Anal. Chem. 1982, 5 4 , 2085-2089
by using an efficiency estimated by just a few samples to correct the radioactivity values in the homogeneously quenched parts of t.he set. LITERATURE CITED (1) Bell, C. G., Jr., Hayes, F. N., Eds. "Liquld Scintillation Countlng"; Pergamon Press: New York, 1958. (2) Birks, J. B. "The Theory and Practice of Eicintillation Counting"; Pergamon Press: Oxford, 1964. (3) Horrocks, D. L., Ed. "Organic Sclntillators"; Gordon and Breach: New York. - . 1988. (4) Bransome, E. D., Jlr., Ed. "The Current Status of Liquid Scintillation Countlng"; Grune and Stratton: New,York, 1970. (5) Horrocks, D. L., Pw~g,C.-T., Eds. Organic Scintillators and Liquid Scintillation Counting": Academic Press: New York, 1971. (6) . . Hendee, W. R. "Radloactive Isotopes in Biological Research"; Wlley: New York, 1973. (7) Horrocks, D. L. "Applications of Liquid Scintillation Counting"; Academic Press: New York, 1974. (8) Fox, B. W. I n "Laboratory Techniques in Biochemistry and Molecular Biology"; Work, T. S . , Work, E., Eds.; North-Holland: Amsterdam, 1976; Vol. 5, part 1, pp 1-333. (9) Noujaim, A. A., Ediss, C., Wiebe, L. I., Eds. "Llquid ScintlllationScience and Technology"; Academlc Press: New York, 1976. (10) Soini, E. Sci. Tools 11976, 25(3), 1-11. (1 1) Neary, M. P.; Budd, A. L. I n ref 4; pp 27:3-282. (12) Peng, C. T. I n "Advances in Tracer Methodology"; Rothchild, S.,Ed.; Plenum Press: New York, 1966; Vol. 3, pp 81-94. (13) Peng, C. T. I n ref 4; pp 283-292. (14) Barrows, G. H.; Samols, E.; Becker, B. J. Nucl. Med. 1976, 17, 1017-1 018. (15) Patterson, J. F.; Sauerbrunn, B. J. L.; Battist, L. Anal. Blochem. 1978, 86,707-715. (16) Horrocks, D. L. Int. J . Appl. Radiat. Isol. 1975, 2 6 , 243-256. (17) Neame, K. D. Anal. Blochem. 1976, 9 1 , 323-339.
2085
(16) Rogers, A. W.; Moran, J. F. Anal. Biochem. 1966, 16, 206-219. (19) Gogan, F.; Gogan, P. Anal. Blochem. 1974, 60, 363-371. (20) Rummerfield, P. S.;Goldman, I. H. Int. J. Appl. Radiat. Isot. 1972, 23,353-360. (21) Malcolm, P. J.; Stanley, P. E. Int. J. Appi. Radiat. Isot. 1976, 27, 415-430. (22) Noujalm, A.; Ediss, C.; Wlebe, L. In ref 5; pp 705-717. (23) Bush, E. T. Anal. Chem. 1963, 35, 1024-1029. (24) Sakashlta, M. Radioisotopes 1979, 26, 760-762. (25) Allen, H. J. Int. J. Appl. Radiat. Isot. 1976, 27, 662-663. (26) Carter, G. W.; Van Dyke, K. Anal. Biochem. 1973, 54, 624-627. (27) Simpson, E.; Browning, M. C. K. Clin. Chlm. Acta 1973, 45, 135-143. (28) Moore, P. A. Clln. Chem. (Winston-Salem, N.C.) 1961, 27, 609-611. (29) Gray, P.; Van Reenen, 0.: Potgieter, G. M. Clln. Chem. (Winston-Sa/em, N.C.) 1960, 26, 1886-1887. (30) Benson, R. H. Int. J. Appl. Radiat. Isot. 1976, 2 7 , 867-674. (31) Leaflet from Packard Instrument Co., Downers Grove, IL. (32) Hendee, W. R.; Ibbott, G. S.;Crusha, K. L. Int. J. Appl. Radiat. Isot. 1972, 2 3 , 90-95. (33) Peng, C. T. Anal. Chem. 1960, 32, 1292-1296. (34) Peng, C. T. Anal. Chem. 1969, 41, 16-21. (35) Bennett, C. A.; Franklin, N. L. "Statistical Analysis in Chemistry and the Chemical Industry"; Why: New York, 1954. (36) Brownlee, K. A. "Statistical Theory and Methodology In Science and Engineering", 2nd ed.; Wiley: New York, 1965. (37) Heitzmann, M. W.; Ford, L. A. Anal. Chem. 1961, 5 3 , 1721-1723. (38) Painton, C. C.; Mottola, H. A. Anal. Chem. 1981, 53, 1713-1715. (39) Hope, H. J. Anal. Biochem. 1973, 5 3 , 295-298.
RECEIVED for review April 7 , 1982. Accepted June 14, 1982. This work was carried out as a part of studies supported by Karolinska Institutet, LEO Research Foundation, and the Swedish Medical Research Council (19P-6483 and 13X-2819).
Selectivity of the Potentiometric Ammonia Gas Sensing Electrode M. E. Lopez and G. A. Rechnltz" Department of Chemistry, University of Debsware, Newark, Delaware 197 1 1
The basicity of amine! lnterferents wasi found to be a more important factor than volatllity in determlnlng the selectlvlty of the Orlon ammonlia gas senslng electrode. Theoretical selectlvlty coefficients of seven volatile amine interferents were successfully calculated from fundamental conslderations and were found to bo In excellent agreement with experimental values provided corrections are made for the nonideal response of the Orlon Inner pH-sensing element. A dependence of the pH of the thin fllm of internal electrolyte on the osmolarity of the sample solution and InV,ernalelectrolyte was observed. The generality of the theoretlcal approach was demonstrated with a imethylamlne-based sensor where experimental selectivity coefflcients of arnmonla and dimethylamine agreed well with predlcted values.
Although the interference of volatile amines on the response of potentiometric membrane electrodes for ammonia has been documented in the literature (1-5) and recognized by the manufacturer of commercial electrodes (6),little effort has yet been made to provide a quantitative explanation of observed interference patterns in terms of fundamental parameters. A very recent paper (7) considered the response of the ammonia electrode tcr several amines from an analytical perspective.
The present study was undertaken to provide a systematic evaluation of the potentiometric response to a series of amines and other nitrogen-containing compounds having a range of volatilities and basicities. A model is proposed for the steady-state response of the electrode which takes into account the effect of amine dissociation on the mass action equilibria responsible for the electrode potential in the presence of ammonia or the volatile amines. It will be seen that quantitative agreement between experimentally measured and theoretically calculated selectivity coefficients can be obtained provided correction is made for the nonideal response of the internal sensing electrode and provided the properties of the electrolyte employed as the filling solution are known. The procedures employed and the model proposed here point the way to a quantitative explanation of the interference characteristics of potentiometric gas-sensing membrane electrodes. The model is further tested and confirmed through the construction of a gas-sensing electrode which utilizes a methylamine-based filling solution; such a sensor is seen to have the predicted response to ammonia as an interferent. EXPERIMENTAL SECTION Apparatus and Materials. Potentiometric measurements were made with a Corning Model 12 Research pH meter and a Health/Zenith Model SR-204 strip-chart recorder. The cell was maintained at 25 "C with a Haake constant temperature circulator, Model FS. The Orion ammonia electrode, Model 95-10, using
0003-2700/82/0354-2085$01.25/00 1982 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 54,
NO. 12, OCTOBER
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the Orion ammonia electrode filling solution (95-10-02) or 0.05 M ammonium chloride in 0.05 M sodium chloride aqueous filling solution, was employed. A Corning semimicro pH combination electrode (476050) was used to determine the pH response characteristics of the Orion inner pH-sensing element. Ammonia, methylamine, dimethylamine, trimethylamine, and diethylamine were weighed as chloride salts. The normality of their 0.1 M aqueous solutions was verified by a Gran's plot titration with silver nitrate using an Orion silver sulfide electrode, Model 94-16A, and an Orion double-junction reference electrode, Model 90-02. The normality of 0.1 M aqueous solutions of ethylamine, triethylamine, and n-propylamine was determined by back-titration of excess standard hydrochloric acid with standard sodium hydroxide using methyl red indicator. Reagents. The 0.1 M standard solutions of ammonium chloride and of the amines listed below were prepared with reagent grade chemicals and deionized water. Ammonium chloride was obtained from Fisher Scientific Co., Fair Lawn, NJ. Methylamine hydrochloride, dimethylamine hydrochloride, trimethylamine hydrochloride, diethylamine hydrochloride, ethylamine (70% in water), triethylamine, and n-propylamine were obtained from Aldrich Chemical Co., Inc., Milwaukee, WI. Potassium phosphate buffer, pH 12.5, I = 0.1 M, was used as sample medium in all work. Procedures Used in Selectivity Study. Preliminary work entailed construction of calibration curves for 13 nitrogen-containing compounds in pH 11.5 phosphate buffer using the Orion Model 95-10 ammonia electrode immersed in the sample solution. Further work was restricted to seven aliphatic amines studied in pH 12.5 phosphate buffer. A 0.5 cm thick circular Teflon spacer was employed in all work so as to minimize volatile amine and ammonia loss. Calibration curves were constructed by progressive additions of a standard 0.1 M aqueous solution to 10 mL of phosphate buffer. An ammonia calibration curve was constructed prior to constructing duplicate calibration curves for each amine as detailed below. The gas-permeable membrane was replaced after duplicating calibration curves for each amine. All work was carried out at 25 OC. Duplicate calibration curves for the amines were constructed using the Orion filling solution and with the electrode suspended 0.5 cm above the sample solution by the Teflon spacer. The outer O-ring of the Teflon spacer created a seal between the electrode body and the sample cell wall. Duplicate calibration curves using the Orion filling solution were then constructed with the lower portion of the electrode immersed in the sample solution and the Teflon spacer almost touching the sample solution. Duplicate calibrationcurves were also constructed using an 0.05 M NHICl in 0.05 M NaCl internal electrolyte and with the electrode immersed in the sample solution. Procedure Used To Obtain pH Response Characteristics of Orion Inner pH-Sensing Element. The pH of a series of buffers spanning pH 6.5 to pH 10 was simultaneously measured with the Orion inner pH-sensing element and a Corning semimicro combination pH electrode; the inner Ag/AgCl electrode of the latter served as the reference for both pH electrodes. Linear regression analysis of the two sets of pH readings gave an equation which was used to correct for the observed nonideal response of the Orion inner pH-sensing element between pH 9 and pH 10. The potential difference between the inner Ag/AgCI electrodes of the Orion NHBgas sensor and the Corning pH electrode was determined in 0.05 M NH&l in 0.05 M NaCl electrolyte. Thus, once the meter was calibrated using the Corning pH electrode in pH 7 and pH 10 buffers, pH readings obtained with the Orion NH, gas-sensing electrode were adjusted at a later point to give the actual pH. Procedure Used To Study the Effect of Osmolarity on pH Readines Obtained with Ammonia Sensor. DH 12.5 DhosDhate - bufferslof 0.13,0.21, and 0.41 M were prepared by appropriate additions of potassium chloride. The electrode and sample cell were placed in a glovebag purged with nitrogen. Duplicate curves of pH vs. ammonia concentration were constructed in the three buffers. A dependence on the buffer osmolarity was observed. Procedure Used To Study the Selectivity Characteristics of a Methylamine-BasedSensor. A methylamine-based sensor was constructed by substituting 0.01 M methylamine hydrochloride in 0.01 M potassium chloride for the internal filling
Flgure 1. Schematic of NH, sensor immersed in sample solution: (a) internal pH sensor, (b) internal electrolyte, (c)Teflon spacer, (d) sample solution, (e) gas permeable membrane, (f) magnetic stirring bar, (9) thermostated glass cell.
Table I. Preliminary Screening of Possible Interferents compd no. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 a
compound pyridine phenylhydrazine ammonia benzylamine ethanolamine trimethylamine ethylenediamine ethylamine methylamine cyclohexylamine triethylamine dimethylamine diethylamine piperidine
P,a
torr 20 1
7600 1 1
2300 10 1060 2700 9 80 1800 240 30
pK,
AE, mV
4.60 5.21 9.25 9.35 9.50 9.74 9.93 10.63 10.64 10.65 10.67 10.73 10.93 11.20
-137 -154 -40 -154 21 -144 55
50 35 51 44 62 60
Values obtained from 1 6 and 17.
solution of the Orion ammonia sensor. Selectivity coefficients of ammonia and dimethylamine were determined following the procedure used for the ammonia sensor. RESULTS AND DISCUSSION The Orion Model 95-10 ammonia electrode consists of a pH combination electrode inserted in a cylindrical plastic tube fitted with a bottom cap and filled with an ammonium salt solution (Figure 1). A thin film of the filling solution is formed between the pH-sensitive glass tip and a gas-permeable membrane when the electrode is assembled. The discriminating ability of the hydrophobic, microporous membrane, mean pore size 0.2 pm, results from the properties of Teflon. The microporous structure creates an air layer through which only neutral, gaseous species can diffuse (8). However, organic liquids, such as surfactants, can also penetrate if able to overcome the surface tension at the solution-membrane interface (6). The results of the initial selectivity study of 13 compounds carried out with this electrode are shown in Table I. Ah' values listed in the third column give a rough estimate of the response of the compound relative to ammonia when both are present at M concentrations. A positive AE indicates a larger potential response to the compound than to ammonia and that the selectivity coefficient would be greater than unity. The compounds show wide variations in pK, (5.2 to 11.2), partial pressure (1to 2700 torr), and structure (aliphatic, aromatic, and heterocyclic). Calibration curves constructed
ANALYTICAL CHEMISTRY,, VOL.
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Table 11. Selectivity Properties of Volatile Amines compd no.
experimental compound ammonia trimethylamine propylamine ethylamine methylamine triethylamine dimethylamine diethylamine
P,torr 7600 2300 300 1060 2700 80 1800 24 0
pKa 9.26 9.74 10.53 10.63 10.64 10.67 10.73 10.93
k",,
BPot
'k"
3+
B~~~
1 2.8 16.3
1 2.3
17.8 20.5 24.4 37.8
17.4 20.5 22.1 25.5 33.3
theoretical
k NH,, B~~~ 1 3.3 17.8 19.6 22.1 22.8 27.0 41.8
~ N H , B~~~ ,
1 3.0 18.3 23.4 24.2 24.9 29.8 46.1
~ N H , B~~~ , e
1 3.0 17.2 21.6 22.4 23.0 21.3 41.3
Using Orion filling solution with a Using Orion filling solution with electrode suspended in air above sample solution. Using 0.05 M NH,Cl in 0.05 M NaCl with electrode immersed in sample solution. electrode immersed i n sample solution. Theoretical selectivity coefficients (uncorrected). e Theoretical selectivity coefficients corrected for error in pH readings obtained with Orion inner pH-sensing' element.
--
I 20-
0-
-20-
-40-
-60
1 -1ooc
..Io g [C 0M PO U N D ] ,mo I/ L - l o g [COMPOUND] , m O l / L
Figure 2. Response of WH3 sensor using Orion filling solution to (1) benzylamine, (2) ammonia, (3) trimethylamine, (4) cyclohexylamine, (5) methylamine, (6) dimethylamine, (7) triethylamine, (8) ethylamine, (9) piperidine, and (10) diethylamine with electrode immersed In sample solution (data points at 1 X M, not shown).
for those compounds giving a potential response above base line are shown in Figure 2. Table I shows that of the first seven compounds, all having pK, values less than 10, only the potential response to trimethylamine is of the same magnitude as that to ammonia. A larger potential resplonse to the remairing seven compounds, which have pK, values greater than 10, was observed. It appears, then, that a potential response of comparable or larger magnitude than that observed to ammonia requires that the compound be as basic or more basic tlhan ammonia. An interesting observation is that of the three compounds with aromatic functional groups, only benzylamine gave 'a response. However, the response was significantly less than that to ammonia although benzylamine is more basic. In studying the effect of membrane characteristics on the selectivity of the C 0 2 electrode, Kobos et al. observed substantially reduced potential response to aromatic acids upon replacing the commercial silicone rubber membrane with the Orion ammonia microporous membrane (9). Perhaps repulsive interactions prevent (diffusion of these aromatic moieties through the fluoropo1,ymeric membrane. Detailed selectivity studies were limited to the seven basic, volatile amines listed in Table 11. With the Orion filling solution as the internal electrolyte, calibration curves for these
Figure 3. Response of NH3 sensor using 0.05 M NH,CI in 0.05 M NaCl electrolyte to (1) ammonia, (2)trimethylamine, (3) propylamine, (4) ethylamine, (5) methylamine, (6) triethylamine, (7) dimethylamine, and ( 8 ) diethylamine wlth electrode immersed in sample solution.
compounds were constructed with the ammonia sensor suspended in air above the sample solution and also immersed in the sample solution. Experimental selectivity coefficients, kNHs,BPot, for the Orion ammonia sensor determined at 3.2 X M concentrations for the two electrode configurations are listed in Table 11. These coefficients were calculated by use of the Eisenman-Nicolsky equation (IO). where El and E2 are the potential readings observed for the same concentrations of ammonia and amine, respectively, in separate solutions and S is the slope of the ammonia calibration curve. The experimental selectivity coefficients obtained when the electrode was suspended in air or immersed in solution were not significantly different, suggesting that the potential response was largely due to the diffusion of gaseous amine across the membrane with little, if any, liquid amine penetration. Calibration curves for the volatile amines constructed upon replacing the Orion filling solution with an 0.05 M NH4C1in 0.05 M NaCl electrolyte are shown in Figure 3. Experimental selectivity coefficients determined at 4 X loy3M concentrations are listed in Table 11. The closeness of the experimentally deterwined selectivity coefficients obtained with the
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ANALYTICAL CHEMISTRY, VOL. 54, NO. 12, OCTOBER 1982
Orion filling solution and the 0.05 M NH4C1in 0.05 M NaCl electrolyte is not surprising since they have the same osmolarity, 0.02 M ( I I ) , and approximately the same ammonium salt concentration, 0.05 M. The Orion filling solution is purported to be a saturated ammonium picrate solution of about 0.05 M (12). Steady-state response times observed with the electrode suspended in air using the Orion filling solution were a minimum of 1 h at 5 X M amine concentrations. Slopes ranged from 93 mV/decade to 67 mV/decade. Response time5 observed with the electrode immersed in the sample solution, again using the Orion filling solution, averaged 20 min for the same concentrations. The slopes ranged from 65 mV/decade to 59 mvldecade. Parallel experiments were conducted with a filling solution of known composition. Steady-state response times observed, for the same amines at the same concentration, upon replacing the Orion f i g solution with an 0.05 M NH4Cl in 0.05 M NaCl electrolyte and the electrode immersed in the sample solution were shortened to 6 min and appeared to be independent of the nature of the amine. The slopes of the amine calibration curves were similar to those observed using the Orion filling solution and the electrode immersed in the sample solution. As seen in Figure 3, the response of the Orion ammonia sensor to the amines listed in Table I1 is larger than that of ammonia, the relative magnitude increasing with the basicity of the amines. However, no correlation of potential response with amine volatility can be made. The steady-state response of the HNU commercial ammonia electrode to the volatile amines was reported in the literature (7) to have the same selectivity pattern reported here. However, when using an automated arrangement, Meyerhoff et al. (7) observed essentially no volatile amine interference on the nonequilibrium response of the ammonia sensor. The better selectivity was attributed to a slower rate of diffusion of the amines relative to ammonia through the gas-permeable membrane. Slower diffusion rates may result in a nonequilibrium distribution across the membrane at lower amine concentrations (ca. 5 X M), perhaps accounting for the slightly higher response slopes observed for the amines compared with the Nernstian slope of 59.7 mV/decade observed for ammonia. However, the response time of the ammonia sensor for a %fold increase M was 2 min, as fast as in amine concentration at 2 X M that for ammonia. In addition, the diffusion at 4 X amine concentration was found to be reversible in that the same starting potential could be obtained once the electrode was reconditioned in water or pure buffer for 10-30 min. The above observations suggest that the observed steadystate response to an amine concentration of 2 X M or greater reflects an equilibrium concentration existing between the two solution phases and the air-layer defined by the microporous Teflon membrane. Calculated amine partition coefficients, expressed as the ratio of the molar concentrations in the gas phase (air layer) and the aqueous phase (sample solution or thin film of internal electrolyte), are of the same magnitude as that of ammonia. The fraction of amine which is present in the thin film of internal electrolyte is the same as that of ammonia at the same sample concentrations, permitting comparison of steady-state responses based on the chemical equilibria existing in the thin film. The hydrogen ion activity in the thin film of internal electrolyte (13) can be calculated upon substitution of the appropriate equilibrium equations
Ka = ("J(H+)/(NH,+) Ka' = (RN)(H+)/ (RNH') into the charge balance equation for the film [NH4+] [RNH+] [H+] = [Cl-] + [OH-]
+
+
(2)
(3)
(4)
8.50k
8.00-
I I
n
I
I 7.50-
L
1
1
l
l
1
l
3.00
1
1
2.50
-log ~AMMONlA],mol/L
Flgure 4. Effect of relative sample osmolarity at (0)0.13 M, (0)0.21 M, and (0)0.41 M on pH profiles of NH, calibration curves obtained with NH, sensor using 0.05 M NH,CI in 0.05 M NaCl electrolyte.
where RN and RNH+ refer to the unprotonated and protonated amine, respectively. The chloride concentration is set equal to the ammonium chloride salt concentration in the internal electrolyte. The presence of added neutral salt such as NaCl can be ignored. The resulting equation is a secondorder polynomial in hydrogen ion activity.
where y and y' are the activity coefficients for NH4+and the protonated amine, respectively. By use of this approach, selectivity coefficients for the amines can be calculated a priori with only knowledge of the molarity of the ammonium ion and the equilibrium constant for the hydrolysis of the protonated amine. The selectivity coefficients are expressed as the ratio of the hydrogen ion activity in the electrolyte film resulting from the same nominal concentrations of ammonia and amine in separate solutions. Selectivity coefficients calculated in this manner can then be compared with experimental values once certain corrections are made. To verify the value of the hydrogen ion activity calculated using eq 5, we experimentally measured the pH of the thin film of internal electrolyte at an ammonia sample concentration of 4 x M. The resulting experimental value of pH 8.45 was higher than the expected theoretical value of pH 8.28. It was also observed that the pH of the thin film increased once the electrode was immersed in pure buffer solution, contradicting the expected and normally observed behavior (14). This latter phenomenon has been reported in the literature for a methylamine air-gap electrode and attributed to C02desorption from the electrolyte film (15). An ammonia calibration curve was then constructed in a glovebag purged with N2 to eliminate C02from consideration;however, identical results were still obtained. It has been observed, incidentally, that a pH reading recorded when the system had reached equilibrium changed to a higher pH value when the system was left intact overnight. It was suspected, then, that the observed pH increase of the electrolyte film was due to the diffusion of water vapor from the sample solution, through
ANALYTICAL CHEMISTRY, VOL. 54, NO. 12, OCTOBER 1982
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Table 111. Selectivity Coefficients for Methylene Sensora compound aimmonia dimethylamine
kMeNH,, BPot
kMeNH,, BPot
0.31 1.2
pot d
~ M ~ N HB, ,
0.06
0.04
1.23
1.21
M. t~ Values interpolated from published results (ref 15). Values Selectivity coefficients determined at 4 x Corrected theoretical selectivity coefficients. experimentally determined using 0.01 M MeNH,Cl, 0.01 M KC1. a
c
10.0-
I
z r
5 & -0 -L i8 -, -I o g [C
o M PO uN
4
,"IO
II L
Figure 5. Response af methylamlne-based sensor using 0.01 M MeNH,CI in 0.01 M KCI electrolyte to (0)ammonia, (0)methylamine, and ( 0 )dimethylamine. the membrane, into the electrolyte film. This would result in a dilution of the weak acid salt and, thus, yield a higher pH value. The osmolarity of the pH 12.5 phosphate buffer was calculated to be 0.13 M while that of 0.05 M NH&l in 0.05 M NaCl was 0.20 M so that diffusion in the above direction would be possible. The same ammonia calibration curve was then consitsucted with 0.21 M and 0.41 M, pH 12.5 phosphate buffers. The dependence of the p H vs. concentration calibration curves on the relative osmolarities of the sample solution and the internal electrolyte is seen in Figure 4. The p H observed a t 4 X Nt ammonia in p H 12.5 phosphate buffer of equal osmolarity to that of the internal electrolyte (0.20 M) was 8.30, which is in excellent agreement with the theoretical value. Selectivity coefficients of the amines calculated using eq 5 are listed in Table 11. They are seen to increase with the basicity of the amine and are all greater than unity, indicating greater selectivity for the amine than for ammonia. In order to compare these theoretical selectivity coefficients with the experimental values obtained by using the 0.05 M NH&l in 0.05 M EX1 filling solution, we had to make a correction for the nonideal response of the Orion inner pH sensor at high pH values (pH 9-pH 10). The corrected theoretical selectivity coefficients are given in the last column of Table 11. These show good agreement with the experimental selectivity coefficients.
To test the general applicability of eq 5 to other gas sensors and, also, the reversibility of the present system, the response of a methylamine-based sensor to methylamine, ammonia, and dimethylamine was similarly studied (Figure 5). The same Orion ammonia sensor was employed, but the internal filling solution was replaced by 0.01 M MeNH3Cl in 0.01 MKCl, an electrolyte composition employed with a methylamine air-gap electrode (25). The experimental selectivity coefficients compare favorably with the theoretical values listed in Table 111. It is of interest that there is reasonable agreement with interpolated values from the literature. The theoretical selectivity coefficients calculated by use of eq 5 are considered to be true equilibrium constants, judging from the reversibility of the amine diffusion through the microporous Teflon membrane. If the membrane were replaced by a more discriminating membrane, equilibrium between the two solution layers may require inordinately long times, in which case, an operational selectivity coefficient may have to be defined. Since the selectivity coefficients of the amine interferents of the Orion ammonia sensor are all greater than unity and the response times for these volatile amines are about the same as that for ammonia, due care should be taken in measuring ammonia in the presence of volatile amines. The effect of a possible amine interference on the determination of a nominal ammonia concentration in a sample can now be predicted on the basis of fundamental considerations.
LITERATURE CITED Midgley, D.; Torrance, K. Analyst (London) 1972, 9 7 , 626-833. Beckett, M. J.; Wilson, A. L. Water Res. 1974, 8 , 333-340. Evans, W. H.; Partridge, B. F. Analyst (London) 1974, 99, 367-375. Gilbert, T. R.; Clay, A. M. Anal. Chem. 1973, 45, 1757-1759. Banwart, W. L.; Tabatabai, M. A.; Bremner, J. M. Commun. Soil Sci. Plant Anal. 1972, 3 ,449-458. "Instruction Manual Ammonia Electrode Model 95-10"; Orion Research, Inc.: Cambridge, MA, 1979. Meyerhoff, M. E.; Fraticelli, Y . M. Anal. Chem. 1981, 53, 1857-1861. Ross, J. W.; Riseman, J. H.; Krueger, J. A. Pure Appl. Chem. 1973, 36,473-487. Kobos, R. K.; Parks, S.J.; Meyerhoff, M. E. Anal. Chem. 1982, 54, 1976-1 980.
Analytical Chemistry Division, Commission on Analytical Nomenclature Pure Appl. Chem. 1976, 48, 129-132.
Arnold, M. A. Ph.D. Dlssertation, University of Delaware, Newark, DE, 1982.
Riseman, J. H.; Krueger, J. A,; Frant, M. S. U S . Patent 3830718, 1974.
Van der Pol, F. Anal. Chim. Acta 1978, 9 7 , 245-252. Masclni, M.; Cremisini, C. Chim. Ind. ( M a n ) 1980, 62,222-230. Hslung, K. P.; Kuan, S. S.,Guilbault, G. G. Anal. Chim. Acta 1976, 84, 15-22.
Felslng, W. A.; Phillips, B. A. J . A m . Chem. SOC. 1936, 58, 1973-1975.
Weast, R . C., Ed. "CRC Handbook of Chemistry and Physics," 53rd ed.; The Chemical Rubber Co.: Cleveland, OH, 1972.
RECEIVED for review May 27, 1982. Accepted July 23,1982. The support of Grant GM-25308 of the National Institutes of Health is gratefully acknowledged.
