Because ISEs respond to the activity of an analyte, and not just its concentration, it is important to control (or account for) the ionic strength of the sample and standard solu- tions. When one has the luxury of working with discrete samples, adding a high level of electrolyte is an attractive way of fixing the ionic strength and, subsequently, the activity coefficient of the analyte. Adding a pH buffer at the same time also helps to control for complexation, precipitation (or protonation, if the analyte is the anion of a weak acid).
In special cases, complexing agents are purposely included in the buffer in order to bind specific metal ions that might interfere with the analysis. For example, iron(III) and alu- minum(III) bind fluoride ions in solution leading to an erroneously low signal for F−when present. Citrate is often added to samples to prevent that problem because citrate competes with fluoride for binding to these metal ions. These additives are often combined into a
“total ionic strength adjustment buffer” (TISAB). These ingredients can be mixed in a solu- tion in advance and then added to all samples and standards in a constant volume ratio immediately before measuring the signal. This approach assumes that the ionic strength of the buffer is so much greater than that of the sample that the ionic strength and pH of the mixture is set by the TISAB. A further bonus of this approach is that the calibration curve can be plotted in terms of the concentration of the analyte rather than the activity.
There is no need to evaluate the activity coefficient of the analyte. Table 3.2 gives example recipes for ionic strength adjustment buffers that were originally formulated for use with
TABLE 3.2 Ionic strength adjustment buffers
TISAB I TISAB II TISAB III
For use when magnesium, calcium, chloride, nitrate, sulfate, or phosphate are present at high levels
For use when iron(III) or silicate are present at high levels
For use when aluminum is present at high levels
500 ml deionized water 500 ml deionized water 500 ml deionized water 57 ml glacial acetic acid 170 g sodium nitrate 17.65 g cyclohexanedi- aminetetraacetic acid (CDTA)
58 g sodium chloride 68 g sodium acetate trihydrate Add 40% NaOH, drop by drop until the salt is dissolved 0.3 g sodium citrate
dihydrate
92.4 g sodium citrate dihydrate
300 g sodium citrate dihydrate 60 g sodium chloride
Dilute to 1 l Dilute to 1 l Dilute to 1 l
Source: Reproduced with permission from Buck and Cosofret [33]. Copyright 1993, IUPAC.
k k fluoride ISEs. An addition of 10 ml TISAB I per 100 ml of sample gives an ionic strength of
about 0.1 M, a good target value.
3.9.2. Potential Drift
In principle, ISE membranes need only a second or less to reach an equilibrium with a new sample, if the system is well stirred. In practice, it is a common experience to have a system take many seconds or even minutes to reach a steady potential. Changes in potential in the same direction that occur over minutes despite constant analyte conditions are usually called “potential drift.” This same term is also applied when repeated measurements with the same standard or sample solution are not reproducible, but trend in one direction. The source of the drift is often attributed to problems with a reference salt bridge [24].
Trustworthy results (as evident from highly linear calibration curves) can usually be obtained by waiting to record a potential until the signal has reached some predetermined rate of change such as 0.2–0.4 mV/min. Alternatively, good analytical data can also be obtained by recording potentials after a specific time interval from the time at which the sensor was first immersed in the sample solution [41].
PROBLEMS
3.1 A calcium ion ISE gives a response of +170 mV in a 100.00 ml sample solution.
The same solution is spiked with 5.00 ml of 0.125 M calcium standard solution.
After thorough mixing the same electrode combination gave a reading of+193 mV.
Assuming a Nernstian response and no interferences, calculate the concentration of Ca2+in the original sample solution.
3.2 A technician puts a new membrane on a nitrate ISE and records the following data for a series of calibration standards in an appropriate ionic strength buffer.
[NO3−] E (mV)
5.50×10−7 −133
6.00×10−7 −133
1.19×10−6 −134
2.50×10−6 −141
6.00×10−6 −150
2.38×10−5 −183
1.18×10−4 −224
1.18E×10−3 −282
(a) Plot the appropriate form of the data for a calibration curve.
(b) Calculate the slope for the electrode response in the linear range.
(c) What is the apparent detection limit for this electrode?
3.3 A nitrate ISE with a new membrane gave the following response to nitrate stan- dard solutions prepared in an appropriate ionic strength buffer. Each solution also
k k contained 2.00×10−2M ClO4−. Graph the data and calculate the apparent selectivity
coefficient for the interfering ion.
[NO3−] E (mV)
5.50×10−7 −140
6.00×10−7 −140
1.19×10−6 −143
2.50×10−6 −146
6.00×10−6 −160
2.38×10−5 −183
1.18×10−4 −224
1.18×10−3 −282
3.4 A colleague brings you her Pb2+ ISE based on a liquid membrane that uses the ionophore, ETH 5435. She said that it exhibited a slow, continuous voltage drift (∼5 mV/min) when she was measuring trace lead levels on pond sediment samples in the field. You begin to speculate on possible causes and quickly list the follow- ing ideas as they come to mind. Describe the type of behavior that each of these issues would cause. Which, if any, might give rise to the behavior that your colleague described.
(a) The membrane may have been punctured, perhaps by something in the sedi- ment.
(b) The silver ions from the reference electrode may have reacted with something in solution to precipitate in the salt bridge.
(c) Your colleague may have forgotten to soak the electrode a standard solution of Pb2+after putting a new membrane on the device.
(d) The temperature may have been changing during the measurement.
3.5 How many Ca2+ions cross the membrane/solution boundary corresponding to a 0.050 V increase in potential for a sensor with a 0.090 cm2 area assuming a capaci- tance of 3×10−6F/cm2?
3.6 The following data were obtained by preparing each standard solution by diluting the appropriate volume of a 1.00×10−2M Cl−stock solution into 250 ml volumetric flasks and diluting to the mark with deionized water. Before measuring, 100.0 ml of each standard was spiked with 1.00 ml of 4 M NaNO3to adjust the ionic strength.
[Chloride] E (mV)
1.00×10−3 −113.0 1.00×10−4 −54.0
1.00×10−5 5.0
5.00×10−6 22.8
1.00×10−6 64.0
5.00×10−7 81.8
k k (a) What concentration of chloride ion would be indicated for a sample that was
also treated with ionic strength buffer and gave a signal of 15.5 mV?
(b) If the sample had not been treated with NaNO3to adjust the ionic strength before measuring, what would the potential reading have been (neglecting the influ- ence on the junction potential for the reference bridge)?
3.7 A cell using a potassium ISE based on a liquid membrane using valinomycin as an ionophore and a silver/silver chloride electrode gives a voltage of−0.273 V in a solution of 7.50×10−6M KCl standard solution (using an ionic strength buffer).
Assuming that the electrode responds ideally, calculate the cell potential for the same conditions plus the presence of 2.00×10−2M HCl as an interfering electrolyte.
3.8 What is the maximum activity of sodium ions that can be tolerated in order to mea- sure K+ ions at activities as low as 1×10−4M with at most a 10% error using a valinomycin liquid membrane?
3.9 If the liquid junction potential for a combination pH electrode drifts by 5 mV, what is the relative error in the apparent H+concentration?
3.10 Consider using an ISE for determining bromide ion based on AgBr/Ag2S crystalline membrane. TheKspfor AgBr is 7.7×10−13. What is the approximate detection limit for bromide ion in the absence of interfering ions? What would the bromide detec- tion limit be in the presence of 0.01 M KCl, given thatKPOTBrCl=3×10−3?
3.11 Using the Henderson equation and table of diffusion coefficients in Appendix C, calculate the junction potential for a salt bridge with 3 M NaCl solution contacting a sample solution with 0.2 M HNO3.
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