Potentiometric Determination of an Overall Formation Constant Using an lon-Selective Membrane Electrode A Laboratory Exercise for Analytical Chemistry Niegomir Radid and Josipa Komljenovid Faculty of Technology, University of Split, 58000 Split, Croatia
The development and application of ion-selective membrane electrodes (ISME'S)continues to be a rapidly growing area of analytical chemistry. However, in the standard analytical chemistry textbook (13)little attention is given to this subject, and there is a gap between the available teaching sources and the research literature.
where k' denotes a constant part of the term in buffered solution.
A Study of an AluminumFluoride Complex in a WaterlAcetonItrile Mixture
This experiment introduces the student to a potentiometric study of aluminum-fluoride complex formation in waterlorganic solvent mixtures using a fluoride iou-selective electrode (FISE) with an LaF3 membrane. This is used to determine free fluoride ion in solution. The reaction studied is shown below.
Its overall formation constant pi is given by
In this experiment, p; in waterlacetonitrile mixtures is determined. The Titration
During a potentiometric titration, the concentration of aluminum species in a slightly acidic media can be represented by the following equation.
The exact pH value of the reaction solution is measured with a glass electrode, and the value is used to calculate k'. The Potential of the Fluorine ISE
The ~otentialof the FISE as a function of "free" fluoride conce&ation at constant pH and ionic strength is given by an empirical form of the Nernst equation. The slope S and the constant K can be calculated from experimental data. Combining eqs 5 and 4 we get the potential of the electrode before the equivalence point in the following form.
where c~is the total concentration of fluoride in solution. When
of the volume at the equivalence point is added, the measured potential (E') is related to the overall formation constant of the aluminum-fluoride complex.
The Stoichiometry between Aluminum and Fluorine
where PT (i = 1, 2, 3) denotes the stability constant of AI(0H)j. The logs of the corespondi~gconstants (4)for
are
As shown by experiments in mixtures withother organic solvents (5.61,the stoichiometnc ratio between alum~num and fluoride in their complexes depends on the aluminum concentration. tA1F6I3-is the complex in solution for the concentration of a l & n u m chosen in this experiment, that is, (5 x 104 - 5 x lo9) M. The quantity of aluminum in the reaction solution can be calculated from the 1:6 ratio according to the following equation.
log p; = -4.3 log p; = -9.3 log p i = -15
where Vq is the added volume of titrant at the equivalence point. Equation 7 then becomes
Combining eqs 1and 3, we get Volume 70 Number 6 June 1993
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Representative Data for the Calculpon Overall Formation Constant of the [AIFs] Complex in WaterIAcetonitrile Mixtures
coated magnetic bar in a double-walled cell maintained at 25 'C. Measurement and Calibration
I~
1
0
~ o.iooo
54
28.2
8 ~ 1 0 ~
0.1000
54
28.4
5x10~
0.1000
54
28.8
'At the begining of the titration.
Experimental Chemicals All chemical reagents used in this experiment were of analytical grade and were used without further purification. Solutions were prepared with doubly distilled water. Solutions Standard Sodium Flwride: A 0.5 M solution, c = 0.1, was prepared from sodium fluoride that had been dried for several hours at 125 'C. It was made up using a polypropylene calibrated flask. AluminumOZI) Standard Solution: Standard aluminum solution, c = 0.5 gL, was prepared by weighing and dissolving the appropriate amount of aluminum wire (99.99% aluminum) in 10 mL of hydrochloric acid (sp gr: 1.18). The solution was then transferred into a 500.0-mL calibrated flask. It was diluted to the mark with water and mixed. Other solutions of aluminum were prepared fmm the standard solution by adding hydrochloric acid and water. Buffer Solution: Acetate buffer, pH 5, was made by diluting glacial acetic acid (60.0 mL) and sodium acetate (16.30 g) to 1,000.0 mL. Equipment
Pipet 5.0 mL of acetate buffer and between 2.0 and 10.0 mL of the chosen aluminum solution h t o the thermostated reaction cell. In this way various amounts of aluminum, but not less than 10 p a l , can be transferred to the cell. Add water and acetonitrile to bring the volume to the same total volume (37.0 mL) for each experiment, making the vol% of acetonitrile in the range from 54% to 68%. Start the stirrer, and add the standard fluoride solution with a delivery speed of about 0.1 mumin. Calculations and Graphing Plot the experimental potential values versus the volume of titrant. From these data, calculate the volume of titrant added at the equivalence point. Take the pH value of the reaction solution. Calculate k', and estimate E'. Calculate the fluoride concentration for each potential value after the equivalence point, and construct the plot of E versus pF (calibration e w e ) . Estimate the S and K values. From the collected data, calculate the formation constant of [A&13- complex using eq 9. Results and Discussion A titration curve of aluminum with fluoride in water/ acetonitrile mixture buffered with acetate buffer is shown in Figure 1. The concentration of aluminum in solution (cAl) is determined using eq 8. Using the procedure described above and the exact pH value of the reaction solution (pH = 4.08), we estimated the potential value (Fig. 1). E' = 72 mV
A plot of the electrode potential versus -log [Fh obtained with the experimental data after the equivalence point is shown in Figure 2. The plot is linear with slope S = 56 mV @FT' The constant of the electrode is
The emf measurements of an FISE (Orion 94-09) in combination with a single-junction reference electrcde (Orion Model 90-02) were made with a Mimprocessor Ionalyser (Orion 901). The pH values were checked with a glass eleetrode and 701 pWmV meter (Orion 91-02). During the measurements the solution was stirred with a PTFE-
0.8
Representative values of the overall formation constant of the [AlFcl" com~lexin waterlacetonitrile media are shown in the table. '
1.6 1
"E. p. "(0.1 M N~F)'"'~ Figure 1. Potentiometric titration curve of aluminium(lll)with fluoride in water + acetonitrile mixture, pH was 4.08. 510
Journal of Chemical Education
Figure 2. Calibration curve calculated with the experimental data recorded after equivalence point. '.
Literature Cited
3.Wilard.H. H.;Mpmtt,L. L.;Dean,J.A.;Settle,F.A.Instr.mnfolMefho&ofAnolyxis. 6th d.; D. van Nmkand: New Yorh.1981; p 640.
1. Skwg, D.A.: West,D. M.PrinerjksafInstnrmolAnolysis. Znded.: Saunders:Philadelphis, 1980: p 396.
4, Kragte,,, J,hlyst
1. Pet-, D. G.;Hayea, J. M.; HieRlje,G.M.Chamid &porntian ondM-unmnts: Saundem: Philadelphia, 1974; p 367.
1914,99,
-,
6. Baumann, E . WAnaL C h . 1970,42,110-111.
6. Radih, Nj.:P-,
D.:B ~ ~ IM. ~C J.,E l e t m m l . Chem. 1988,248,87-90
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