Results obtained on 13 different organic liquids are compared in Table 111. Functional groups such as alcohols, amines, esters, ketones, and aldehydes are represented. I n general, the agreement bctween the near infrared and Karl Fischer results is good, especially considering that only single determinations were made. I n the case of morpholine, the low results were caused by insufficient drying; the dried material contained 0.043% water by the Karl Fischer method. However, morpholine can be more efficiently dried by using a larger amount of Molecular Sieve 4%.The relatively poor results with n-butylamine are probably caused by the broadness of the water band and interference from the primary amine peak, making accurate measurement of the absorbance difficult. Excluding the few extreme errors, the accuracy is about *0.02~o, absolute, in the range 0.02 to 1% water. The absorptivity of the water band ranged from 0.52 ml. per gram-em. in triethylamine to 1.81 ml. gram-cm. in dimethylphthalate. No special precautions were taken to exclude air from the samples and standards, except that all were kept in closed containers whenever possible and the usual precautions for handling hygroscopic materials were observed. 'The various data indicate that the near infrared determination of water with batch drying by Molecular Sieve 4.i is applicable t o many organic systems containing alcohols, esters, ketones, aldehydes, and amines. The sensitivity and accuracy of the method are a t least partly dependent upon the hydrogen-bonding property of the solvent, with the poorer hydrogen-
Table 111.
Comparison of Near Infrared and Karl Fischer Methods
HzO,
Compound Acetone Acetonitrile n-Butylamine Diethylamine Dimethylphthalate Ethyl acetate Ethylene glycol 2-Propanol Methyl Cellosolve 2-Methylpentaldehyde Morpholine
wt. %
by KF 0.50 0.56 0.49 0.17 0.68 0.27 0.79 0.18 0.81 0.17 0.46 0.25 0.62 0.20 0.51 0.13 0.52 0.17 0.65 0.22 0.58 0.40
Phenyl Cellosolve
0.55 0.13 0.49
H20, wt. % by NIR 0.52 0.59 0.48 0.16 0.65 0.31 0.66 0.17 0.75 0.17 0.48 0.23 0.61 0.21 0.50 0.13 0.53 0.15 0.66 0.24 0.59 0.36 0.51 0.13 0.50
bonding solvents having better sensitivity and accuracy. The most critical aspect of the method is the drying step, and the results will automatically be low by the amount of water not removed from the dried sample. LITERATURE CITED
(1) Chapman, D., Nacey, J. F., Analyst
83,377 (1958). (2) Cordes, H. F., Tait, C. W., ANAL. CHEM.29,485 (1957).
Deviation of NIR from KF, % as Ht0 +o. 02 +0.03 -0.01 -0.01 -0.03 + O , 04 -0.13 -0.01 -0.06 0.00 f0.02 -0.02 -0.01 +0.01 -0.01 0.00 +o. 01 -0.02 0.01 +0.02 +0.01 -0.04 -0.04
Relative deviation, yo
+
0.00 +o. 01
+4
-6 -4
$15 - 16 -6 -7 '+4 -8 -2
?; '+2
- 12 $2
+9
+2 - 10
-7
'+2
(3) Streim, H. G., Boyce, E. A., Smith J. R., ANAL.CHEM.33,85 (1961). Division of Analytical Chemistry, 142nd Meeting, ACS, Atlantic City, N. J., September 1962. R. L. MEEKER F. E. CRITCHFIELD E. T. BISHOP Research and Development Dept. Union Carbide Chemicals Co. Division of Union Carbide Corp. South Charleston, W. Va.
indirect Procedure for the Determination of Tin(l by Potentiometric Titration SIR: The quantitative determination of tin in the bivalent state is a rather tedious analysis as the tin(I1) ion is very sensitive t o air oxidation. To titrate tin(I1) directly with an oxidizing agent such as iodine, potassium iodate, potassium permanganate, or potassium dichromate, it is necessary t o remove all dissolved oxygen from the titrant and to maintain an oxygen-free atmosphere above the solution during titration. The need for these precautions is particularly acute when either permanganate or dichromate solution is used because each has the ability to induce the air oxidation of tin(I1) ions ( I , 14). A review of all direct methods of analysis for tin(I1) is given by Kolthoff and Elving (11). The need for titrating in an oxygen-
free atmosphere can be avoided by use of a n indirect procedure. Indirect methods of analysis are often applied to strong reducing agents which are susceptible to air oxidation. By adding an excess of some readily reducible species, an equivalent amount of an intermediate is formed which can then be titrated in the presence of air. A typical reagent for this type of procedure is a solution of iron (111)chloride. Use of an indirect method for estimating tin(II), in which iron(I1) is produced as an intermediate, was first reported by Druce (3) ; however, his procedure, which involved titrating the iron(I1) intermediate with potassium dichromate and using diphenylamine as an indicator, was capable
of onlv ~ 2 %accuracv. Another procedire u t i l k i g the &me strategy was reported recently by Donaldson and Moser (g), who employ ceric sulfate and 1,lO-phenanthroline-ferrous sulfate as the titrant and indicator, respectively. An accuracy of approxtin(I1) content is imately *0.2% obtained by use of this procedure. The only other mention of indirect tin(I1) analyses, via the iron(I1) route, is found in the reference texts of Kolthoff and Sandell (la)and Laitinen (IS) ; however, no procedural details are given in either instance. This paper presents an alternate procedure which involves potentiometrically titrating the iron(I1) intermediate with a standard solution of potassium dichromate. VOL. 34, NO. 1 1 , OCTOBER 1962
151 1
EXPERIMENTAL
Reagents. Iron(II1) chloride solution, 3.0 mg. of F e per ml., was prepared by dissolving 14.6 grams of reagent grade iron(II1) chloride-6hydrate in 100 ml. of dilute hydrochloric acid (5395) and diluting to 1 liter with water. Standard potassium dichromate solution, 0.1000N, was prepared by dissolving 4.9030 grams of purified (twice recrystallized and dried a t 200" C.) reagent grade potassium dichromate in water and diluting to exactly 1 liter. Nitrogen gas was used for all deoxygenating procedures ; however, to remove any traces of oxygen that might be present, the gas was passed through a train of wash bottles containing vanadium(I1) ions and amalgamated zinc by the procedure of iMeites and Meites (15). Apparatus. A Fisher Titrimeter equipped with platinum and calomel electrodes was used for t h e potentiometric titrations. Procedure. The details of this procedure through the point of dissolving the sample are similar to those employed by Farnsworth and Pekola (6)for the iodimetric titration of tin(IIj .' A volume of 175 ml. of 6AVhydrochloric acid was added to a 250-ml. Erlenmeyer flask. Several glass beads were added and the flask was stoppered with a two-holed rubber stopper. A glass tube reaching almost to the bottom of the flask was inserted through one hole of the stopper. This tube sewed as an inlet for nitrogen gas and the second hole in the stopper served as an outlet for escaping gas. The flow of nitrogen gas was begun and the flask was placed on a hot plate. The solution was boiled gently for 10 minutes. The combination of boiling and purging with nitrogen gas was used to remove all dissolved oxygen from the solution. A 100- to 200-mg. sample of the tin(I1) compound was weighed into a microsize glass cup. Without interrupting the flow of nitrogen gas, the rubber stopper was lifted out of the flask just far enough to allow the sample cup to be dropped into the solution. The flask was immediately recapped and the solution swirled to help dissolve the sample. An excess of iron(II1) chloride solution was introduced into the flask through the second or outlet hole using a pipet. This solution was also deoxygenated beforehand by gentle boiling and purging with nitrogen gas. The amount of this reagent required to constitute a n excess was determined by the persistence of the yelloy color characteristic of the iron(II1) ion. The solution was then cooled to room temperature. The flow of nitrogen gas was maintained during the cooling period. Finally, the cooled solution was transferred quantitatively to a beaker and the iron(I1) which had been formed was titrated with a standard solution of potassium dichromate. The equivalence point was determined potentiometrically and the tin(I1) content of the original sample was then calculated. 15 1 2
ANALYTICAL CHEMISTRY
Table 1.
Sample
Sample Tin(ll) Analyses Theoretical Found
Wt., mg.
%
Wt., mg.
%
RESULTS
Examples of some typical results obtained by the use of this method are given in Table I. Excellent results were obtained on the samples of pure electrolytic tin. Deviations for the tin(I1) salts occurred probably because these compounds were not loo'% pure. (Another possible explanation for the low values, based on an effect caused by the fluoride ion, will be discussed later.) DISCUSSION
The first step in this procedure is represented by the equation: Sn+*+ 2Fe+S + S n f 4 2Fe+*
+
The quantitative nature of this reaction is readily verified by calculations based on the standard electrode potentials for the couples Sn(I1)-Sn(IV) and Fe(I1)-Fe(II1). The equilibrium constant is loz1 and a t the equivalence point, [Sn+4]/[Sn+2] = 1.3 X 10'; nevertheless, the reaction actually should not be written as taking place simply between the individual ions, as it was shown by Gorin (7) that no reaction occurs in perchloric acid medium. Thus, because the reduction proceeds only in the presence of chloride (or other halide) ions, the correct reaction is one involving chloride complexes of the ions. I n this reaction, not only the equilibrium is favorable, but also the rate. Studies by Duke and Peterson (4) have shown that the reaction is first order in both iron(II1) and tin(I1) and fourth order in free uncomplexed chloride ion concentration. As a result, Kolthoff and Belcher (10) suggest that when reducing iron(II1) with tin(I1) in 1.5N hydrochloric acid medium, a large excess of a chloride salt should be added. Quantitative reduction was obtained in this study without the addition of extra chloride ion, however, because of stronger acid conditions than those mentioned. There are several possible objections to the use of strong acid solutions for this reduction. First, according to Kolthoff and Belcher (9), a slightly acid solution (O.&l;Zr, of iron(I1) is stable and atmospheric oxidation is very slow; however, in neutral or strong acid media the oxidation rate is substantially increased as represented by 4Fe+* + O2 + 4H+ + 4Fe+3
+ 2H2O
The results of our investigations indicate, however, that this rate of oxidation is still sufficiently slow, even in strong acid solution, t o be considered negligible during the short time (10 to 15 minutes) necessary to complete the titration of the iron(I1). A second consideration regarding the use of strong acid solutions is the question of their suitability as a medium for the dichromate titration. Tsubaki (16) reported that 3.5N hydrochloric acid is the maximum concentration tolerable during the titration of iron(l1) with dichromate. The fact that 6,V acid was used successfully in our procedure may seem to be inconsistent with these findings; however, the original 6 N concentration of our acid has been reduced to the acceptable level as a result of the boiling and the addition of rinse water during the t'ransfer of vessels. Of course, if necessary, further dilution would be possible just prior to the titration. Unnecessary dilution should be avoided, however, because the rise in potential observed a t the equivalence point in the potentiometric titration decreases in magnitude as the acidity decreases. I t has been reported by Hostetter and Roberts (8) that small m o u n t f i of iron(I1) are produced by the boiling of solutions of iron(II1) chloride. This could lead to erroneously high tin(I1) assays as some iron(I1) would then be present in addition t o that formed by the reducing action of tin(I1); however, a study of the amount of reduction caused by the boiling of the iron(II1) solution during the deoxygenating step revealed that it is negligible for the procedure described. When employing this method for the determination of the tin(I1) content of samples containing either fluoride or phosphate, it is necessary to add an m o u n t of iron(II1) chloride considerably greater than the calculated theoretical amount before the yellow color characteristic of an excess will persist because both fluoride and phosphate ions form colorless complexes with iron(II1). These complexes evidently do not take part in the reaction with tin(II), and hence sufficient iron (111) must be added to satisfy both the complexation and reduction steps before a quantitative amount of iron(I1) will be formed; an excess of iron(II1) will then be visible in solution. The presence of these complexes also lorers the oxidation potential of the Fe(I1)Fe(II1) system and thereby affects the magnitude of the rise in potential observed a t the equivalence point in the titration with dichromate. An observation was also made that the results obtained in the presence of fluoride ion were always low. A possible explanation could be that fluoride ions induce the air oxidetion of iron(I1) ;
however, the results obtained from several inert-atmosphere titrations of solutions containing fluoride did not conclusively verify this explanation. More likely, this error in determining the tin(I1) content of fluoride-containing samples occurs because a portion of the tin(I1) is complexed by fluoride ions in the form of SnF3- ions. This complex ion would be more stable than the corresponding chloride complex, and i t is possible t h a t SnF3ions would markedly reduce the rate of the reaction with iron(II1). This view is substantiated by the findings of Duke and Pinkerton (6),who report that the order of effectiveness in promoting reaction between tin(I1) and iron(II1) is chloride < bromide < iodide. Although no data on fluoride liere given, it is obvious that the order observed is that of decreasing complexing affinity. Therefore, if sufficient time were not allowed for complete reduction, a low result could be obtained 11hen SnF3- ions are present. The accuracy of this procedure compares favorably with that of other methods for tin(I1) analysis. The number of steps involved is greater than for a direct procedure, but elimination of the necessity for maintaining an oxygen-free atmosphere during titration is an advantage which compensates in part for the extra steps. This method is also less empirical in nature than the conventional tin(I1) pro-
cedures in which the titer of the oxidant is applicable only when all titrations are performed in exactly the same manner and at approximately the same speed. These restrictions are necessary because errors due t o atmospheric oxidation are assumed to occur; however, by developing a pattern of titration, the errors during standardization against tin metal and those during an actual titration are then comparable and it can be assumed that these errors will cancel. I n the present method all errors due to atmospheric oxidation have been eliminated. Thus the empirical tin standard is not necessary. Use of this method is limited, of course, to samples containing no strong reducing agents other than the tin(I1) ion, or, if such additional reductants are present, then a prior separation of the tin(I1) is necessary before the procedure can be employed, another disadvantage of this method in comparison to iodimetry is the interference of ferrous iron. Oxidizing agents other than potassium dichromate could be used for the titration of the iron(I1) intermediate; however, this reagent was chosen because of its prolonged stability and because it serves as its own primary standard. LITERATURE CITED
(I) Bray, W. C., Ramsey, J. B., J . Am. Chem.’Soc. 55, 2279 (1933).
(2) Donaldson, J. D., Moser, W., J . Chem. SOC.84, 10 (1959). (3) Druce, J. G. F., Chem. A’ews 128,
273 (1924). (4) Duke, F. R., Peterson, N. C., Zowa St. coll. J . Sci. 32, 89 (1957). (5) Duke, F. R., Pinkerton, E. C., J . Am. Chem. SOC.73, 3045 (1951). (6) Farnsworth, M., Pekola, J., AI.. CHEW31, 410 (1959). ( 7 ) Gorin, M. H., J . Am. Chem. SOC. 58, 1787 (1936). (8) Hostetter, J. C., Roberts, H. S., Zbztl. 41, 1348 (1919). (9) Kolthoff, I. M., Belcher, R., “Srolumetric Analysis,” Vol. 111, p. 602, Interscience, New York, 1957. (IO) Ibid.,p. 622. (11) Kolthoff, I. M., Elving, P. J., “Treatise on Analytical Chemistry,” Part 11, Vol. 3, p. 328, Interscience, New York, 1961. (12) Kolthoff, I. M., Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” p. 568, Macmillan Co., New York, 1953. (13) ,L$tinen, H. A,, “Chemical Analysis, p. 373, McGraw-Hill Book Co., New York, 1960. (14) Lenssen, E., Lowenthal, J., J . Prakt. Chem. 76, 484 (1859). (15) Meites, L., Meites, T., ASAL. CHEV. 20, 984 (1948). (16) Tsubaki, I., Runseki Kagaku 3 , 253 (1954). RONALD W. COLLIXS~ \frILLIAM H. NEBERG.4LL Chemistry Department Indiana University Bloomington, Ind. Present address: Inorganic Research Department, Wyandotte Chemicals Corp., Wyandotte, Mich.
A Simplified Technique for Analysis of Binary Mixtures by Second-Order Differential Reaction Rates SIR: Recently Siggia and Hanna
reacted.
A plot of ln(b-z),(a-z)
(1, 4 ) presented an interesting graphical
us. t , therefore, yields a straight line
method for the analysis of binary mixtures by means of second-order differential reaction rates. They employed the expression, initially presented by Lee and Kolthoff ( 2 ):
after all al has reacted, and extrapolation to zero time (t = 0) yields the intercept In(b - & / ( a - a:) from which a: can be calculated. Also Equation 1 is valid only when b # a. Recently a graphical method for second-order differential kinetics which was restricted t o the case, 6 = a, has been presented (3). Single-point and double-point techniques for this case, which greatly reduce the amount of experimental data required for this analysis, were developed. Although these techniques are somelvhat less accurate (3) than graphical methods, they are less laborious and therefore more suited t o routine analysis of large numbers of similar samples. The purpose of this communication was to develop single- and double-point techniques for the method of Siggia and Hanna (1, 4) which is not restricted to
lid
=
1 (b - a)
[In(a( b -- x) X )
where b equals the initial concentration of reactant and kq is the rate constant of slow reacting species, a2. The quantity a is given by a = a? ai
+
where a: and a! are the initial concentrations of faster and slower reacting species, al and a2, respectively, and x is the total amount of a, and az reacted at any time, t. Equation 1 holds only for times after a, has “completely”
b = a, and is, therefore, somewhat more general. Single - Point Method. Solving Equation 1 for a: yields: a -b aP =
1-
(‘e.?) e(b-Q)kPt
(2)
(Gx) e(b-Q)kzl b - x
Equation 2 can serve as a means of performing the analysis by substituting in the value of x a t a single time t , during the reaction. This time is restricted in t h a t it must be any time after al has reacted to completion but prior to completion of the reaction of a2. To apply this method, the quantities kp,6, and a must be measured or known. [The total quantity of the two species being determined must be measured independently, as is the case in the Siggia and Hanna method (1, 4).] This technique is in a sense temperature dependent because the rate constant VOL. 34, NO. 1 1 , OCTOBER 1962
1513
