In the electrode-calibration method, K is determined by measuring Ecell for one or more standard solution of known pM or pA. The calibration is performed just before the determination of pM or pA for the unknown. There is an assumption that K is unchanged when standard is replaced by the unknown analyte solution. Some time instead of calculating K using emf value of standard solution, a calibration graph is plotted using measured emf of standard solutions against log a (concentration), pM or pA. The sample is then treated in the same way as the standards and the concentration (or activity) read directly off the calibration graph. For getting fast results, a graph between cell potential and log a is plotted and for this we can use graph paper called semi-log paper for drawing calibration curve and concentration of a sample can be read off directly from the graph.
3.3 Basic Theory
An ion selective electrode generates a difference in electrical potential between itself and a reference electrode. The output potential is proportional to the amount or concentration of the selected ion in solution.
The concentration is a measure of the number of ions in a specific volume. The definition assumes that all of those ions behave in the same manner. However, ions do not always behave similar to one another: some are effective i.e. exhibit properties associated with that ion, and some are not effective. The number of effective ions is called the activity of the solution. It is therefore reasonable to assume that the electrode will measure the activity rather than the finite concentration of the ions. In dilute solutions though, the ionic activity and concentration are practically identical but in solutions containing many ions, activity and concentration may differ. This is why dilute samples are preferred for measurement with ISE's.
It is possible to 'fix' the solution so that activity and concentration are equal. This can be done by adding a constant concentration of an inert electrolyte to the solutions under test. This is called an Total Ionic Strength Adjustment Buffer (T.I.S.A.B.). Thus the ion selective electrode will measure concentration directly. Activity can also be an important quantity to measure; for instance, it is the activity of calcium in blood that is physiologically important, and not the concentration.
The measured electrode potential, E, is related to the activity of an ionic species by the Nernst equation.
Where Eo = a constant for a given cell
R = the gas constant
T = the Temperature in Kelvin
n = the ionic charge
F = the Faraday constant
and the expression RT/nF is termed the Slope Factor
For example, when measuring Potassium ions, (i.e. n = +1), the slope factor at 298K (25°C) has a value of 59.16 mV. This is termed the Ideal Slope Factor, and means that for each tenfold change in Potassium concentration, an ideal measuring system will sense a mV change of 59.16.
The measurement of slope factor gives an indication of the performance of the electrode system.
If ion selective electrodes are not cleaned after use, and are subject to long term neglect, then the accuracy of the system is lost. This loss of performance can be monitored by a steady decrease in measured slope value during the calibration of a system.
A number of factors including reference junction blockage, electrolyte loss, electrode interference and the use of incorrect calibration solutions will all contribute to ‘low slope values’. All of these must be considered when there are doubts about the system performance.
Direct measurements are particularly useful for ions which have specifically designed electrodes. However, the accuracy can be increased by using different methods of measurement, e.g. standard addition, known subtraction or titration. It is possible to use the ion selective electrode as an end point indicator, just as litmus or universal indicator can be used as indicators for acid/base titrations.
Furthermore, it is even possible to measure the concentration of an ion for which there is no specific electrode, e.g. the measurement of aluminium concentration where there is no aluminium electrode. By carefully considering the chemistry of ions such as aluminium one can create a system by which it is possible to determine their concentrations. For example, aluminium fluoride is insoluble, therefore the addition of a sodium fluoride solution precipitates the aluminium out of solution as aluminium fluoride. Using a fluoride ISE, the concentration of sodium fluoride added to aluminium solution can be measured. At first there will be no potential as the fluoride precipitates with aluminium. When all the aluminium has reacted further additions of fluoride will provide a sudden change in electrode potential. By drawing a graph of electrode potential versus amount of sodium fluoride added, the concentration of sodium fluoride required to react with the aluminium can be calculated. Therefore, if one knows the ratio in which the aluminium and fluoride react, the aluminium concentration can be found.
3.4 Selectivity, Interferences, activity
As the name suggests, ion selective electrodes are selective to one ion but not specific for it.
This means that other ions in solution may also be sensed by an ISE although it is not designed to do so.
The ion that is to be determined is referred to as the determinant and other ions to which an electrode responds are known as Interferents or interfering ions. It is possible to calculate the preference of an electrode for the determinant over the interferent. This is called the selectivity of the electrode. The preference, expressed as a ratio is called the selectivity coefficient, or ratio. Each electrode has its own set of selectivity coefficients. For example:
Meaning that the preference for K+ (potassium) over Na+ (Sodium) for this electrode is 1 to 2.6 x 10-3 or 385:1. This means that the electrode is 385 times more selective to K+ than Na+.
These coefficients are not always constant and the manufacturer's specification should be consulted.
The ‘pH glass-electrode’ or hydrogen selective electrode is the most responsive of all ISE's, yet measurements of pH to better than ±0.01 pH are known to require considerable care: this corresponds to an uncertainty of ±2%. However, accuracy of 2% or even down to 0.5% can be achieved if the operator follows good laboratory procedures.
It is possible to increase accuracy by using different measuring techniques. Methods known as known addition, sample addition, etc., have been devised to improve accuracy. These techniques reduce the effect on the result of errors in single readings and consequently a balance must be found between accuracy and convenience in the choice of technique.
3.5 Types of Electrode
There are four types of ion selective electrode whose construction and mode of operation differ considerably. These are:
1. Glass body electrode
2. Solid state (crystalline membrane)
3. Liquid ion exchange (polymer membrane)
4. Gas sensing type
1. Glass body electrodes
The most common ISE is the glass-bodied pH combination electrode. The sodium (Na+) combination has a similar construction which houses a glass bulb that is sensitive to sodium ions in solution.
2. Solid state ion selective electrode
The electrode potential of standard and sample solutions is measured across a solid, polished crystalline membrane. The crystalline material is prepared from a single compound or a homogeneous mixture of compounds (for example, the fluoride ISE has a Lanthanum Fluoride crystal).
3. Polymer membrane ion selective electrode
These electrodes use a replaceable membrane cap which has a solid, polymeric membrane containing a selective ion exchanger. The electrode potential of solutions is measured by their effect on the ion exchange material. Due to the complex properties of the ion exchangers, they are subject to more interferences than other ion selective electrodes.
4. Gas sensing type
Electrodes, including the ammonia ISE, use a gas sensing mode of operation. In the case of ammonia, a caustic solution is added to the sample solution to liberate ammonia. The gas permeates through a membrane and changes the pH of the filling solution. The change in pH is proportional to the ammonia concentration. This gives a quantitative measurement of the ammonia in the sample solution.
Figure 1: Types of electrode with diagram (Wiley 2010).
Reference electrodes
The potential of an Ion Selective Electrode can only be measured against a suitable reference electrode in contact with the same test solution. Reference electrodes are electrochemical half cells whose potentials are maintained at a constant value by the chemical equilibria maintained inside them.
Figure 2: Internal structure of Ion selective electrode (Wiley 2010).
It is preferable to use a double-junction reference electrode for ISE applications. Standard reference half cells have KCl based electrolyte filling solutions. This is a distinct disadvantage when, for example, potassium or chloride is being measured. To overcome this, a double junction reference is used in which the escaping KCI is retained in a second chamber containing a non-interfering electrolyte, which in turn escapes into the test solution at the second junction.
3.6 Ions determined by ion selective electrode
The table lists ions selective electrodes which are readily available, concentration ranges, selectivity ratios to the major interfering ions and recommended outer bridge solutions for double junction reference electrodes.
Figure 3: Possible interferences with reference solutions stated (Stanford 2008).
3.7 Solid State Electrodes
The most successful example is the fluoride electrode. The membrane consist of a single crystal of lanthanum fluoride doped with some europium(II) to increase the conductivity of the crystal. Lanthanum fluoride is very insoluble, and this electrode exhibits Nerstian response to fluoride down to 10-6M (19ppb). This electrode has at least a 1000-fold selectivity for fluoride ion over chloride, bromide, iodide, nitrate, sulphate,monohydrogen phosphate, and bicarbonate anions and a 10-fold selectivity over hydroxide ion. Hydroxide ion appears to be the only serious interference. The pH range is limited by the formation of hydrofluoride acid at the acid end and by hydroxide ion response at the alkaline end; a pH range of 4 to 9 is claimed.
A useful solution for minimizing interferences with the fluoride electrode consist of an acetate buffer at pH 5.0 to 5.5, 1M NaCl, and cyclohexylenedinitrilo tetraacetic acid (CDTA). This solution is commercially available as TISAB (total ionic-strength buffer). A 1:1 dilution of samples and standards with the solution provides a high ionic strength background, swamping out moderate variation in ionic strength between solutions. This keeps both the junction potential and the activity coefficient of fluoride ion constant from solution to solution. The buffer provides a pH at which appreciable HF formation is avoided and hydroxide response is not present. CDTA is a chelating agent, similar to EDTA, that complexes with polyvalent cation such as Al3+, Fe3+, and Si4+, which would otherwise complex with F- and change fluoride acitivity.
3.8 Ion Activity Level
The measured potential corresponding to the level of nitrate ion in solution is described by the Nernst equation.
E = Eo + S * log (A)
E = measured electrode potential
Eo = reference potential (a constant)
A = nitrate ion activity level in solution
S = electrode slope (about -57 mV per decade)
S = (2.3 R T) / nF
R and F are constants, T = temperature in degrees K and
n = ionic charge
The level of nitrate ions, A, is the activity or “effective concentration” of free nitrate ions in solution. The nitrate ion activity is related to free nitrate ion concentration, Cf, by the activity coefficient, y.
A=(y)(Cf)
Ionic activity coefficients are variable and largely depend on total ionic strength. The ionic strength of a solution is determined by all of the ions present. It is calculated by multiplying the concentration of each individual ion by the square of its charge, adding all these values up and then dividing by two.
Ionic strength = 1/2 Σ (Ci Zi2)
Ci = concentration of ion i
Zi = charge of ion i
Σ symbolizes the sum of all the types of ions in solutions
If background ionic strength is high and constant relative to the sensed ion concentration, the activity coefficient is constant and activity is directly proportional to concentration. Ionic strength adjustor (ISA) is added to all nitrate standards and samples so that the background ionic strength is high and constant relative to variable concentrations of nitrate. For nitrate, the recommended ISA is (NH4)2SO4. Nitrate interference suppression solution (NISS), a specific solution for removal of nitrate-interfering ions, is recommended for samples with competing ions. Other solutions can be used as long as they do not contain ions that would interfere with the electrode response to nitrate. If samples have a high ionic strength (above 0.1 M), standards should be prepared with a composition similar to the samples. Reference electrode conditions must also be considered. Liquid junction potentials arise any time when two solutions of different composition are brought into contact. The potential results from the interdiffusion of ions in the two solutions. Since ions diffuse at different rates, the electrode charge will be carried unequally across the solution boundary resulting in a potential difference between the two solutions. In making electrode measurements, it is important that this potential is the same when the reference is in the standardizing solution as well as in the same solution; otherwise, the change in liquid junction potential will appear as an error in the measured specific ion electrode potential.
The most important variable that analysts have under their control is the composition of the liquid junction filling solution. The filling solution should be equitransferent. That is, the speed with which the positive and negative ions in the filling solution diffuse into the sample should be nearly as equal as possible. If the rate at which positive and negative charge is carried into the sample solution is equal, then no junction potential can result. Optimum Results filling solutions are specifically design
3.9 Calibration Theory
Calibration is carried out by immersing the electrodes in a series of solutions of known concentration and plotting a graph of the mV reading versus the log of the activity (or the actual activity on a logarithmic X-axis). This should give a straight line over the whole linear concentration range. However, as noted above, activity is difficult to determine in complex solutions and it is generally more useful to plot concentration units. In this case the effect of variable activity coefficients in solutions with high ionic strength can be minimized by adding Ionic Strength Adjustment Buffer to all standards and samples - but note the limitations to this detailed in the previous chapter. Nevertheless, it must be noted that if the samples to be measured are likely to have a total ionic strength of less than about 0.01M for monovalent ions (0.001M for divalent ions) then the activity effect should be insignificant and it may not be necessary to add ISAB. However, it must be noted that ISAB may be useful when using a double junction reference electrode with a non-equi-transferrent outer filling solution, in order to compensate for drift in the liquid junction potential, and in general most ISE systems give a stable reading more quickly in high ionic strength solutions.
The slope of the calibration graph is the mV response per decade of concentration change. This is typically around 54 mV/decade for monovalent ions and 27 for divalent ions and will have a negative value for negative ions - i.e. a higher concentration means more negative ions in solution and therefore a lower voltage.
a) Linear Range.
The linear range of the electrode is defined as that part of the calibration curve through which a linear regression would demonstrate that the data points do not deviate from linearity by more than 2 mV. For many electrodes this range can extend from about 0.1 Molar down to 10-6 or even 10-7 Molar.
b) Total Measuring Range.
The total measuring range includes the linear part of the graph as shown below together with a lower curved portion where the response to varying concentration becomes progressively less as the concentration reduces. Samples can be measured in this lower range but it must be noted that more closely spaced calibration points are required in order to define the curve accurately and the percentage error per mV on the calculated concentration will be progressively higher as the slope reduces.
c) Limit of Detection.
For monovalent ions, the IUPAC definition is: that concentration at which the measured potential differs from that predicted by the linear regression by more than 18 mV. The practical limit of detection can be calculated by plotting a calibration graph using several standards at the lower end of the concentration range, and below it. Say 100, 10, 1, 0.1, 0.05, 0.01 ppm - i.e. at least two to define the linear slope and two to show the position of the horizontal section below the limit of detection, where the electrode is unresponsive to concentration change. The limit of detection is then defined by the crossing point of the two straight lines drawn through these points.
Figure 4: Calibration curve of Ion selective electrode (Thermo Scientific 2008).
4. Procedure
4.1 Calculation and Preparation of Cl- and NO3- stock solutions
The mass of NaCl has been weighed out accurately to prepare 250mL of 0.1 M stock solution.
The NaCl was dissolved carefully with deionised water and diluted to the mark of 250-mL volumetric flask. The above steps were repeated for NaNO3.
4.2 Preparation of 10-2M standard solution
The 10-2M standard solutions were prepared accurately from the 0.1M NaCl and NaNO3 in 200mL volumetric flask.
4.3 Preparation of 10-3, 10-4, 10-5M standard solutions
Following, 10-3, 10-4, 10-5M standard solutions were accurately prepared from 10-2M NaCl and NaNO3 in 100-mL volumetric flask.
4.4 Measurement using ion-selective electrode
Foremost, the electrode leads were connected to the meter. Next, 50mL of 10-5M standard solution and 1mL of TISAB (total ionic strength adjustment buffer) solution were pipetted to the small plastic beaker. Following, we have placed the electrode into the beaker while the solution was stirred by a magnetic stirrer. Afterwards, the value was recorded when the reading has stabilized. Subsequently, we have rinsed and blotted the electrode before we have repeated the above procedures for concentrations 10-4, 10-3, 10-2 and the unknown solution (The sequence of repetition was kept from the dilute to the concentrated standard solutions).
5. Results and Calculation
Nitrate
Chloride
6. Discussion
6.1 Determination of the concentration of unknown solutions
From the results we have obtained from the calibration curve, we can interpolate the concentration of the unknown sample with a known electrochemical value.
6.2 Determine which pool is the victim likely to be drowned
Let me clarify the fact that the linear square regression to obtain the best fitting line is to minimize the residuals (differences between the values from standard and the line). However, the method has uncertainty in both the slope and interception value and we may combine the uncertainty and conduct a t-test to give a further insight to determine which pool is the victim likely to be drowned.
This standard deviation can be used to calculate the standard deviation of the slope and the y-intercept using the formulas.
The confidence limits can be calculated from the t-statistic for n-2 degree of freedom. Table of t-statistics are available in any statistics textbook or even our lab manual.
The confidence limit with a 99% confidence interval, we can use tn-2 = 2.576.
Nitrate Test
Chlorine Test
The above t-test is to ensure that the probability of the sample having the sample concentration is roughly at 99% confidence. Hence, the probability of rejecting hypothesis of victim was drowned in pool A, B, C, E while it is true is less than 1%. Conclusively we can state that pool D is the only possible case at 99% confidence.
6.3 Ionic Strength, Activity coefficient and Purpose of T.I.S.A.B
Ionic Strength and Activity Coefficients: Activity versus Concentration.
Ion selective electrodes measure the concentration of ions in equilibrium at the membrane surface. In dilute solutions this is directly related to the total number of ions in the solution but at higher concentrations, inter-ionic interactions between all ions in the solution (both positive and negative) tend to reduce the mobility and thus there are relatively fewer of the measured ions in the vicinity of the membrane than in the bulk solution. Thus the measured voltage is less than it would be if it reflected the total number in the solution and this causes an erroneously low estimate of the concentration in samples with a high concentration and/or a complex matrix.
Ionic Strength is a measure of the total effect of all the ions in a solution. It is the sum of the molar concentration multiplied by the square of the valency of all the ions. The effective concentration measured at the electrode head is known as the activity of the ion. In general chemical terms it is the number of ions taking place in any chemical reaction – measured in concentration units. The activity coefficient is the ratio of the activity divided by the concentration. This is a variable factor and depends on the valency and ionic radius of the measured ion and the total ionic strength of the solution.
The activity coefficient is always less than one and becomes smaller as the ionic strength increases; thus the difference between the measured activity and the actual concentration becomes higher at higher concentrations. This effect causes two main problems in ISE measurement. Firstly, when constructing a calibration graph using concentration units, the line is seen to curve away from linearity as the concentration increases (it remains straight, up to the highest concentrations if activity units are used). Thus, if concentration units are used, it is necessary to measure many more calibration points in order to define the curve more precisely and permit accurate interpolation of sample results. Secondly, it is most likely that the sample solutions will contain other ions in addition to the ion being measured and the ionic strength of the samples may be significantly higher than that of the standards. Thus there will be an incompatibility between the calibration line and the measured samples leading to errors in the interpolated results.
It is possible to calculate the activity coefficient for the primary ion in a simple pure solution where the composition and relative concentration of all the ions is known. Thus the measured activity can be converted into concentration results for simple solutions - but in most practical applications this is not possible, or very difficult and time-consuming.
The Ionic Strength (I) can be calculated from I = 0.5 x Sum (ci x Zi2)
Where c is the concentration in Moles and Z is the valency.
The Activity Coefficient (f) can then be found from:
-Log(f) = [(0.51 x Z2 x SQR(I)) / (1 + (3.29 x d x SQR(I))] - (0.1 x Z2 x I)
Where: Z = the ionic charge, I = the ionic strength of the solution, d = the ionic radius in nanometres.
Note that it is generally accepted that this formula is only accurate up to about I = 0.1 Molar. At higher ionic strength other factors come into play which make the calculation of activity coefficients virtually impossible and thus most ISEs cannot be used reliably above this concentration.
For samples with high ionic strength, there are five possible methods which can be used to avoid the error introduced by the difference between activity and concentration.
1) Bring the ionic strength to the same level in both the calibrating standard solutions and the samples by adding a suitable Total Ionic Strength Adjustment Buffer (TISAB) to both.
2) Dilute the samples to a level where the ionic strength effect is insignificant – but make sure that the detected ion is still within the linear range of the electrode.
3) For samples with complex but known matrix, make up the standards in a similar solution which does not contain the detected ion, or any which would interfere with the measurement.
4) Use the Activity Coefficient to calculate the concentration from the activity. As noted above, the activity coefficient can be calculated for simple solutions with known concentrations of all the ions, but this is not possible in many practical applications, where the samples may have a complex or unknown matrix.
5) Use the Standard Addition (or Sample Addition) Method where the voltage is measured before and after a measured small volume of standard (or sample) is added to a larger measured volume of sample (or standard) and the ionic strength is not altered significantly. This section is further discussed in the recommendations.
Total Ionic Strength Adjustment Buffers
A useful solution for minimizing interferences with the ion electrode consist of an acetate buffer at pH 5.0 to 5.5, 1M NaCl, and cyclohexylenedinitrilo tetraacetic acid (CDTA). This solution is commercially available as TISAB (total ionic-strength buffer). A 1:1 dilution of samples and standards with the solution provides a high ionic strength background, swamping out moderate variation in ionic strength between solutions. This keeps both the junction potential and the activity coefficient of ion constant from solution to solution. The buffer provides a pH at which appreciable conjugate acid formation is avoided and hydroxide response is not present. CDTA is a chelating agent, similar to EDTA, that complex with polyvalent cation such as Al3+, Fe3+, and Si4+, which would otherwise complex with anion and change anion acitivity.
The most common way of overcoming the effect of the variable ionic strength of the solutions is to attempt to make them all the same. Theoretically, this can be done by adding, equally to all standards and samples, another solution of high ionic strength which does not contain the ion to be measured, or any likely interference. These solutions are known as Total Ionic Strength Adjustment Buffers (TISAB). The idea is that they are added in sufficient quantity to completely swamp the ionic effects of the host solutions and hence give a uniform ionic strength in all samples and standards. In this case the straight-line calibration curve can be constructed using concentration units and the unknown concentrations can be read directly from the calibration graph. Note that there is no need to recalculate the standard concentrations after adding TISAB as long as all standards and samples are treated in the same way.
It must be noted, however, that most recipes for the addition of ISAB only produce an increase in ionic strength of 0.1M and thus will only be effective if the IS of the original sample is much lower than 0.1M. Furthermore, the difference between activity and concentration is relatively small at low Ionic strengths and can often be ignored in many practical applications – see table above.
Most ISE suppliers also supply TISABs. These vary in composition depending on the detected ion and occasionally contain other components which actively suppress interfering ions and/or control the pH value. In many cases the composition of these solutions are carefully guarded ‘trade secrets’ and the reagents are simply labelled as e.g. "TISAB for nitrate electrode".
6.4 Reproducibility
Sensitivity (Slope factor)
We notice from the Nerstian equation the slope factor at 298K (25°C) has a value of 59.16 mV for both chlorine and nitrate (n=-1). This is termed the Ideal Slope Factor, and means that for each tenfold change in Potassium concentration, an ideal measuring system will sense a mV change of 59.16. However, from the results above the slopes for nitrate and chloride are 27 and 45.8 respectively which translate into 46% and 77% of the ideal Slope respectively.
I have look though several literatures and here are some of the causes of the low-slope value and the possible corrective actions to be done.
Most solid-state electrode will last six months in normal laboratory use. On-line measurement might shorten operational lifetime to several months. In time, the response time will increase and the calibration slope will decrease to the point calibration is difficult and electrode replacement is required.
Therefore, I would encourage a regular calibration and check on the slope factor to ensure that the slope factor is with 102-95% of 59.16mV. Otherwise, due to the lack of sensitivity it may lead to inaccurate and a greater range of uncertainty that may arise from the ion-selective electrode.
Temperature influences
Samples and standards should be at the same temperature, since electrode potential are influenced by changes in temperature.A l °C difference in temperature results in a 2% error at the l0-3 M level. Because of the solubility equilibria on which the electrode depends, the absolute potential of the reference electrode changes slowly with temperature. The slope of the nitrate electrode, as indicated by the factor “S’ in the Nernst equation, also varies with temperature. Table 4 gives values for the “S” factor in the Nernst equation for the nitrate ion.
The operating range of the nitrate ion electrode is 0 °C - 40 °C, provided that temperature equilibrium has occurred.
If the temperature varies substantially from room temperature, equilibrium times up to one hour are recommended
Electrode Response
Plotting the electrode potential against the nitrate concentration on semi-logarithmic paper results in a straight line with a slope of about 55 mV/decade.
The time needed to reach 99% of the stable electrode potential reading, the electrode response time, varies from one minute or less in highly concentrated solutions to several minutes near the detection limit.
Therefore, we should always allow the reading to reach its equilibrium before taking its reading. We were also encouraged to take either the lowest, highest or the average of both to ensure consistency in the way we take our reading. Otherwise, it will cause discrepancy in our values and low linearity value when we plot our calibration curve or interpolate from it.
Noisy or unstable reading (reading continuously or rapidly changing) may be due to a several factors and I have also included the corrective actions that ought to be taken to ensure a stable reading.
In some cases, I have noticed that the value either constantly increases or decrease instead of fluctuation and reaching equilibrium and this phenomenon is known as drift (reading slowly changing in one direction). I have listed the causes and corrective actions to be acted upon to effective resolve this issue.
Taking Readings and Minimising Drift Effects.
There are several schools of thought as to the best way to take measurements. Some authorities suggest that the solutions should be stirred slowly with a magnetic stirrer at 50 - 100 rpm during immersion of the electrodes - but care must be taken to ensure that there is no heat exchange between the stirrer and the solution and that the stirrer is always at the same speed. Big differences in mV readings can occur if the speed is varied. Moreover, the magnitude of this effect is related to concentration (most variation in dilute solutions) and hence will affect the slope of the calibration graph. Some prefer to take a reading whilst stirring, others suggest that it is better to switch off the stirrer and take a reading in a still solution - but this can greatly increase the time for the measurement because it may take several minutes for the mV reading to stabilise whilst stirring, and several more to stabilise after the stirrer is switched off. Alternatively, and more easily, the solution can simply be swirled manually after immersion of the electrodes (to ensure good, homogeneous contact between solution and membrane- i.e. no air bubbles) and then left to stand. This avoids the problems of heat transfer and the inconvenience of adding and removing the magnetic stirrer every time. It must be noted however that some electrode systems stabilise more quickly and give more reproducible measurements when stirred whilst others give better results in still solutions - this is sometimes mentioned in the electrode operating instructions.
The time at which the measurements are taken can also vary depending on the characteristics of the particular type of electrode system being used and the balance between time constraints and precision requirements. In some cases it is best to wait for a stable mV reading (this can take several minutes). In others it is better to take all readings after a pre-specified time after immersion. Generally, the mV reading changes rapidly in the first 10 or 20 seconds as the ISE membrane equilibrates with the solution, then more slowly and exponentially as the reference electrode liquid junction poential stabilises. Always taking a reading after say 1 or 2 minutes (depending on which electrode system is being used) should ensure that all are taken in the shallow part of the stabilisation curve where only small and insignificant changes are occuring. A third alternative is to observe the drift in reading as the electrodes equilibrate after immersion and then take a reading at the point where the direction of drift is definitely reversed - i.e. a different electrochemical process begins to dominate - but this last effect is not common.
If the highest possible precision is required then it is suggested that each operator should make his own simple experiments by comparing the reproducibility of repeated measurements on the same sample using each method discussed above, in order to determine which is most suitable for his particular application. But note that whichever method is preferred, the same procedure must be used for all standards and samples in order to ensure that the electrode response will be as uniform as possible.
Maintenance and Storage
Lastly, to ensure optimum electrode performance, the ISE must be maintained and stored correctly.
6.5 Recommendation (Incremental Techniques)
These methods can potentially yield even more accurate and precise results than direct potentiometry because the calibration and sample measurement stages are made essentially at the same time and in the same solution (so that Ionic Strength and temperature differences are not significant) - and the electrodes remain immersed throughout the measurements so that hysteresis, memory, and variations in the reference electrode liquid junction potential are eliminated. They are paticularly useful for samples with high ionic strength or a complex matrix.
In analytical situations where a large number of very different measurements have to be made with different electrodes, it would be laborious to calibrate at several points for only two or three samples. In such situations the incremental addition techniques are more suitable.
There is a common misunderstanding that these incremental methods will also help counteract the effects of Ionic Inteference - but this is not so. If interfereing ions are present in sufficient concentration, then they will always add to the signal produced by the primary ion and increase the measured voltage. The measuring system has no way of distinguishing between the signal produced by the primary ion and that from the interferent. Thus significant concentrations of interfereing ions will always produce a falsley high reading no matter what method of analysis is used.
In principle, all incremental addition techniques operate on the same basis. The electrode potential of a known volume of unknown solution is measured. A small volume of a known solution is added to the first volume and the electrode potential re-measured, from which the potential difference (E) is found. By solving the following equation the unknown concentration can be found:
Summary of Advantages over Direct Potentiometry
-
The electrodes remain immersed throughout the process so that there is little change in the liquid junction potential of the reference electrode (which can often be changed by several millivolts when the electrodes are removed from one solution and placed in another) between calibration and sample measurement - and therefore this source of measurement error is virtually eliminated.
-
Calibration and sample measurement are both made essentially at the same time and in the same solution so that ionic strength and temperature differences between standard and sample are reduced and TISAB is not generally required.
-
Once the approximate concentration for the samples is known, the calibration (slope) can be "fine tuned" by analysing a standard with a concentration that lies within the range of the samples (and is at the same temperature) and then adjusting the slope and re-calculating the results until the standard gives the correct answer.
- Measuring the slope at or very near to the sample concentration means that these methods can be used with old or worn electrodes which may not be completely linear over their whole range, as long as the slope is stable and reproducible over the limited range of the samples.
The main disadvantage of these methods is that it is necessary to mix together accurately measured volumes of standard and sample, and they involve more complex calculations than simple direct potentiometry. This makes them a more skilled and time-consuming procedure and they have traditionally been less popular with most ISE users. It must be noted, however, that these analyses can now be made much more easily and quickly using software to calculate a suitable volume and concentration for the standard, measure the electrode potentials, and calculate the results - and by using automatic syringe pipettes to dispense the measured volumes.
Further slight disadvantages are that this calculations are only valid within the linear range of the electrode and hence these methods cannot be used for low-concentration samples, near to the detection limit, and the approximate concentration of the sample must be known before commencing the analysis in order to choose an appropriate standard concentration and suitable volumes for the two solutions. Nevertheless, if the sample concentration is completely unknown, it can easily be determined by making a quick Direct Potentiometry measurement first. This can be done using an old calibration graph - but the most accurate results can only be obtained if the electrodes are calibrated using two standards which span the expected range of the samples immediately prior to analysis.
7. Conclusion
Conclusively, we can state that the calibration of the ISE was successful despite having low slope value and experiencing severe drifting effects. In the discussion section I have suggested several corrective actions that should be taken to resolve the low slope value and drifting effects. Furthermore, I have established a t-test and determine the source of the unknown at 99% confidence. The results of the t-test for nitrate and chloride have unanimously concluded that the unknown corresponded only to sample D.
8. References
Internet:
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Vedyadhara, 2010. Unit 2 Potentiometry-I [online]. Available from: [Accessed 23 June 2011].
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Frederiksen, 1997. Instruction Manual and Experiment Guide for PASCO scientific model [online]. Available from: [Accessed 24 June 2011].
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Alliance Technical Sales, 2010. A guide to ion selective measurement [online]. Available from: [Accessed 24 June 2011].
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NICO2000, 2011. A beginners guide to ion-selective electrode measurement [online]. Available from: [Accessed 24 June 2011].
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Thermo Scientific, 2008. User guide: Nitrate ion selective electrode [online]. Available from: [Accessed 24 June 2011].
Books:
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G.H. Jeffery and J. Bassett, 1989. Vogel’s Textbook of Quantitative Chemical Analysis. 5 ed. UK: Longman Scientific & Technical.
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Gary D. Christian, 2004. Analytical Chemistry. 6 ed. USA: Wiley.
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