John T.
W O C ~
University of Connecticut Storrs, Connecticut 06268
The Response of a Sodium-ion Selective Electrode in Acid and in Alkaline Medium
Ion-selective potentiometric electrodes are now important tools of analytical chemistry. The applications and diversity of these electrodes are increasing rapidly, and are the subject of several recent reviews.' An important characteristic of an electrode that is sensitive to ion i in the presence of a foreign ion j is the selectivity ratio K,,. A simple experiment to stress this point can be carried out with a Beckman sodium-ion electr~de.~Unlike many other ion-selective electrodes, this electrode is comparatively inexpensive. Although it responds rapidly to changes in sodium ion
activity ( a ~ . +this ) , electrode has K m + / a += 0.028 and is in fact more sensitive to am+than to a~.+. 1. Inta each of seven 10C-ml volumetric flasks, pipet 10 ml of a solution that iis 0.100 M in each ammonia and ammonium chloride (pH > 9). Add sufficient 1.00 M or suitably diluted sodium chloride solution to produce sodium ion concentrations of 100, 50, 10, 5, 1, 0.5, and 0.1 mM when made up to volume with water. Transfer the solutions ta marked beakers and note the temperature. 2. Insert'the sodium-ion electrode (NsE) and the saturated calomel electrode (SCE) into the 0.1 mM solution, swirl, leave for 2 to 3 min, then read the potential, E, of the NaE electrode versus SCE (E will be negative). In turn, repeat with the other six solutions in order of increasing concentration, 3. Take the ionic strength, u, as the sum of the concentrations of ammonium chloride and of sodium chloride and calculate o~.+in each of the solutions. Use the equations
and -log f m + = 0.50 uL/a/(l
+us)
Plot E against log ax.+. The plot should he linear as in the figure (open circles). 4. By successive dilution of a solution that is 0.100 M in each acetic acid and sodium acetate (pH < 5), prepme a second set of solutions that me 0.1-100 mM with respect to sodium ion. 5. Note the temperature and measure E with the NaE and SCE as before. Then replace the NaE with a normal (i.e., pH) glass electrode and measure the pH of the seven acetate solutions." 6. Take u .as the concentration of sodium acetate and calculate aN.+ as before. Plot E against log anr+(see figure, squares). Calculate aa+ in each acetate solution from pH = -log aa*,' then plot E against log (ma+ a~+/0.028). If the temperatures in the two sets of runs did not differ by more than about lac,the points (solid circles in the figure) should fit the line defined by the points applicable to the ammoniacd solutions. Read off the values of EO (versus SCE) and S in the equation E = Eo Sloga~.*.
+
+
Respenre of sodium-ion electrode in dkaline and in acid solution. Open circles, NoCl in NHs-NHLl solution; solid circler, CHsCOOH-CH&OONo rolutionr, oa+ allowed for; squares, CH3COOH-CHJCOON. solutianr, OH + ignored.
J. D. R., Lab. Pract., 18, 1056 G. J., O m , R. B., AND THOMAS, (1969); RECHNITZ,G. A,, Anal. Chem., 41,109A (1969). Beckman Instruments Inc., Fullerton, Calif. 92634, bulletin 7145, sodium-ion electrode 39278. The smdl efiect of dilution u .m n pH . illustrates an important property of a hufier solution. 'This relationship is, of coune, not exact. pH is measured by reference to standard hufiers and has no simple theoretical significance. See: MATTOCK, G., in "Advances in Analytical Chemistry and Instrumentation," (Editor: RICILLEY, C. N.), John Wiley & Sons, Ine., New York, Vol. 2,1963, pp. 3941.
Volume 47, Number 8, August 1970
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