Determination of nitrate ion levels in water samples by use of an ion selective electrode.
Many recent developments in analytical measurement are based on methods utilizing ion selective electrodes. In many applications, ion selective electrode methods are being substituted for existing analytical methods with resulting increases in efficiency and simplicity of measurement.
Ion selective electrodes are usually classified by the type of ion-sensitive membrane used in their construction, and by the ion for which they are selective. The membrane in a solid-state electrode is a thin section of an inorganic crystal. Liquid ion-exchange membranes consist of an inert porous plastic substrate which is saturated with a water immiscible organic solvent containing an inorganic salt of the ion be measured. A glass membrane is a thin section of glass which is formulated to be responsive to the ion of interest.
The nitrate ion selective electrode used in this experiment is of the liquid-membrane type and can be used for the rapid determination of nitrate in solution. It is particularly useful for the measurement of nitrate in water samples
The response of an ion selective electrode, i.e., the potential developed, is a function of the ionic activity of a species in solution. When the activity increases, the electrode potential becomes more positive if the electrode is sensing a cation, more negative if the electrode is sensing an anion. When the electrode potential of an electrode is measured against a suitable refer-ence electrode (e.g. a standard calomel electrode), its relationship to ionic activity is logarithmic.
The behaviour of the electrode potential is determined by the Nernst equation, Equation 1:
|where||E||is the measured total potential of the system|
|Ea||is the portion of the total potential due to choice of reference electrodes and internal solutions|
|R||is the universal gas constant (8.314 J.K-1.mol-1)|
|F||is Faraday's constant (96,485 C.mol-1)|
|n||is the charge on the ion|
|T||is the temperature in Kelvin|
|a||is the activity of the ion in the sample solution|
For a tenfold change in ionic activity, the electrode potential (at 25C) changes by 59.2 mv if the ion being measured is monovalent, 29.6 mv if the ion being measured is divalent. In practice, the electrode slope may be slightly less than theoretical which may lead to measurement errors if compensation is not provided.
While ion selective electrodes respond to ionic activity, it is often desired to know the total amount, or concentration of a species in solution. Activity is related to, but is not necessarily the same as concentration. Two factors contribute to this difference. First, activity refers only to free, unbound ions. A portion of the total concentration may be bound or complexed by other ions in the solution and therefore does not exist in the free ionized state.
Second, ionic activity depends on the total ionic strength of the solution, which is a measure of its overall ionic composition. In general, as ionic strength increases, the ratio of ionic activity to the ionic concentration (that portion of the total concentration which is uncomplexed) for a given species changes. This ratio is called the activity coefficient.
The difference between activity and total concentration then depends on the total ionic strength of the solution and presence or absence of complexing species. For solutions which contain only uncomplexed ions, the total concentration is simply the ionic activity divided by the activity coefficient.
A number of techniques are available for determining concentration even though the electrodes sense activity:
Provided that the electrode has been calibrated (i.e. its e.m.f. with respect to the standard calomel electrode, s.c.e., is plotted against log (concentration) then the sample solution can be analysed by a single direct potentiometric measurement by interpolation from the calibration curve. The standard solutions used for calibration usually differ by a factor of ten and must be close to the expected activity or concentration of the sample.
Direct measurement methods are therefore useful where samples are essentially pure solutions of the ion sensed, or have a relatively high and constant total ionic strength, which can be reproduced in the solutions used for calibration.
Known increment and known decrement methods utilize the principle that the total concentration of a species in a sample can be determined from the change in electrode potential that occurs when the concentration is changed by an incremental amount. The incremental change in concentration can be made either by adding a known amount of the species to increase the concentration (known increment), or by adding a reagent which stoichiometrically reduced the concentration by precipitation or complexation (known decrement).
Known increment and known decrement theory is as follows: If Cs is the concentration of the species being measured, DC the incremental change in concentration caused by adding or removing an increment of the ion of interest, S the Nernst factor or electrode slope, and DE the change in electrode potential caused by DC, then
for known increment method
for known decrement method.
Potentiometric titrations using ion selective electrodes permit the determination of either ions which can be sensed by the electrodes, or species which react with these ions to form other ions, complexes, or precipitates. The electrodes are used to detect the endpoint of the titration, which is indicated by a rapid change in the potential developed in the electrode system.
Electrode titrations permit the analysis of samples that cannot be titrated by conventional colorimetric methods because of sample colour or turbidity. Also, they generally permit a greater degree of precision than direct potentiometric methods. There are two reasons for this. First, the electrodes are used to detect a change in potential rather than measure an absolute value. Second, this change in potential occurs at a much greater rate than the response slope associated with direct potentiometry. This results in greater measurement precision.
In this experiment you will use the direct measurement approach using an Orion Model 701 pH/mV digital meter and the Nitrate Ion Electrode. Instruction Manual provides instructions on the use of the electrode with this meter (pp. 8-11 for calibration and normal measurements; limits of detection and interferences are covered in pp. 24-27). The procedure is described in the Nitrate Ion Electrode Instruction Manual, which also contains a good discussion of possible sources of interference and error. The sections of this manual that are important to this practical have been duplicated and are located on the bench.
The electrode response to nitrate ion is almost that described by the Nernst equation (Equation 1) over a wide range of nitrate concentration but at low concentrations (< 1 ppm) the sensitivity is reduced. The overall calibration line is therefore a curve.
The activity of an ion is proportional to the concentration of that ion in solution. For very dilute solutions, it is possible to replace activity by concentration, as the proportionality constant approaches one.
Determine the calibration curve using pure sodium nitrate standards and ammonium sulphate to maintain an effectively constant ionic strength. The ionic strength of a solution is defined by Equation (2).
|where||I||is the ionic strength|
|zi||is the charge on the ion i|
|mi||is the molality of the ion i|
It is important that the ionic strength of the solutions being analysed is roughly the same. The ionic strength of a solution is a measure of the number and strength of interactions between ions in a solution. In this experiment a constant ionic strength is maintained by adding 2 mL of ionic strength adjuster [(NH4)2SO4] to 100 mL of solution. (The ionic strength adjustor should be added before the solution is made up to 100 mL.)
Prepare solutions containing 100, 70, 10, 7, 1, 0.7 and 0.5 ppm nitrogen (not nitrate) from the stock standard sodium nitrate solution. This stock solution contains 100 ppm nitrogen. The solutions you make up should all contain 2 mL of ionic strength adjuster.
NOTE that in water analysis nitrate ion is often reported as ppm N rather than ppm NO3-. As the molar weight of NO3- is 62 and that of N is 14, a 10-3 M solution of NO3- will contain 14 mg.L-1 of N and therefore will be 14 ppm as N.
The electrode calibration can be affected by changes in pH and as natural water always contains carbonate the solutions should be adjusted to a pH of 4 during the calibration and measurement of unknowns. Use a pH meter to determine the pH of the solutions and adjust to a pH of 4 by adding 0.01 M sulphuric acid drop-wise until the solutions are within 0.2 units of pH 4.
Measure the EMF of the electrode/reference electrode immersed
in the stirred calibration solutions at 298.2 K (25°C).
In order to estimate the errors in a single measurement, make five readings of the EMF at both a high and a low nitrate concentration. Plot the measured EMF against the logarithm of the nitrate ion concentration (as ppm N) using semi-logarithmic graph paper. What is the slope of the linear portion of the calibration graph? How is this slope related to the Nernst equation (Equation 1)? (Remember, for dilute solutions, the activity of the solution can be replaced by the concentration.)
In order to see the effect of chloride ions on the electrode calibration, redetermine the calibration curve in the presence of 200 ppm chloride ion. This is achieved by preparing a new set of solutions, each containing 200 ppm chloride (from 1000 ppm chloride stock solution) and various concentrations of nitrate (from 70 to 0.5 ppm nitrogen). Again, the 2 mL of ionic strength adjustor should be added before the solution is made up to 100 mL.
What are the differences in the calibration curve with and without chloride present.
Estimate the lower detection limit for nitrate with and without chloride ion; i.e. where does the calibration curve become non-linear?
From the comments in the Manual and your calibration curves could you use the nitrate ion electrode for the determination of nitrate in sea water. (What is the ionic composition of normal sea water?).
Natural water samples containing nitrate ion, especially at low concentrations, must be treated to prevent bacterial activity reducing the nitrate concentration before measurement. Boric acid is used as a preservative to prevent this and should be added to your samples at the time of collection. The amount recommended is 2 mL of 1 M boric acid to 100 mL of sample.
Measure the nitrate ion concentration of the water samples (2 mL of ionic strength adjuster should be added to 100 mL of each sample and the pH should be adjusted as previously mentioned before any EMF measurements are made).
Comment on the concentrations of nitrate in the samples.
Nitrate in a soil sample can be measured by shaking 10 g of soil in 25 mL of 0.01 M ammonium sulphate solution: there is no need for the constant ionic strength adjustment as the 0.01 M ammonium sulphate will effectively be the same in the solutions. Shake for 15 minutes. There may be no need to centrifuge the mixture - place the nitrate ion and reference electrodes directly in the soil extract suspension.
Harris, D.C. "Quantitative Chemical Analysis; 5th Edition"; W.H. Freeman and Company: New York..
Skoog, D.A.; West, D.M.; Holler, F.J. "Fundamentals of Analytical Chemistry; 5th Edition"; Sanders College Publishing: New York, 1988.
Christian, G.D. "Analytical Chemistry; 4th Edition"; John Wiley and Sons: New York, 1986.
Atkins, P.W. "Physical Chemistry; 5th Edition"; Oxford University Press: Oxford, 1994.