar systems, these equations have erroneously been combined into a single equation (6, 7). Although of little consequence in this case, such combination leads to misinterpretation of results at higher concentrations. Because of the extreme
Table 1.
Sample No, 1 2
3 4 5 6 7
8
9
10 11
12 13 14 15 Av.
202
Total NaHS08 Concn. 598.2 299.1 119.6 59.82 29.91 23.93 17.95 11.96 5.982 2.991 2.393 1.795 1.196 0.598 0
dilution (10-6 t o gram mole per liter) of all samples, complete ionization of the complex and all salts was assumed. The sum of un-ionized sulfurous acid and sulfite ion can be shown to be trivial at these concentration levels. The effect of p H on the system is twofold, both direct and indirect. The direct effect is shown in Equation 2, which shows that a n increase in H + will tend to shift the ionized complex t o the un-ionized state. The indirect effect is shown in Equation 1, which shows that a n increase in H + will tend to withdraw the bisulfite ion from the system, resulting in a decrease in the amount complexed. Therefore, changes in p H will have a large effect on the system. However, if the concentration level of all components is low, as is the case in this work, so that it can be assumed that all species are present in ionic form, the variations in p H are small and may be assumed inconsequential. The justification for the assumption is apparent from the fine fit of the data. It was also assumed that the ultraviolet absorption of the complex is negligible. The high absorptivity of furfural is undoubtedly due t o conjugation of the aldehyde group with the double bonds in the heterocyclic ring. I n the sodium bisulfite-furfural complex, the aldehyde group is not present, and the complex would be espected to exhibit a spectrum similar to that of furan. Furan exhibits a maximum absorption a t 250 mp with a n absorptivity of 1 liter per gram mole per em. (I), compared to 276 mp and 14,850 liters per gram mole per em. for furfural ( 2 ) . Thus, the conjugated double bonds in the furan ring contribute little t o ultraviolet absorption. The supposition that the absorption of the complex is negligible is borne out by the curves of Figure 1, iyhich show no shift
Absorbance of Furfural Sodium Bisulfite Solutions
(Concentration, gram mole/liter X 105) Absorbance a t Uncombined Combined 276 M p Furfural Furfural 0.754 5.280 0.112 4,707 0.197 1.327 3.610 0.360 2,424 2.539 3.495 0.519 1.583 0.661 4.451 4.694 1,340 0.697 5.037 0,997 0.748 0.687 5.347 0.794 0.337 5.697 0.846 0.873 5.879 0.155 5.906 0,128 0.877 5.953 0,081 0.884 5.993 0.041 0.890 6.020 0.014 0.894 6.034 0 0.896
ANALYTICAL CHEMISTRY
HSOaConcn. 582.9 294.4 116.0 56.28 27.33 22.59 16.95 11.27 5.645 2.836 2.265 1.714 1.155 0.584 0
Calcd. Furfural Concn. 5.997 5.987 5.779 5.841 5.902 5.959 6.055 6.066 6.081 6.078 6.066 6.075 6.076 6.062 6.034 6.0039
in the wavelength of maximum absorption. Thus: -4bsorbance at 276 nip
=
A
=
EICI (4)
where €1
molal absorptivity of furfural a t 276
=
nv
C1 = molal concentration of uncombined
furfural in solution
Defining the equilibrium constant for the complex as:
v-here [C?] = molal concentration of complex ion in solution [HSOI-] = molal concentration of uncombined bisulfite in solution
I t is easily shown that:
where C F = iiiolal concentration of total furfural in solution
Thus, for a group of solutions with constant total furfural concentration but with varying amounts of bisulfite ion, the reciprocal of the absorbance should plot as a straight-line function of the bisulfite ion concentration. The value of K is the ratio of the slope to the intercept. The experimental data listed in Table I were fitted by least mean squares to the form of Equation 6 and Equation 7 was obtained. _ -- 1343 [HSOa-]
A
+ 1.1253
17)
From Equation 7 , the value of K is 1193 liters per gram mole. The average value of the total concentration of furfural in the solutions may be calculated from the intercept value, using a value of 14,850 liters per gram mole per em. for the absorptivity of furfural. This is 5.984 X gram mole per liter, which compares to 6.034 X 10-5 gram mole per liter obtained from the absorbance of the solution containing no sodium bisulfite. Recognizing that the total bisulfite concentration of the solution, C801, is the sum of the complexed bisulfite and free bisulfite, (HSOI-), Equation 6 may be rearranged to yield :
CF
=
[€I
+AK A l
[k'(CSh
+ ): + 11 (8)
Thus the total concentration of furfural may be determined when the total concentration of bisulfite in solution and the absorbance are known.
Values for total furfural concentration of the test solutions, calculated by Equation 8, are listed in the last column of Table I. These have a relative standard deviation of k1.587, and a standard deviation, when compared to the average value calculated from the intercept, of =k1.62y0. DISCUSSION
The procedure described may be used for the determination of furfural in the presence of bisulfite ion. Such solutions are frequently encountered when working with sulfite pulping liquors, but the method should be of much greater usefulness, as it can be used to obtain quantitative information on many of
the aldehyde-bisulfite equilibria. Such systems are encountered frequently and, in general, little information is available. Equilibrium measurements may be made by adding small, known quantities of furfural and determining the absorbance of the resulting solutions. From these measurements, the free bisulfite ion in the solutions may be obtained. The furfural is thus employed as a n indicator of bisulfite ion concentration. When used in this manner, the method should be useful for solving many problems. LITERATURE CITED
Miller, F. A,, in Gilman, H., “Organic Chemietry,” Vol. 111, 1st ed., p. 167,
(1)
Wiley, New York, !953:
( 2 ) Root, D. F., “Kinetics of the AcidCatalyzed Conversion of Xylose to
Furfural,” unpublished Ph.I>. thesis, Chemical Engineering Department, University of Wisconsin, 1956. (3) Root, D. F., Saeman, J. F., Harris, J. F., Keill, W. K., Forest Products J . 9, 1.58 f l%W\. , - - - - I
(4)-Sheppard, W. A., Bourns, A. N., Can. J . Chem. 32, 4 (1954). (5) Shriner, R. L., Land, A. H., J . Org. Chem. 6, 888 (1941). (6) Stewart, T. D., Donnally, L. J., J . Am. Chem. SOC.54, 2333, 3555 (1932). f7) Tomoda. Y.. J . SOC. Chem. Ind. ’ (Japan)30, 747 (1927). RECEIVEDfor review August 10, 1961. Accepted November 6, 1961. Study supported in part by Sonoco Products CO., Hartsville, N. C.
Solvent Extraction with Quaternary Ammonium Halides ARTHUR M. WILSON, LILLIAN CHURCHILL, KENNETH KILUK, and PAUL HOVSEPIAN Chemistry Department, Wayne State University, Detroit 2, Mich.
b The solvent extraction of metal ions b y quaternary ammonium halides dissolved in 1,2-dichIoroethane i s described. The extraction i s dependent upon the ion association of the quaternary ammonium ion with labile chloro anionic complexes of Fe(lll), Co(ll), Zn(ll), and TI(III) or with the oxyanions o f Hf(lV), Ta(V), and Mo(VI) in concentrated hydrochloric acid or lithium chloride. The effects o f metal ion concentration, contacting time, and type of quaternary ammonium halide are investigated with the cobalt ion. Extractions o f for cobalt are possible only if the distribution ratio of the quaternary ammonium halide itself i s very large. The value of the distribution ratio of the cobalt ion i s dependent upon the quaternary ammonium ions’ size, shape, and type o f organic groups. An analogy between solvent extraction with quaternary ammonium halides and anion exchange of metal ions in concentrated hydrochloric acid is demonstrated.
>soy0
I
RECENT YEARS, many solvent extraction systems have been described which utilized high molecular \\-eight amines. General review of this field have been written by Coleman et al. (4), Moore (IS), and Morrison and Freiser (14). I n general, metal anionic complexes and simple and complex oxyanionic metals and nonmetals have been extracted from acidic aqueous media. The following paragraph, quoted from Moore’s review (13), is an excellent N
the same general form (17) that Kraus et al. (9) observed with Dowex 1; thus, many investigators have drawn a formal analogy and called liquid “Smith and Page (19) first reported extractions with high molecular weight that the acid-binding properties of amines “liquid anion exchangers.” high molecular weight amines depend If the extraction mechanism is ion on the fact that acid salts of these association to produce neutral species, bases are, in general, essentially init is reasonable that high molecular soluble in water but readily soluble in weight quaternary ammonium halides organic solvents, such as chloroform, benzene, or kerosine. The extraction should be more efficient as liquid anion reactions are of the following ionexchangers. I n fact, they should be association type: operative even us. basic aqueous media, 1. The organic solvent containing as Katekaru and Freiser (8) have the amine can extract an aqueous acid demonstrated with the extraction of to form an amine salt in the organic the anionic tris(8-hydroxyquinolinato) phase : complex of calcium(I1) a t p H 12.2(R3S)o Ha+ A , - S ( R I N H - ~ - ] ~ 12.9 into methyl isobutyl ketone. Other authors have used quaternary where RBN = 3 high molecular weight ammonium compounds to charge-neuamine tralize anionic metal complexes and A = an anion of either a simple acid or a complex hence enhance the percentage of metal metal acid, like FeC14extracted. Clifford et al. (3) have o = organic phase utilized Arquad 2C-75 (a 75% solua = aqueousphase tion in isopropyl alcohol of dialkylI n alkaline solution, the extraction is dimethylammonium chloride where each reversed. of the two alkyl groups has 8 to 18 2 . An amine sult in the organic carbon atoms, averaging about 16 phase can undergo anion exchange with carbon atoms) to charge-neutralize the an ion in the aqupous phase: anionic 8-quinolinol complex of uranyl (RaNH+A-)o Ba-CRJJH+B-).+ A,ion to extract it more efficiently into methyl isobutyl ketone. Maeck et al. The order of preference in the amine organic solution is similar to that in (11) have utilized tetrapropylammoanion exchange resins : nium nitrate to charge-neutralize the anionic trinitrate complex of uranyl C10,- > ?io$-> C1- > HSO4- > F-(4).” ion for extraction into methyl isobutyl ketone. Since less polar solvents than This exchange of the chlorometal commethyl isobutyl ketone can extract plex anions for the simple anion apuranyl complexes-viz., uranyl nitrate parently occurs a t the interface and into pentaether (16) and uranyl quinoproduces distribution ratio us. hydrolinolate into chloroform (It?), it is chloric acid molarity curves which have
explanation of the mechanism of distributions with high molecular weight amines.
+
+
+
VOL. 34, NO. 2, FEBRUARY 1962
203
