observed difference in performance may be attributable to slight eccentric positioning in the case of Curve 111. Even when the bulb only approximately matches the cavity, however, the flow velocity situation will be better than when no insert is used. Referring to Figure 4, a comparison can be made to see what flow rate value with the insert will give a quality of measurement equivalent to that obtained a t 125 ml/min without the insert. On Curve I1 or 111, this value corresponds to a flow rate less than 10 ml/min. Thus, equivalent performance can be obtained with the insert a t a flow rate under one-tenth that without the insert. Use of the insert should therefore be an effective means of conserving sample solution when this is in short supply. Conservation of added base reagent is also an important benefit that can be achieved. O t h e r Ion Measurements. I t should be possible to apply the flow velocity principle to achieve benefit in the case of certain other ion-selective electrodes. Solid-state and liquid membrane electrodes operating as electrodes of the first type should be amenable to sensitivity improvements. In these electrodes, the same ion as that to be measured, is bound within the structure of the sensing element. Normal solubility mechanisms will release some of this ion to the sample solution, thereby establishing under ordinary circumstances a lower limit to sensitivity. By flowing fresh sample solution rapidly past the electrode, concentration increase from solubility can be minimized. For example, the lanthanum fluoride electrode in measuring unbuffered fluoride ion a t low levels is a likely candidate. Lingane (12) has shown in studying performance of the lanthanum fluoride electrode in batch measurements that stirring of the solution extends the lower limit of response to fluoride ion, below the theoretical solubility limit, by minimizing accumulation of fluoride ion a t the sensor surface. Srinivasan and Rechnitz (13) have observed a similar effect. A stirring effect of this kind will depend, of course, on the sample solution not having reached solubility equilibrium with the sensor material. In the case of continuous measurements, fresh sample solution is always striking the sensor, so that maximum benefit should be realized. The flow velocity technique may not prove useful for increasing measurement sensitivity in those cases where the mechanism of electrode response depends on interposed
solubility product equilibration, as in measurement of lanthanum ion with the lanthanum fluoride electrode. O t h e r Flow Effects. A word of caution should be expressed in applying the high flow velocity principle described in this paper. Electrode systems, if not properly constructed, can produce spurious voltages induced by sample flow. Trouble can arise a t the liquid junction. Also, in sodium ion measurement, the design of the flow passageway is important to prevent reference electrode solution from reaching the sensor bulb through turbulence. Recommendations have been given for coping with these problems (6). In sodium ion measurements, the high conductivity of sample water resulting from addition of base should safeguard the system from streaming potential effects. In other measurements, however, where sample conductivity is low, as in pH measurement of nearly pure water, streaming potential effects can be serious and should be dealt with in accordance with good practice (14, 15). If the system is working properly, flow variations should have virtually no effect on the voltage reading, as in Curve I11 of Figure 4. ACKNOWLEDGMENT Appreciation goes to R. D. Gillen who reviewed the literature and manuscript and offered helpful comments. LITERATURE C I T E D (1)W. 6.Gurney, Elect. Work$ March 23, 1964,p 125. (2) D. Hawthorn and N. J. Ray, Analyst(London), 93, 158 (1968). (3)H. M. Webber and A. L. Wilson, Anawst (London),94, 209 (1969). (4)E. L. Eckfeldt. W. E. Proctor, Jr., W. D. Howie, and W. A. Lower, Proc. 29th Inter. Water Conf., Engineers Society of Western Pennsylvania, Pittsburgh, November 19-21, 1968,p 109. (5) A. A. Diggens. K. Parker, and M. H. Webber, Ana/yst(London), 97, 198 (1972). (6) E. L. Eckfeldt and W. E. Proctor, Jr.. Anal. Chem., 43,332 (1971). (7)E. L. Eckfeldt. Anal. Chem., 47, 2309 (1975). (8)R. G. Bates, "Determination of pH. Theory and Practice", Wiley, New York, 1964,pp 296-302. (9)G.A. Perley. Anal. Chem., 21, 559 (1949). (10)D. Hubbard, J. Res. Natl. Bur. Std., 36,365 (1946). (11) E. L. Eckfeklt, /SA Trans., 9,37 (1970). (12)J. J. Lingane. Anal. Chem., 40,935 (1968). (13) K. Srinivasan and G. A. Rechnitz, Anal. Chem, 4kW (1968). (14) "pH of Water and Waste Water", ArneiiEan Society for Testlng and&&-terials, Philadelphia, Designation D 1293-65(1965). (15) P. Van den Winkel, J. Mertens. and D. L. Massart, Anal. Chem., 46 1765 (1974).
RECEIVEDfor review May 30, 1975. Accepted August 5 , 1975.
Unusual Temperature Effect in Low Level Sodium Ion Measurement with the Glass Electrode Edgar L. Eckfeldt' Corporate Research Department, Lee& & Northrup Company, North Wales, Pa.
Several publications discuss the implementation and ' utility of industrial sodium ion measurements (1-5). Requirements for making successful measurements in the vicinity of 25 "C at levels down to 0.1 ppb have been described (6, 7). In the course of a recent study investigating conventional temperature compensation of a sodium ion glass electrode system (L&N 7971-1 sodium ion analyzer and 7070-05 industrial sodium ion monitor), an abnormally large temperature effect was noted when a sample solution Present address, 6 Lindenwold Terrace, Ambler, P a . 19002.
of low sodium ion concentration (about 1 ppb) was heated to 60 "C. The cause was traced to erroneous response of the electrode to hydrogen ion. The present note explains the effect and issues a word of caution. Moreover, advantage may be taken of the effect to significantly improve the sensitivity of the measurement and/or to save on the amount of base reagent added to the sample solution. The lower in concentration one wishes to go in measuring sodium ion, the more stringently one must repress the interferences coming from hydrogen ion and from the cation produced by the base that is added to suppress hydrogen
ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975
2309
Table I. Values of the Ionization Constant of Water at Several Temperatures Temperature, "C
-log K,
0 25 50 60 75 100
14.944 13.997 13.262 13.017 12.691 12.258
KJX
ioL5)
1.1376 10.069 54.70 96.16 203.5 552.5
ion. A relatively strong base reagent is needed, and it should be chosen so that the electrode will be insensitive to the cations it produces (6). The reaction of base with water produces hydroxyl ions, which operate through the K , equilibrium to suppress hydrogen ion activity: aH+ =
KW -
(1)
aOH-
When hydroxyl ion activity is made large in value, hydrogen ion activity, of course, becomes small. The effect of temperature was measured on a 0.02M solution of diisopropylamine, a preferred base reagent for sodium ion measurement. From 0 to 60 "C, the pH change with temperature is nearly linear. In going from 25 to 60 "C, pH decreases about one pH unit, which corresponds to a tenfold increase in hydrogen ion activity. In going from 25 to 0 "C, pH increases by almost 0.7 pH unit. If suppression of hydrogen ion activity is only marginally adequate a t 25 "C, a rise in temperature will upset the situation. Instead of following sodium ion, the electrode will respond to the increase in hydrogen ion activity and give an erroneous upscale indication. The hydrogen ion activity of almost any base solution is strongly temperature dependent, giving a sharp rise in hydrogen ion activity as temperature rises (8). Equation 1 shows that hydrogen ion activity depends on the values of both hydroxyl ion activity and K,. Although temperature usually has some effect on the value of the ionization constant of the base, the hydroxyl ion activity of many base so-
lutions does not change drastically with temperature. The main cause of pH change is the change in value of K,. Values of K , at various temperatures are listed in Table I (9, IO). About a ninefold increase in value of K , occurs in going from 0 to 25 "C, and another ninefold increase takes place from 25 to 60 "C. The large observed change in pH of the diisopropylamine solution is thus largely attributable to change in K , value over the temperature range. It will be advisable to make low level sodium ion measurements at low rather than at high temperature. One should avoid heating the sample solution. Although use has not been made of the principle, it would appear desirable to cool the sample solution to the vicinity of 0 "C if one is striving to make the lowest possible concentration measurement of sodium ion. This way of suppressing hydrogen ion activity is achieved without corresponding increase in base cation concentration, which is an advantageous objective in its own right. Conversely, to maintain constant hydrogen ion activity, lowering of temperature offers a way of conserving on the use of base reagent, since less base will be required at lower temperature to hold the same hydrogen ion activity as that needed a t higher temperature. Another possible advantage of operating at low rather than high temperature is the probable lesser interference at low temperature from alkali metal ion leaching from the electrode (7). LITERATURE CITED (1)W. B. Gurney, Nect. World, March 23,1964,p 125. (2)D. Hawthorn and N. J. Ray, Analyst(London),93, 158 (1968). (3)H. M. Webber and A. L. Wilson, Analyst (London), 94,209 (1969). (4)E. L. Eckfeldt, W. E. Proctor, Jr., W. D. Howie, and W. A. Lower, Proc. 29th Inter. Water Conf., Engineers Society of Western Pennsylvania, Pittsburgh, November 19-21,1968,p 109. (5) A . A. Diggens, K. Parker, and H. M. Webber, Analyst(London),97, 198 (1972). (6)E. L. Eckfeldt and W. E. Proctor, Jr., Anal. Cbem., 43, 332 (1971). (7)E. L. Eckfeldt and W. E. Proctor, Jr., Anal. Cbem., 47,2307 (1975). (8)E. L. Eckfeldt, /SA Trans., 9, 45 (1970). (9)H. S. Harned and R. A. Robinson, Trans. Faraday SOC.,36, 973 (1940). (IO)R. A. Robinson and R . H. Stokes, "Electrolytic Solutions", 2nd ed., Academic Press, New York, N.Y.. 1959,p 544.
RECEIVEDfor review May 30, 1975. Accepted August 5 , 1975.
Ion-Selective Electrode Determination of Complex Formation Constants in Sub-Micromolar Silver Nitrate Solution Arthur L. Cummings and Keith P. Anderson Department of Chemistry, Brigham Young University, Provo, Utah 84602
The Nernstian response of silver ion sensitive membrane or electrodes in solutions containing between 10-l and molar silver ion has been well established (1-5). Furthermore, the Nernstian electrode response has been shown to extend to levels of free silver ion below molar (2, 6 ) in solutions wherein most of the silver present is bound by strong complexing agents ("buffered" solutions). In unbuffered solutions, deviations from Nernstian response have or been reported when silver ion is less than molar (3, 5 ) . This paper reports the observation of deviations from Nernstian electrode response in solutions in which most of the silver is strongly complexed, but the total formal concentration of silver is less than molar. A method is presented by which a correction for the non2310
Nernstian response can be determined as a function of formal silver concentration. The method is applied to the potentiometric determination of stability constants of bromoargentate complexes. EXPERIMENTAL Reagent grade sodium bromide and silver nitrate were dried at 110 "C prior to the preparation of solutions. Doubly distilled or distilled-deionized water having a specific conductance of less than 1.5 X mho/cm a t 25 O C was used. Apparatus. Solutions were prepared and dispensed using a Mettler P1200 or a Mettler P160 top-loading, constant-load balance. Polyethylene gravimetric burets (7) were used to dispense the solutions. Potential differences between an Orion Sulfide Ion Activity Electrode, Model 94-16A, and an Orion Double-Junction
ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975
