described here also has obvious manipulative advantage, especially in remote handling, over currently-used models that require graphite powder insulation for the production of adequate fusion temperatures when using 4 to 5 kW induction heating generators. Other improvements to the apparatus previously described (8) include: a diaphragm pressure regulator used instead of the liquid manostat; two vacuum thermocouple gauges used to determine the degree of nitrogen removal from the charcoal traps; and two gas flowmeters employed to detect gas linkage in the portion of the analytical train confined to the hot cell. Procedure for Changing Crucibles. The procedure for remotely changing crucibles is not difficult. After removing the 45° viewing mirror and loosening the clamp on the 35/25 O-ring joint, the upper part of the joint is rotated out of the way, and, with forceps that engage the inner lip of the crucible, the latter is lifted out and placed in an active waste container. A new crucible is inserted into the chamber. The chamber is closed and the train is again readied for use by adding 8 grams of platinum to the crucible, followed by a sufficient number of heating cycles to reduce the blanks of oxygen, nitrogen, and hydrogen to low constant values.
RESULTS AND DISCUSSION
Recovery data on samples of known carbon content are shown in Table I. These data are presented not to prove the validity of this well-established method but to show that remote handling does not necessarily prevent the acquisition of useful results in the microgram range. Recovery results obtained on oxygen, nitrogen, and hydrogen standards were essentially the same as reported previously (8). Some advantages of the modified equipment are not necessarily confined to remote handling. The carbon furnace
Table I.
Recovery Data
on
Samples of Known Carbon Content
Sample
Added
LECO Standard
90
C, Mg Measured 89 88 88 83
% 99 98 98 92
Av 97 Std dev ±
BaCOs
190 210 215 201 188
3
98.9 99.0 99.5 200 99.5 186 98.9 Av 99.2 Std dev ±0.3 188 208 214
chamber permits preheating a combustion crucible to a constant low blank (in the microgram range) before the sample is added. Sample addition is made without exposure of the ignition vessel (crucible) to the atmosphere. The same degassed crucible can be used for several successive samples. Likewise, the modification of the fusion furnace chamber is generally advantageous in that it eliminates the problems of crucible alignment, of sample losses by missing the crucible, and of carbon-SiO¡> interaction.
Received for review September 27, 1967. Accepted January 15, 1968. This paper is based on work performed under the auspices of the U.S. Atomic Energy Commission.
Potentiometric Titration of Perchlorate with Tetraphenylarsonium Chloride and a Perchlorate Ion Specific Electrode R. J. Baczuk and R. J. DuBois Hercules Inc., Bacchus Works, Magna, Utah
organometallic salts have been used to the perchlorate anion. Both electrometric and gravimetric methods have been employed. This paper reports a potentiometric precipitation titration of perchlorate with tetraphenylarsonium chloride. The titration is followed with a perchlorate ion specific electrode and a double-junction calomel electrode. Best results are obtained with solutions between pH 4 and 7. The method is simple, rapid, and free from common interferences. However, extremely large amounts of some simple anions distort curve shapes and require adjustment in the titrant standardization. Accuracy and precision of the method for assaying simple perchlorate salts are equivalent to or better than available methods. Based on replicate analyses of three perchlorate samples, overall 95% confidence limits were ±0.16%. In addition, the perchlorate electrode was found to respond linearly to permanganate, dichromate, and periodate ions over an appreciable concentration range. A number of measure
amperometric titration with tetraphenylphosphonium chloride. Morris (2) described an amperometric determination of perchlorate with tetraphenylstibonium sulfate. As early as 1939 Willard and Smith (3) proposed the use of tetraphenylarsonium chloride for perchlorate analysis, based on a potentiometric back-titration with triodide ion. Recently, Baczuk and Bolleter (4) developed a conductometric precipitation titration of perchlorate with tetraphenylarsonium chloride, suitable for assay purposes and free from most common anionic interferences. However, there has been no published work employing any of these salts as titrants in conjunction with an ion specific electrode for the determination of perchlorate. The availability of an ion specific electrode for perchlorate (Orion Research, Inc.) suggested the possibility of a direct (1) H. Nezu, Bunseki Kagaku, 10, 561 (1961); Chem. Abstr., 56,
26/(1962).
The application of organometallic salts such as tetraphenylarsonium chloride to the measurement of perchlorate has been reported by many workers. Nezu (7) employed an
(2) M. D. Morris, Anal. Chem., 37, 977 (1965). (3) . H. Willard and G. M. Smith, Ind. Eng. Chem., Anal. Ed., 11, 186 (1939). (4) R. J. Baczuk and W. T. Bolleter, Anal. Chem., 39, 93 (1967). VOL. 40, NO. 4, APRIL 1968
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685
MOLES / LITER
Figure 1. Response of the perchlorate electrode to ammonium perchlorate solutions A. B.
Log activity vs. potential Log concentration vs. potential
potentiometric measurement of this anion. The electrode was reported to be suitable for single measurement determinations of perchlorate at concentrations of 10-1M to 10_5M, between the pH of 4 and 10-11. While the electrode was sufficiently selective (free from interferences of other common anions), preliminary work showed that a single or direct potentiometric measurement of perchlorate was not sufficient for assay purposes. Potential responses had to be determined with an accuracy better than a hundred microvolts. Because of the problems of residual junction potentials, asymmetry effects, and long-term drift, the accurate measurement of these small voltages is presently not feasible (5). On the other hand, the use of this electrode as an indicator in a potentiometric precipitation titration of perchlorate with tetraphenylarsonium chloride offered the possibility of a rapid, accurate, and specific means of measurement. The application of the latter approach to the determination of perchlorate is reported in this paper. EXPERIMENTAL Reagents. Tetraphenylarsonium chloride was purchased from J. T. Baker Chemical Co. and tetraphenylarsonium chloride hydrochloride from Eastman Organic Chemical Co. or Hach Chemical Co. Titrant solutions, 0.05M in tetraphenylarsonium chloride, were prepared by dissolving either form of the reagent in an appropriate amount of distilled water and adjusting the solution to about a pH of Oc5 with dilute sodium hydroxide or hydrochloric acid. casionally, a lot of tetraphenylarsonium chloride hydrochloride was obtained which was somewhat impure. This was evidenced by the presence of an appreciable amount of The addition of Celite to the a water-insoluble material. aqueous preparation of the reagent, followed by filtration through Whatman No. 42 filter paper, eliminated the problem. Fisher certified reagent ammonium perchlorate was used for standardization of the titrant. Apparatus. A Metrohm Potentiograph, Type E-336, was used for both automatic and manual titrations. An Orion (5) G. A. Rechnitz, Rec. Chem. Progr., 26, 4 (1965).
686
·
ANALYTICAL CHEMISTRY
perchlorate ion specific electrode, Model 92-81, was used as the indicator electrode. A double-junction calomel electrode, Beckman 40452, was used as the reference. The outer compartment was filled with an ammonium nitrate solution. This double-junction electrode prevented potassium ions from forming a potassium perchlorate precipitate in the capillary of the commonly-used, single-junction calomel electrode. Procedure. The perchlorate sample to be titrated was dissolved in about 75 ml of distilled water. A sample of about 1.5 to 2.0 millimoles of perchlorate usually was taken for analysis. With 0.05M titrant solution, this sample size provided a 30- to 35-ml titration. The sample was titrated potentiometrically with the 0.05M tetraphenylarsonium chloride using the perchlorate ion specific electrode and double-junction calomel electrode. Manual titrations were performed in the usual fashion. The potential of the stirred solution was allowed to come to a constant value before a reading was taken. Generally, stable potential readings were obtained within 15 seconds of reagent addition even near and after the end point. For automatic titrations, delivery rates no higher than 2 ml per minute were used near the end point. Stirring was performed at such a rate that no vortex or accompanying bubbles were formed in the solution. If formed, bubbles frequently became attached to the membrane surface of the perchlorate electrode and caused unstable potential measurements. Standardization was performed with Fisher certified reagent ammonium perchlorate. End points were located by rectangulation. In the presence of large amounts of other anions, standardization should be performed on solutions containing approximately equivalent amounts of these anions. RESULTS AND DISCUSSION
The Orion perchlorate ion specific electrode operates in a somewhat similar to the conventional glass pH electrode. Both electrodes develop a potential across a thin layer of conductive material. However, with the perchlorate electrode a water immiscible liquid ion exchanger, held in place by an inert semipermeable polymer disk, is substituted for glass. According to the manufacturer (Orion Research, Inc., Instruction Manual, Perchlorate Ion Electrode, Model 92-81), this electrode exhibits a high specificity for the perchlorate ion and responds in Nernstian fashion to this ion between concentrations of 10-1M and 10~bM. The only serious interference reported is for the hydroxide ion which has a selectivity constant equivalent to that of perchlorate. The electrode is recommended for use from a pH of 4 to 10-11. Electrode Response. The potential-concentration response of the electrode for perchlorate was determined in this laboratory. The potentials of solutions from 10~lM to 10-6M in ammonium perchlorate were measured. The temperature was maintained at 25° C. Curves, representing changes in potential vs. both the log perchlorate concentration and log perchlorate activity, are presented in Figure 1. No effort was made to maintain the same ionic strength in each test solution. The activities of the solutions were calculated, using the reported activity coefficient of perchlorate for the particular ionic strength of each test solution (Orion Research Instruction Manual). A linear potential change was observed for perchlorate at concentrations between 10-W and 10-W. In this range, a 10-fold change in concentration resulted in a potential change of 57 mV and a 10-fold change in activity showed a corresponding change in potential of 59 mV. The response time of the perchlorate electrode in stirred solutions was observed to increase as the perchlorate concentration decreased. A constant potential reading was obmanner
Table I.
Recovery of Ammonium Perchlorate at Various pH Values
Initial pH of solution
NH4CIO4 added, mg
0.8 2.0 4.0 5.0 7.0 9.0
213.7 209.6 203.2 202.9 207.1 208.7
sample
MILLILITERS
OF
0.05M TETRAPHENYLARSONIUM CHLORIDE
Figure 2. Theoretical and experimental titration curves 200 mg ammonium perchlorate in 65 ml water A. B. C.
Actual titration curve Predicted Nernstian response curve Predicted response curve based potential data from Figure 1
on
for
log activity perchlorate/
tained almost instantaneously for 10_ lM solutions, but at the 10~5M level about 15 seconds were necessary for stabilization. Titration of Perchlorate with Tetraphenylarsonium Salts. The precipitation titration of perchlorate with tetraphenylarsonium chloride was easily followed using the perchlorate ion electrode. Preliminary titrations of ammonium perchlorate solutions were performed using both tetraphenylarsonium chloride and the less costly form of the reagent, tetraphenylarsonium chloride hydrochloride. Both salts were found to give equivalent titrimetric results. A typical titration of 200 mg ammonium perchlorate in 75 ml of water with 0.05M tetraphenylarsonium chloride is shown in Figure 2, curve A. End point potential breaks of about 150 to 200 mV were observed. The titration curve obviously was not symmetric. Figure 2, curve B, was then prepared to illustrate the shape expected if the electrode had behaved in Nernstian fashion. The concentrations of perchlorate in solution beyond the end point were obtained from the solubility product constant (Ksp) of tetraphenylarsonium perchlorate, determined by an ultraviolet method previously reported (4). The solubility of the tetraphenylarsonium perchlorate was determined in such a way as to approximate the Ksv expected at the end point of the titration—i.e., in the presence of an amount of ammonium chloride equal to that at the end point. The resultant Ksv was 8.2 X 10~s. Corrections to activities were then made. The respective emf values were obtained from the Nernst equation and plotted to complete curve B. To determine if the asymmetry of curve A, Figure 2, was due solely to the electrode response, Figure 2, curve C, was prepared. To obtain this curve, perchlorate activities were selected at several points in the titration and plotted against the corresponding emf’s taken from Figure 1. This plot, curve C, illustrated the best titration curve shape obtainable with the electrode. As can be seen, curves A and C were identical. Thus, the shape of the titration curve was mainly caused by the deviation from Nernstian behavior of the electrode at low perchlorate concentrations rather than poor control of titration variables. Effect of pH on Perchlorate Recovery. The effects of pH on recovery were determined by titrating known amounts of
NH4CIO4 recovered, mg 243.0 213.3 203.6 202.8 206.8 211.1
Error, %
+ 13.7 + 1.8
+0.2
0.0 -0.1 + 1.2
perchlorate at several values of pH. The results of this study are shown in Table I. Acceptable recovery of perchlorate was found in the pH range between 4 and 7. As mentioned, the hydroxyl ion shows a selectivity equal to perchlorate for this electrode. In the region of the end point, if the hydroxyl concentration approaches and/or exceeds the concentration of perchlorate, a significant distortion of the titration curve and an error could be introduced in the titration. This behavior was demonstrated by the slightly high recovery at pH 9. However, the recovery at pH 9 was greater than expected from just distortion of the titration curve by the hydroxyl ion effects. The tetraphenylarsonium ion was suspected of forming an associated ion pair with the hydroxide at high pH’s which would also increase the volume of titrant required. Titrations performed in solutions of pH 2 or less exhibited useful breaks but recoveries of perchlorate were somewhat erratic. Although recoveries were high, they actually varied with different lots of reagent in these highly acidic solutions. No experimentally-verified causes can be offered to explain this behavior. Effects of Stirring and Titration Rate. As described above, the response time of the perchlorate electrode was found to increase as the perchlorate concentration decreased. However, the response speed of the electrode was found to be sufficiently rapid for both manual and automatic titrations, provided the automatic titrations did not exceed a titrant delivery rate of 2 ml per minute in the end point region. Streaming or mixing potentials were reported for some anion specific electrodes (6). In our work, the stirring rate of the titration solution was not found to be critical as long as there was some mixing. Stir rates were varied over a wide range in manual titrations with no observable effect on the potentials. However, complete cessation of stirring caused a gradual decrease in the potential, probably caused by a diffusion of perchlorate from the membrane into the sample solution. On the other hand, very high stir rates caused a noise in the potential reading. This was attributed to vortexing and the formation of bubbles about the tip of the perchlorate electrode. Interferences. Previous conductance work (4) demonstrated that the determination of perchlorate with tetraphenylinterarsonium chloride was relatively free from common ferences. These earlier studies were performed on 150-mg quantities of ammonium perchlorate together with an equivalent amount of the salt in question dissolved in 600 ml of water. No interference was observed with chloride, chlorate, bromide, brómate, fluoride, sulfate, nitrate, and chromate. (6) G. A. Rechnitz, 20th Annual Summer Symposium, Analytical Chemistry, Claremont, Calif., June 21, 1967. VOL. 40, NO. 4, APRIL 1968
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687
Table II.
Recovery of Ammonium Perchlorate in One-to-One Salt Mixtures
Salt added, mg
Salt
205.4 208.5 205.2 208.1 208.5 206.3 208.5 205.5
KC1
KClOg
KBr
KBrOs KNOs K2Cr04 NaF K2S04
Table III.
E NH4CIO4 added, mg 208.8 204.7 203.2
208.0 207.8 204.3 210.4 208.2
NH4CIO4 recovered, mg 211.3 245.1 203.6 210.2 217.1 207.1 210.2
Error, %
+ 1.2 + 19.7 +0.2 +1.1 +4.7 + 1.4 -0.1 -0.2
207.7
Comparative Reliability Data
_%
Potentiometric titration with tetraphenylarsonium chloride
Sample
99.17 99.40 99.26 99.20 99.33
American potash NH4CIO4, lot 90-2
NH4CIO4
Titanium(III) reduction (7) 99.15 99.30 99.23
X
=
99.27% 99.79 99.90 99.88 99.70 99.76
99.23% 99.60 99.64 99.68
X
=
99.81% 99.50
99.64%
American potash NH4CIO4 lot 9015-00006
Mallinckrodt
99.51 99.50 99.42
Reagent grade KCIO4
lot
7053
X
=
99.48%
100.40 100.38 100.78
100.52%
=
The perchlorate electrode was reported to respond to some anions to varying degrees and large quantities of foreign salts influenced the activity of the sample solution (Orion Research Instruction Manual). Therefore, an interference study was conducted to demonstrate the effects of large concentrations of several anions on the recovery of ammonium perchlorate. One-to-one mixtures containing 200 mg of ammonium perchlorate and the various salts were placed in 75 ml of water and titrated. These concentrations, of course, would never be encountered in practice but demonstrated the electrode behavior under extreme circumstances. Prior to testing, the solutions were checked to assure the proper pH range—i.e., pH 4 to 7. The anions tested were: chloride, chlorate, bromide, brómate, fluoride, sulfate, nitrate, and chromate. The results of the study are shown in Table II. As can be seen, most of the salts caused a slight increase in the perchlorate recovery but this was due to (with one exception) distortions in the titration curve rather than precipitation. This disNitrate caused the most serious distortion error. tortion was evidenced by a decrease in both magnitude and slope of the end point break. These distortions were considered to be an inherent shortcoming of the electrode. The potential response relationship ·
ANALYTICAL CHEMISTRY
=
e
—
—
F
In (Acio4~ + KiCi)
The coefficient Ki is the selectivity constant for an interfering ion at a concentration Ci and Acio4~ is the activity of the perchlorate ion in the test solution. Although the value of KiCi essentially does not change in a titration where precipitation of the interfering ion does not occur, the net effect of KiCi is a suppression in the range of the potential response for perchlorate. Inasmuch as the titration curves are asymmetric in the region of the end point, smaller potential responses have a significant effect on end point location. Also of importance is the effect of Ci on the total ionic strength and, hence, activity of perchlorate in the sample solution. The activity of perchlorate is decreased by Ci, causing higher potentials throughout the titration. The net effect again is a suppression of the potential break, caused by the limiting potential response of the electrode. In actual assay work, Ci of course would be very small and would have no effect on the titration curve shape or recovery of perchlorate. The large positive error in the recovery of potassium chlorate was mainly caused by a nonquantitative precipitation of tetraphenylarsonium chlorate. This precipitation occurred in high concentrations of chlorate only. Chlorate did not interfere with the titration when present in less than 0.02M concentrations. [This concentration was never reached in earlier conductance work (4), hence, the interference was not noted.] In addition to these common anionic impurities listed in Table II, dichromate (at pH l), periodate, iodide, and permanganate were also checked for they had been found to precipitate with tetraphenylarsonium chloride in previous conductance work. Permanganate and periodate were precipitated stoichiometrically in a mole-for-mole reaction. The dichromate ion in highly acidic solutions precipitated quantitatively also but in the ratio of two moles of titrant to one mole of dichromate. The iodide interference was not quan=
Potentiometric method Overall std dev 0.073%
688
of the electrode to perchlorate and to interfering ions may be represented in the following manner:
titative. Standardization of Titrant. The titrant was standardized against high purity ammonium perchlorate. Considerable was also obtained by titration against potassium success permanganate and potassium dichromate but these salts did not provide any particular advantages other than the possibility of a referral to a primary standard. As described, the interference studies showed changes in curve shape when large amounts of certain anions were present. This presented some difficulty in end point location. Therefore, special care should be taken when perchlorate is to be measured in the presence of high concentrations of these anions. The standardization should be performed on perchlorate solutions containing the appropriate salts of the anions added in roughly equivalent amounts to that present in the test solutions. Aqueous solutions of tetraphenylarsonium chloride and tetraphenylarsonium chloride hydrochloride, adjusted to a pH of 4 to 7, gave equivalent results as titrants. The hydrochloride form, however, is appreciably cheaper and is therefore, preferred. Precision and Recovery. Reliability data are given in Table III for replicate analyses of three perchlorate samples by the potentiometric tetraphenylarsonium chloride titration. Also included are data obtained by the titanium chloride method (7) to serve for comparison. The two ammonium perchlo(7) E. A. Burns and R. F. Muraca, Anal. Chem., 32, 1316 (1960).
rate lots were purchased to a federal specification (Mil-A192A) and were certified to have a minimum ammonium perchlorate content of 99%. The potassium perchlorate was reagent grade material. As can be seen, a favorable comparison was obtained between methods for analyses of both AP lots. The slight discrepancy in the case of potassium perchlorate cannot be explained although the titanium chloride method gave results over 100%. An overall standard deviation of 0.073% was obtained for the potentiometric method. This compared favorably with the conductance
determination of perchlorate for which a standard deviation of 0.074% was obtained (4). Applications of the Potentiometric Perchlorate Titration. The primary use of this potentiometric precipitation titration is for assay of various perchlorate salts. In addition, testing of a wide range of other perchlorate-containing samples is feasible. For example, this titration was successfully used for rapidly measuring ammonium perchlorate in certain solid propellants. The perchlorate in both composite and composite-modified double-base propellant formulations was determined following a water extraction of ground samples. The use of aqueous solutions for analysis was necessary since most organic solvents attack the perchlorate ion electrode. Surprisingly, the perchlorate electrode was found to respond to permanganate, periodate, and dichromate. This was verified by preparing individual solutions of each of these ions at various concentration levels and measuring the potentials with the perchlorate electrode. Plots of emf vs. log concentration are shown in Figure 3. No corrections were made to activities. As can be seen, the plots are quite similar to the perchlorate plot and are linear over about the same concentration range. The dichromate measurements were made at pH 1. In addition, solutions of these ions were quantitatively measured by a potentiometric precipitation titration with
M0LES/LITER
Figure 3. Perchlorate electrode response curves for Cr20 72~, C104", ICV, Mn04A. B. C.
D.
Dichromate 1.0M in acid Perchlorate Periodate Permanganate
tetraphenylarsonium chloride and the perchlorate electrode. Thus, the electrode might be useful as an indicator for potentiometric titrations involving these ions. Received for review December 4, 1967. 12,1968.
Accepted January
Potentiometric Determination of Acid Groups in Acrylic Polymers and Fibers J. Ray Kirby and A. J. Baldwin Chemstrand Research Center, Inc., Research Triangle Park, Post Office Box 731, Durham, N. C.
A procedure is described whereby acidic groups in acrylics at the microequivalents level may be measured by nonaqueous potentiometry. A purified ethylene
carbonate-propylene carbonate mixture is used as solvent. Tetramethylammonium hydroxide in this solvent provides a superior titrant. Use of platinum electrodes permits titration of acid groups in polymers in the presence of sodium and potassium salts, thus avoiding the problem of alkali metal poisoning encountered with glass-calomel electrodes. Titration of acid groups present as salts is performed after ion exchange treatment with appropriate resins. The technique is sufficiently sensitive to permit differentiation of acid groups of different strengths in the same polymer sample. Typical titration curves are given. The need to measure functional groups in polymers on a quantitative basis has assumed a role of increasing importance in recent years as the science of polymers has progressed. Quantitative determination of such functional groups is particularly difficult because of the extremely low level at which
they exist in the polymer. Polymers in a molecular weight range suitable for preparing synthetic fibers quite often have less than 100 ^eq of functional groups per gram of polymer. However, such analyses may play an integral part in understanding the mechanisms of polymer initiation and termination, the role of the catalyst, the degree of incorporation of functional groups and, in general, one’s ability to characterize the polymer. A broad knowledge of the types and amounts of functional groups is important to the dye chemist in order to develop suitable dyestuffs and dyeing techniques for yarns and fabrics. In the specific case of acrylics, the type and quantity of acidic groups present are of primary interest. A quantitative measure of the acid group content of acrylics is particularly useful in predicting the basic dye acceptance—i.e., the degree of acceptance of basic dyes (cationic dyes) by acidic sites on the polymer or fiber. In this paper the term acrylic polymer is to be considered synonymous with acrylic fiber unless otherwise specified. VOL. 40, NO. 4, APRIL 1968
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