Low Level Sodium Ion Measurement with the Glass Electrode Edgar L. Eckfeldt and William E. Proctor, Jr. Corporate Research Department, Leeds & Northrup Company, North Wales, Pa. Power plants and other industries need continuous measurement of sodium ion, both at higher concentrations and also in the range from 0.1 to 1.0 ppb, to determine quality of process water. Laboratory and field studies have shown that sodium, even at the lower level, can be successfully monitored over long periods of time if attention is given to several conventionally overlooked factors. A strong base is preferred to economically reach the high pH of sample (11.0-11.6) needed for adequate suppression of hydrogen ion; also preferred is one that produces a cation (dimethylamine, or better, diisopropylamine) that exerts minimal electrode interference. Other troublesome interferences and ways to avoid them are discussed, including potassium and silver ion interference from the reference electrode solution during both normal and stopped flow sample conditions and electrical noise originating at the liquid junction.
measure down to 1 ppb with a maximum uncertainty of 1.8 ppb. Sodium concentration changes at lower levels than this, however, are said to be discernible. In the other study, Webber and Wilson ( 4 ) claim that the electrode can be made to follow the Nernst equation down to about 1 ppb by controlling the p H of the sample and by using continuously flowing samples. They found standard deviations of 0.4 to 1.2 ppb in measuring low concentrations of sodium. The present authors, along with W. A. Lower and W. D. Howie of Duquesne Light Company, published results of an extensive field study of the glass electrode for measuring sodium ion in various power plant applications (5). The work featured several innovations, including the use of base agents other than ammonia. It was concluded that measurement of sodium ion, even in the low range of 0.1 to 1.0 ppb, should be no more difficult to carry out than conventional p H measurement. A measurement as low as 0.07 ppb was reported on one specially purified sample. Those results represent a sensitivity increase over results reported by others, corresponding to at least one order of magnitude lower in sodium ion measurement. Capability of measuring at such low levels is important, because good deionizer effluents and much of the hot well condensate water of high-pressure steam plants contain sodium in the 0.1 t o 1.O ppb range. A novel calibration technique was used which circumvented a problem that has confronted other workers who have attempted to measure very low levels of sodium ion-namely, the problem of knowing the background sodium content of reagent water used to prepare standard sodium solutions. The method determines background sodium content by introducing a known sodium concentration increment and calculating the initial solution concentration using the known logarithmic (Nernstian) response characteristic of the electrode. Adherence of the electrode to theoretical Nernstian response can be checked down to very low sodium ion levels by using small increments in conjunction with larger increments. Details of the calibration technique have been described (5, 6). Although that method of calibration adds greatly to overall measurement confidence, the improvements in electrode reliability and sensitivity are attributable to other factors. The purpose of the present paper is to describe laboratory studies that elucidate the fundamentals involved and to explain preferred operating conditions. Early sections of the paper deal with problems of hydrogen ion and base ion interference; the later sections discuss instrumental considerations, including equipment design factors and other conditions that lead to good performance.
SOhlE AREAS OF MODERN industry require very high purity Of materials to support the advanced technology achieved. Examples are the semiconductor industry and electric power generation. Process water used in these industries may exceed in purity, for example, the distilled water commonly used in many laboratories. To meet the purity requirements, industry needs sensitive, convenient, and reliable analytical techniques. The ubiquitous character of sodium in our environment makes sodium an effective telltale to reveal the condition of a high purity water system and to signal trouble when contamination enters. Flame photometry can serve to determine sodium at rather low concentrations, but the measurements involved are not easy to carry out, nor is the method easily and inexpensively adaptable to continuous monitoring instrumentation. The sodium selective glass electrode offers an attractive alternative. Although many people have contributed, Eisenman ( I ) has contributed outstandingly to theoretical and experimental knowledge of the electrochemistry of cation-sensitive glass electrodes (see Reference I for a comprehensive discussion). Gurney ( 2 ) pioneered the use of the glass electrode for measuring sodium in process water of power plants. In contrast to flame photometry, glass electrode measurements are easy to carry out, and a successful monitor of this type offers considerable practical advantage. Accordingly, this electrode measurement has generated much interest and is the basis of several commercial instruments now on the market. Two careful recent studies have shown the current status of low-level sodium ion measurement with the glass electrode. Following usual practice, ammonia was used in those investigations as the base agent to raise pH. In one of the studies, Hawthorn and Ray (3)concluded that the electrode follows the Nernst equation over concentrations ranging from 25 ppb to 25 ppm. Below 25 ppb, corrections to compensate for background sodium in the water used to prepare solutions are said to be needed in order to allow use of the electrode to
Glass electrodes are not perfect sensors for sodium ion measurement and, depending on selectivity considerations, may respond to other ions present in the sample water ( I ) .
(1) G . Eisenman, “Advances in Analytical Chemistry and Instrumentation,” C . N. Reilley, Ed., Wiley, New York, N. Y., 4, 213 (1965). (2) W. B. Gurney, Elec. World, March 23, 1964, p 125. (3) D. Hawthorn and N. J. Ray, Analyst, 93, 158 (1968).
(4) H. M. Webber and A. L. Wilson, Analyst, 94, 209 (1969). (5) E. L,, Eckfeldt, W. E. Proctor, Jr., W. D. Howie, and W. A. Lower, Proc. 29th Inter. Water ConJ, Engineers Society of Western Pennsylvania, Pittsburgh, November 19-21,1968, p 109. (6) E. L. Eckfeldt, ISA Trans., 9, 37 (1970).
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ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
HYDROGEN-ION, BASE-ION INTERFERENCE
Some of the substances commonly present in power plant sample water might be expected to exert little or no interference effect, and this has been confirmed by experimental evidence (see the rather extensive studies of interferences reported in References 4 and 5). Electrode theory indicates, however, that monovalent cations might be troublesome ( I ) , and certain important ones have heretofore been overlooked. For low level sodium ion measurement, sample pH must be raised by addition of a base, as already noted, to avoid hydrogen ion interference. A satisfactory value for the ratio of H+ to Na+ is often quoted at and even lo-* has been advocated (7). Not generally appreciated is the interfering effect arising from the cation that comes from the base agent used to raise pH. If a hydrolyzable base, B, is added to water, it reacts to produce hydroxyl ion: B
+ HOH e BHOH +BH+ OHKB
(1)
For every hydroxyl ion produced, a corresponding base cation, BH+, is generated. The situation may be even worse if the sample water initially contains some acid, such as HC1. With a sample of this kind, a reaction of the base with the acid must first take place, as follows, before Reaction 1 can occur: B
+ H+ + C1-
+
BH+
+ C1-
JI
(2)
In this case a surplus of base cations over hydroxyl ions will result. In any case a rather high base cation activity will exist in the sample solution at the high pH value needed for low-level sodium ion measurement. Whether or not this constituent will interfere will depend on its activity level and on the selectivity characteristics of the electrode. Eisenman ( I ) has described the response of cation-sensitive electrodes to ammonium ion, as well as to a number of substituted ammonium ions. Rechnitz and coworkers investigated the response of a general cation-sensitive electrode to ammonium ion (8) and to various alkyl-substituted ammonium ions (9). They used neutralized hydrochloride salts in a THAM buffer for their measurements. The cations present were the same as those which would have been formed at lower concentration by adding the corresponding bases to water. They found that the electrode did respond to ammonium ion and several alkyl-substituted ammonium ions. These observations do not pertain directly to the problem of sodium ion measurement, because for sodium ion measurement one would normally choose an electrode that had good selectivity for sodium, rather than one of general cation response. Also, the solution conditions were quite different. Nevertheless, it was clear that the base cation interference must be considered, and experimental evidence in the laboratory and field has amply confirmed this view. EXPERIMENTAL
Substances of moderate price which would produce strongly basic aqueous solutions were sought. A number of organic amines appeared quite suitable for the purpose and certain of these were studied experimentally. These substances were added in gaseous or vapor form, thereby forestalling possible contamination from dissolved sodium compounds as might occur if the base were added directly in liquid form. Ex(7) M. Hams, M. de Heaulme, and P. Morin, Bull. SOC.Chim. France, 1963, 2658. (8) S. A. Katz and G. A. Rechnitz, Z . Anal. Chem., 196, 248 (1963). (9) G . A. Rechnitz and G . Kugler, ibid., 210, 174 (1965).
Figure 1. Experimental arrangement perimentation focused attention on the base cation effect, but hydrogen ion was not overlooked. Special efforts were exerted to keep sodium in the reagent water at the lowest possible level in order to avoid its effect. The equipment is illustrated in Figure 1. The vessels and lines were of polyethylene construction throughout, and the interior of the system was kept closed from the ambient. The flow cell K contained the sodium ion sensor M (LN 117,201 glass electrode). The lithium alumino silicate glass of the electrode bulb was especially developed for the purpose. This electrode has good selectivity for sodium ion and will give many months of satisfactory field service with little or no maintenance. The reference electrode L of saturated calomel type had a fused-in porous ceramic junction and was internally pressurized through line N by means of air pressure kept at 5 lb/in*. The range resistor was selected to provide 500 mV full scale on a Speedomax G recorder readout device. In most of the work the zero adjuster of the pH meter was set to give a recorder scale of +50 to -450 mV. The chemical bases studied were used as received from suppliers. Matheson Company supplied the gaseous bases in pressurized cylinders, monomethylamine (98.0%), monoethylamine (98.5 %), dimethylamine (99.0 %), and trimethylamine (99.0 %). Pennsalt Chemicals supplied anhydrous ammonia. Liquid bases in the laboratory work were supplied as follows: Eastman Organic Chemicals, 1396 diisopropylamine, 616 triethylamine, and P6760 N-methylmorpholine ; Fisher Scientific Company, D-46 diethylamine. In the field work the liquid bases (diethylamine and diisopropylamine) were the usual commercial grade supplied by Union Carbide Chemicals in 5-gal containers; the dimethylamine was the same as in the laboratory work. Reagent water was prepared by starting with distilled and deionized water in the elevated 5-gal bottle A (Figure 1). The water was passed in series through the ion exchange columns C and D and into the 5-gal bottle E. Next, the water completely filled the mixing flask F and flowed through tube H back into bottle A , which was now located in the indicated lower position. The water was circulated at least twice through the ion exchangers in this manner before being ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
333
I
0.10
\
I
I
I
I
I
rSODlUM
-\ \ AMMONIUM ( N H ~ ) Y
0
1
I
I
I
I
1
2
3
4
5
NEGATIVE LOGARITHM OF CATION MOLARITY 12 II 10 9 p H O F BASE SOLUTION
Figure 2. Electrode response to various cations A
B C D E
F G
J
DI ISOPROPYLAMINE DIMETHYLAMINE TRIETHYLAMINE N-METHYLMORPHOLIN TRIMETHYLAMINE METHYLAMINE ETHYLAMINE AMMONIA
with the flow of nitrogen entering flask F. In the case of liquid bases, transport system I1 was used; the base was put into vessel T and nitrogen flow through tubes V , W , and S carried vapor of the base into flask F. The hot water bath U accelerated the rate of transport of base. Nitrogen gas flow promoted mixing in flask F and prevented backflow into tube S . The ion concentrations of individual batches of base in water thus prepared were determined after passing the solution through the measurement cell K by measuring the pH of the effluent at P. For electrode measurements, solutions flowed a t a rate of about 100 ml/min and were at room temperature (approximately 25 "C). Higher concentrations of base cation, between 0.1 and 1.0 molar, were prepared by a somewhat different procedure, also designed to avoid contamination from sodium ion. Multiply-deionized, distilled water contained in a 5-gal polyethylene bottle was acidified by passing HCl cylinder gas over the surface of the water, while keeping the water briskly stirred with a magnetic stirrer. A portion of the hydrochloric acid solution thus prepared was run into flask F, where it was treated with excess base, by means of transport system I or I1 until its pH was in excess of 10.5. Small samples were withdrawn to carry out the pH test. Also, samples of the original HCl solution were withdrawn for titration with standardized base in order to determine the HCl concentration and, therefore, the approximate base cation concentration in the respective final solutions. The base cation solutions thus prepared were passed in turn through the cell for electrode measurement, but in each instance enough solution was retained in flask F to allow making up a quantitatively diluted solution, using deionized water. The diluted solutions were likewise submitted to electrode measurement. Additionally, electrode response measurements were made on several solutions of sodium chloride and potassium chloride which had been rendered basic with amine. Response to hydrogen ion was obtained on solutions of measured pH containing ammonium chloride and different amounts of acetic acid. To check calibration stability of the electrode, voltage measurements were made at the beginning and/or end of each working period using a basic 0.075M sodium chloride solution. RESULTS AND DISCUSSION
,
I I I I 2 3 4 NEGATIVE L O G A R I T H M OF C A T I O N M O L A R I T Y 12 II IO 1
0
9
pH OF E A S E S O L U T I O N
Figure 3. Electrode response with different base reagents retained in bottle E, preparatory to making test solutions. The treatment rinsed the system and lowered the sodium content of the water to the vicinity of 0.1 ppb. In preparing a base solution, a portion of the water in E was run into the flask F. A flow of cylinder nitrogen gas entered through tube R and passed into flask F (transport system I). Gaseous base entered through tube Q and mixed 334
ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
Results of Tests. Response characteristics of the electrode are summarized in Figures 2 and 3. Electrode potential values are referred to the saturated calomel electrode throughout and the IUPAC sign convention is used. Figure 2 portrays the electrode response situation in general and will help to explain the nature of the interference problem and what can be done to achieve low level sodium ion measurement. The approximate way in which the electrode responds in the presence of each of a number of different bases is shown in Figure 3, and this information will be helpful in choosing a suitable base. Discussion of Amine Tests. In all but the tests specifically on sodium solutions, the experimental precautions precluded sodium ion as a factor except for the unavoidable low background level. A background of 0.1 ppb corresponds to a voltage somewhat more negative than -0.400 volt. The voltage value corresponding to dotted line A of Figure 2 serves to represent a postulated low concentration of sodium ion which one may wish to measure. Since the electrode is more responsive to hydrogen ions than to sodium ions, the hydrogen ion content of even a rather high pH sample solution may introduce interference. The dotted line C of Figure 2 (relating to the pH abscissa) was arbitrarily drawn to represent hypothetical response to hydrogen ion concentration, which at each voltage point is
ljl00 that of the competitive sodium ion concentration determined by extrapolation of the theoretical sodium ion line. Although actual slope values are in some disagreement, the right-hand branches of the diisopropylamine and ammonia curves of Figure 2 nevertheless correlated with the hypothetical hydrogen response line C. The experimental results of Figure 3 substantiate the existence of a region of hydrogen on the right-hand side where a ion dominance-namely, number of the curves with upward slope lie nearly on top of each other. The dotted line C lies a short distance from the hydrogen ion-dominated portion of the experimental curves. F o r satisfactory sodium ion measurement, one needs a hydrogento-sodium ion ratio closer to ljl000 than to the ljl00 used in plotting line C. These results demonstrate that solution p H is a n important factor, if really low level sodium ion measurement is to be achieved. Solution p H must be elevated into the region of p H 11.0 to 12.2. The 117,201 electrode satisfactorily withstands protracted exposure to such strongly basic solution. In general each base of Figure 3 produces a unique curve. The differences presumably arise from individuality of the respective base cations. The straight line portion of the ammonia line, with a nearly theoretical Nernstian slope, definitely demonstrates a n ammonium ion response of the electrode. Although exhibiting more random behavior, the curves for the other bases nevertheless indicate corresponding base cation response. The dotted line B has been arbitrarily inserted in Figure 2 to illustrate the hypothetical base cation effect. The line has Nernstian slope and is placed close to the cation response branch of the diisopropylamine curve. If the intersection of the hypothetical dotted lines B and C lies below the postulated sodium ion line A , then sample conditions can theoretically be found to enable measurement of that presumed sodium ion concentration. On the other hand, if the sodium ion concentration lies below the intersection point, then conditions satisfactory for the measurement will theoretically not be attainable. As Figure 3 shows, the selection of the particular compound to be used as base reagent will markedly influence the position of the electrode response line arising from the base cation, and hence along with the p H factor will determine the lowest sodium ion level that can be measured. Intersection of line C with a n extrapolated ammonium line, for example, will take place at a much higher level than that obtained by extrapolating the cation branch of the diisopropyl line. This explains why diisopropylamine has proved in practice t o be a n efficient base for low level sodium ion measurement. The other bases tested (Figure 3) show cation response lines at somewhat less negative voltages than the diisopropylamine curve and hence might be expected to be inferior base agents. The difference, however, in the case of dimethylamine is not great. Sodium ion measurements were made in the field on a series of sample solutions, each of which was separately tested with diisopropylamine and dimethylamine. The results are summarized in Figure 4. At low sodium ion levels a n appreciable difference exists in the performance of the two bases, which disappears a t concentrations above 1 ppb. These tests likewise show the superiority of diisopropylamine for making very low sodium ion measurements. Individuality in behavior of the various bases undoubtedly is related to differences in molecular properties. In the study of alkyl-substituted ammonium ions referred t o earlier (9), cation size was considered the important factor-the larger
I
I
/I /
/
,
/'
/
'I I
I 0. I
I
1.0 SODIUM P P 0
( D I ISOPROPYLAMINE)
Figure 4. Comparative measurements with dimethylamine and diisoprop ylamine on common samples the cation, the less its effect in producing an electrode response. Recently, Rechnitz has suggested that adsorption phenomena may be more important than size (10). Although some exceptions may be noted, the present study nevertheless confirms the merit of having a base cation of large size (a base of large molecular weight) to avoid interference in low level sodium ion measurement. Matheson Company has furnished analytical information on the identity and approximate concentration of impurities typically present in their amines (ZZ). They give the ammonia impurity in respective bases as follows: methylamine, 0.0 %; ethylamine, 0.6 % ; dimethylamine, 0.0 %; trimethylamine, 0.0%. The appreciable amount of ammonia in ethylamine, contrasted with its indicated absence in the three other amines, may account for the poor showing of ethylamine in the present study. In the absence of ammonia impurity, the ethylamine curve G might have appeared lower on the graph, possibly below the methylamine curve F, in which case its position would have agreed with the cation size principle. No special attempt was made in the study to purify the bases (gas or liquid) prior to testing. Although a preliminary purification step would not be acceptable in practical situations, the manner of introducing the bases in practice (5), particularly in the case of liquids, should result in progressive purification of the supply through fractional distillation of the more volatile ammonia impurity. It should be noted that electrode response to low-level hydrogen ion and cations of the organic amines studied was somewhat less definite and less reproducible than in the case of the ammonium, potassium, and sodium ions. The electrode takes time t o equilibrate. Also, the undissociated base molecules may or may not exert a complicating influence. Furthermore, some of the observed curvature and nonNernstian character of the data may be attributable to mixed response of the electrode, which is trying to adjust its potential, for example, to the competing influences of sodium ion (line A ) , base cation (line B), and hydrogen ion (line C). Other work has shown satisfactory electrode response to changes in (10) G. A. Rechnitz, discussion following presentation of the
present paper at the Pittsburgh Conference, Cleveland, Ohio, March 5 , 1970. (11) Matheson Gas Products, East Rutherford, N. J . , Data Sheets on monomethylamine, dimethylamine, trimethylamine, and monoethylamine, communication received May 5 , 1970. ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
335
Tablc I. Value of Dissociation Constant and Consumption of Base calculated at solution pH of 11.3.~ . - Quantities . __ _ . Dissociation Degree of Grams of Pounds of coilstant, 25 "C dissoclalloll Solution base per hase used ( x 106) Mol wt of hase molarity liter per month" ~~
Base reagent Morpholine 0 . 244b 87.1 0.0012 Ammonia 1 79c 17.0 0.0088 Trimethylamine 5,45r 59,l 0.027 Dimethylamine 52' 45.1 0.21 Diisopropylamine 112J 101.2 0.36 a At assumed sample flow rate of 100 ml per min (1.6 gal per hour). b A. R. Ingram and W. F. Luder, J . Anier. Clrcm. Sw., 64, 3043 (1942). c H.S. Harned and B. B. Owen, ibid., 52, 5079 (1930). N. F. Hall and M. R. Sprinkle, ibid., 54, 3469 (1932). I
-0.50
I
AMMONIA ADDED TO WATER DI ISOPROPYLAMINE ADDED
TO AN AMMONIA SOLUTION A
-4
i? -P 40
PH
Figure 5. Ammonium ion suppression by addition of diisopropylamine sodium ion concentration a t low levels, when pH and the amine concentration are kept constant (5). Advantages of a Strong Base. A large dissociation constant of the base reagent will be helpful in producing the desired high p H of the sample solution. A greater degree of dissociation results from a constant of larger value. The situation is illustrated by the approximate data of Table I. These data were calculated assuming that the listed base reagents were added to water in each case in amounts to produce a solution p H of 11.3, an arbitrary value within the recommended pH range. Bases at the bottom of the table with large dissociation constants exhibit a much greater degree of dissociation than the bases at the top of the table. Accordingly, as the table shows, a much lower solution molarity is needed to produce the specified pH. Needed amounts of base expressed on a weight basis will reflect the molecular weight, but this factor is not as important as the dissociation constant. The last column of Table I presents base consumption rates for a sample flow rate that might be reasonably established in practice. The amount of morpholine needed to keep equipment running is prohibitively large, thus ruling out the use of this base and any other base with a correspondingly low dissociation constant. If ammonia were not excluded for other reasons, its dissociation constant might be regarded as sufficiently large to make this a practical base, especially in view of its low molecular weight. To avoid base cation interference, as previously described, a base of larger molecular weight is needed and a dissociation constant larger than that of ammonia will be desirable. The constant for trimethylamine ( 5 X 10V) might be regarded as minimal. The advantage of using a much stronger base is illustrated by the tabulated data for dimethy lamine and diisopropylamine. 336
ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
1.64 0.23
0.075 0.0097 0 0056 I
143 3.9 4.4 0.44 0.57
~
1360 37
42 4.3 5.4
For these bases, only a relatively small amount of each compound is needed to reach the specified pH, thus making thew compounds economical to use as base reagents. As pointed out previously, ammonium ion is an interference in measuring sodium at low levels. Ammoni~iniion often is present in process sample water, and its concentration may be several orders of magnitude higher than that of the sodium ion to be measured. A high hydroxyl ion concentration, resulting from thc use of a strong base reagent, will substantially remove ammonium ion from the solution, making it go into the form of undissociated ammonium hydroxide. Calculation indicates that of the total ammonium (NH4+ and ",OH) present in a solution of p H 10.5, only about 6% will be in the ionized form. At 11.0, the ionized fraction drops to about 2%, while at pH 11.5 the ionization further diminishes to about 0.6% (12). Experimental evidence of the beneficial effect of the strong base diisopropylamine in suppressing ammonium ion concentration is shown in Figure 5. PREFERRED OPERATING CONDITIONS
Several factors in addition to the choice of base reagent are important in achieving good performance. Some of these are related directly or indirectly to the flow of sample solution. Cell. In a flow cell of conventional type, we have observed interference arising from potassium ion coming from the reference electrode, even though the reference electrode was located downstream from the sensor electrode. The explanation appears to be turbulence generated in the flowing sample stream by electrodes that are mounted at right angles to the flow path. The turbulence causes some regions of the flowing solution to move in a direction opposite to the nominal flow direction. If potassium chloride solution leaking from the reference junction is caught in one of these eddies, it may be carried upstream to the bulb of the sodium electrode, where it can generate a significant interference voltage of a sporadic or cyclic nature. This type of potassium ion interference was eliminated by using the cell illustrated in Figure 6. Sample water enters pipe A at the bottom of the sodium electrode compartment, promptly contacts the sensor bulb, and rises through the annular space surrounding the electrode stem. At the crossaperture B the solution flows into the reference electrode compartment and descends to the bottom of the annular space, where it contacts the junction of the reference electrode. The solution then passes into the temperature compensator compartment, where it rises through the annular (12) D.S. McKinney, ASTM Proc., 41, 1285 (1941).
space and discharges from outlet C. (In the present study, a thermometer replaced the temperature compensator shown.) It will be noted that the advancing flow of solution through the nonturbulent annular spaces adjacent to the electrodes will prevent the backward passage of potassium chloride solution. The more dense potassium chloride solution will not flow upward against the current and is carried away with the discharge. When solution flows through the cell, its level is determined by the height of outlet C, thereby assuring sufficient liquid in aperture B to maintain satisfactory electrical conductance between the electrodes. If the sample flow stops, as sometimes happens in plant practice, the solution in the reference electrode compartment leaks through the small opening D, until the level drops to that point. When this happens, the tips of both electrodes remain immersed in solution, but solution continuity between them is broken. Consequently, potassium chloride from the reference electrode cannot reach the sodium electrode, as it otherwise could by diffusion through the quiescent solution bridging between them. This feature of the cell provided automatic indication of down time, because with the meter used, the indication went off scale when the solution circuit opened. Furthermore, reliable measurement begins immediately when sample flow resumes, because the electrode is already adjusted to sample conditions and does not need time to dispel the effect of potassium ions. A calomel rather than a silver reference electrode is used to avoid any possibility of interference from silver ion. The electrolytic conductivity of the amine-treated sample water is relatively high. Therefore, no trouble was encountered with streaming potential effects, which probably would be encountered with this cell in measuring pH of high purity water. Liquid Junction. In sodium ion measurement, the liquid junction can be a source of serious trouble unless this item is properly designed and operated. The high pH of the sample solution, combined with flow-induced head fluctuations that result in transport variations of potassium chloride away from the junction, can make a poor junction produce electrical noise. Many of the erratic electrode flow effects in p H and specific ion measurements reported from the field are undoubtedly related to faulty liquid junctions. Bad conditions are illustrated at A , B , and C of Figure 7. These are traces of actual sodium ion chart records. As shown, a poor junction can cause excursions of the trace of 5 or 10 mV or more, occurring sporadically or in rapid succession. The seriousness of an erratic junction is illustrated by scale E, which indicates the percentage error introduced into the measurement by excursions in traces A , B, and C. Two factors were found to be important in greatly improving the operation of the liquid junction. The solution pressure difference across the junction should be much greater than that usually employed. Instead of the usual 5- or 10cm head differential, the head difference should be approximately 300 cm. Since it was inconvenient to elevate the reference solution reservoir this much, the desired condition was achieved by applying 5 lb/in.* gas pressure to the reservoir, using a cylinder of compressed inert gas, equipped with a regulator set at the desired pressure. The other important factor is the construction of the junction itself. A small porous plug of ceramic sealed into the
N
Figure 6. Cell arrangement
E
Figure 7. Liquid junction effect on output record reference electrode tube is preferred. The porosity is chosen to give a solution flow under operating pressure of about 2 ml per day. At this flow rate, a 500-ml supply of solution will last in the field more than 6 months before needing replenishment. With the junction described, very satisfactory field performance was observed over a period of many months. The trace shown at D,Figure 7, is typical of the entire period. ACKNOWLEDGMENT
Some of the information reported, particularly that of Figures 4 and 7, is based on field experience gained at Duquesne Light Company. Expressions of appreciation go to Walter A. Lower, who was in charge of the field work, and to his associates. RECEIVED August 31, 1970. Accepted December 2, 1970. Paper presented at the Twenty-first Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Cleveland, Ohio, March 5,1970.
ANALYTICAL CHEMISTRY, VOL. 43, NO. 3, MARCH 1971
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