Ion-selective Measurement (ISE)

Ion-selective Measurement (ISE) Page 1 of 39

Table of contents 1 Basic information 1.1 Introduction to ion-selective measurement 1.2 Reference electrode design and principle of operation 1.3 Concentration and activity 1.4 Diffusion voltage, diffusion potential 1.5 Stirrer 1.6 Selectivity 2 Description of ion-selective electrodes 2.1 Solid-state electrodes 2.2 PVC membrane electrodes 2.3 Gas-sensitive electrodes 3 Calibration 3.1 One-point calibration 3.2 Two-point calibration 3.3 Multi-point calibration 4 Sample preparation 4.1 ph value 4.2 Interfering ions 4.2.1 Precipitation 4.2.2 Complexing 4.2.3 Cooking 4.3 Trace analysis 4.3.1 Shifting the nonlinear region of the characteristic curve 4.3.2 Extraction 4.3.3 Ion exchangers 5 Measurement methods 5.1 Direct potentiometry 5.2 Incremental methods 5.2.1 Standard addition 5.2.2 Double standard addition 5.2.3 Standard subtraction 5.2.4 Sample addition 5.2.5 Sample subtraction 5.2.6 Standard addition with blank value correction 5.3 Titration methods 5.3.1 Direct titration 5.3.2 Indirect titration 5.3.3 Back titration 5.3.4 End point indicator Page 2 of 39

6 Maintenance and special information 6.1 Electrode contamination 6.2 Cleaning the electrodes 6.3 Storing the electrodes 6.4 Faults during measurement and calibration 6.5 Measuring ranges 7 Bibliography 8 Appendix: Interfering Ions Page 3 of 39

1 Basic information 1.1 Introduction to ion-selective measurement Ion-selective measurement is an analytical method that frequently offers advantages over other methods. The most important advantages are: 1. Relatively low purchase costs 2. Rapid arrival at measuring results 3. Use in turbid and coloured solutions in which other methods such as photometry cannot be applied. This list of advantages could easily be extended. Of course, this method also has some disadvantages over other methods. At an error rate of 2-5%, the measuring accuracy is not particularly high. Nevertheless, it is perfectly adequate for many applications. Moreover, the method unfortunately cannot be used for every substance. It generally does not work for organic substances. However, it is particularly well suited for measuring ions, i.e. for the measurement of salts. Ions, the components of salts, are electrically charged particles. When an electrode is introduced into an ion solution, the ions interact with the electrode surface. Interaction means that ions accumulate on the electrode surface. This can be a two-part adsorption of the ions. Alternatively, ions may penetrate the electrode material. Finally, ions may become detached from the electrode and enter into the solution. An electrode is an electrical conductor that is immersed in the solution at one end and connected to a millivoltmeter at the other. In its simplest form, it is nothing more than a simple metal wire. Incidentally, substances that form free mobile ions in aqueous solutions are called electrolytes. Often this term is also used to refer to the ion solution itself. What is the significance of the interaction of the ions with the electrode? Ions are electrically charged particles. They carry an electrical charge onto the electrode, thereby generating a voltage between the electrode and the solution. In the ion-selective measuring technique, as shown below, the magnitude of this voltage is a function of the concentration of the ion to be measured. However, our measuring device in its present form is not yet capable of measuring a voltage at all since the second input of the device has not yet been defined. It would be ideal if we could simply connect the second input directly to the solution. However, this cannot be accomplished without the use of an electrical conductor. If we were to use such a conductor and immerse it in the solution, the same processes that take place at the surface of the first electrode would also occur at this conductor. Experience in the field has provided us with a similar solution. The second input is connected to a reference electrode. Essentially, this reference electrode operates in a manner similar to that already described above. However, it exhibits a constant voltage difference between the electrolyte output, with which the reference electrode is immersed in the solution, and the connection to the device. Thus, we have a fixed reference point with respect to which we can measure the voltage at the first electrode. If the surface of the first electrode is now coated with an ion-selective substance, i.e. with a substance that only interacts with a specific ion species, we then have the basic component needed for working with the ion-selective measuring technique. The ion-selective electrode and the reference electrode are collectively known as an electrode. If they are physically separated, they are referred to as a two-probe electrode. If they are combined in a single probe, they are known as a combination electrode. Page 4 of 39

As already indicated above, ion-selective measurement is about measuring voltage changes in order to draw conclusions about ion concentrations. To do so, we need to establish a physicalmathematical relationship between these two parameters. This is provided by the Nernst equation as the expression: RT E = E0 + 2,303 log c nf E Measured voltage between the measuring and reference electrodes E 0 Measured voltage between the measuring and reference electrodes at a concentration of c = 1 R Universal gas constant (R = 8.314 Joule mol -1 K -1 ) T Absolute temperature in Kelvin (the absolute temperature is defined as T = 273.15 + t. The units are Kelvin K. t is the numerical value of the measured temperature in C. F Faraday constant (F = 96485 C mol -1 ) n Electrical ionic charge c Ion concentration R and F are constants and therefore not of further concern. Since temperature appears in the Nernst equation, the measured voltage is a function of the temperature. The ionic charge n influences a parameter we call the electrode slope. The Nernst slope is given by the expression: RT 2,303 nf Depending on the ionic charge, the slope has the following values in millivolt at 20 C and 25 C: Ionic charge Slope 20 C Slope 25 C Examples +2 29.08 29.58 Copper (Cu 2+ ), lead (Pb 2+ ) +1 58.16 59.16 Sodium (Na + ), potassium (K + ) -1-58.16-59.16 Fluoride (F - ), chloride (Cl - ) -2-29.08-29.58 Sulphide (S 2- ) In practice, many electrode types actually achieve these slopes. Only a few do not. If the slope decreases over time, this generally indicates ageing of the electrode. Finally, we should discuss the expression log c. While it may be difficult for some readers to understand this expression, this should not be grounds for rejecting the ion-selective measuring technique. After all, log c is usually not needed when performing standard measurements. Also, any good calculator today is capable of calculating the value. Example: Using a calculator that features the log function, enter the numerical value for concentration c and press the log key. The number that appear on the display is the value of log c and can be substituted in for log c in the Nernst equation. (A brief remark is in order regarding concentration c: For small and very small concentrations, the Nernst equation applies without limitations. Deviations arise for larger concentrations. Later we shall see that this problem can be solved by introducing the activity parameter.) Page 5 of 39

We will now try to plot the characteristic curve of an ion-selective electrode. This will make it easier to perform analyses later on. Also, this task is a good example that will help you understand this measurement technique. The characteristic curve of an ion-selective electrode represents the relationship between voltage E and concentration c. In other words, it is a graphic representation of the Nernst equation. Because of the log c expression, it is not practical to plot the concentration c itself on an axis. It is much more convenient to use the value of log c. The plot then takes on the form of a straight line. This procedure is easy to master. Using semi-logarithmic paper, we plot log c on the logarithmic axis and voltage E on the other. An example illustrates how this is done. Now let us move on to the practical part. We will record the characteristic curve of a bromide electrode. To do so, we require the following equipment and chemicals: Instrumentation 1 millivoltmeter (resolution of 0.1 mv, input impedance 10 12 ) 1 bromide electrode (BR 500, Br 501) 1 reference electrode (R 502) 1 stirrer with stirrer bar 1 electrode stand 6 plastic beakers (150 ml volume) 8 graduated flasks (100 ml volume) 1 pipette (1 ml volume) 1 pipette (10 ml volume) Bromide standard with 10000 mg/l bromide Conducting salt solution: solution of 5 mol/l or 425 g/l sodium nitrate (The reason for using conducting salt will be explained later on.) We dilute the bromide standard to prepare two additional standard solutions, one of 100 mg/l bromide and the other of 1 mg/l bromide. Preparation of the 100 mg/l bromide standard Add 1 ml of 10000 mg/l bromide standard to a 100 ml graduated flask. The flask is filled to the calibration mark with deionized water and mixed well. Preparation of the 1 mg/l bromide standard Add 1 ml of 100 mg/l bromide standard to a 100 ml graduated flask. The flask is filled to the calibration mark with deionized water and mixed well. Measuring solutions are prepared in a similar manner according to the following table: Concentration of measuring solution mg/l bromide Volume of standard solution ml Standard solution mg/l bromide Volume of conducting salt solution ml Measured values mv 1000 10 10.000 2-73.7 100 1 10.000 2-15.8 10 10 100 2 43.3 1 1 100 2 101.9 0.1 10 1 2 152.8 0.01 1 1 2 170.3 Page 6 of 39

To prevent mistakes, the preparation of the 1000 mg/l bromide measuring solution is described here in detail. Preparation of the 1000 mg/l bromide measuring solution Add to a 100 ml graduated flask: 10 ml 10000 mg/l bromide standard 2 ml conducting salt solution Fill the flask to the calibration mark with deionized water and mix. The other solutions are prepared in a similar manner according to the table. The measuring solutions are now poured into plastic beakers and placed on the stirrer one after the other to be measured. The electrodes connected to the millivoltmeter are dipped into the solution with the aid of the electrode stand. The solution is stirred while readings are taken. The reason for stirring is described later on. The measured value usually drifts at the beginning of a particular measurement. After some time, the readings stabilise and the final value can be read and recorded before the next solution is measured in a similar manner. To avoid transferring substances from one measurement solution to the other, the electrodes are rinsed with deionized water after each measurement and the residual water is blotted with a soft paper towel. Surely not every reader will be able to reproduce these steps immediately. Therefore, the table also contains the values we ourselves measured in millivolts. How can we graphically represent the characteristic curve of the bromide electrode? A look at the semi-logarithmic graph paper reveals that the lines have a 1-mm spacing in one direction only. The voltage values are plotted in this direction, i.e. values of approx. 80 to approx. 180 mv. We recommend that 1 mm be used to represent 1 mv. In the other direction the line spacing varies. This is the logarithmic scale. The lines are arranged in such a way that when the concentration is plotted, the log c logarithm automatically generates a linear plot. This is precisely what we require and why the line spacing varies and repeats periodically. The same applies to the plotted numbers 2 to 10. For each number 10, we draw a short line and write below it the concentration of our measuring solutions, beginning from the right and moving along in ascending order. We can now plot the measuring points as usual. Fig. 1 shows this graph. Figure 1: Characteristic curve of a bromide electrode as a function of the decadic logarithm of the concentration Characteristic curve of bromide electrode 200 150 170.3 152.8 Voltage mv 100 101.9 50 43.3 0 0.01-50 0.1 1 10-15.8 100 1000-73.7-100 log c Page 7 of 39

The following information can be gathered from this plot: 1. The plotted points for concentrations 1, 10, 100 and 1000 mg/l bromide lie in what is almost a straight line. This is the strictly linear part of the characteristic curve. The voltage difference between neighbouring points and, in general, between the concentrations, which have a ratio of 1 to 10, is 58.4 mv. We already encountered a value of this magnitude for the Nernst slope. In this case, however, it is referred to as the electrode slope. Mathematically, the value of this slope is negative, i.e. S = -58.4 mv. This is easy to remember. All negatively charged ions, also known as anions, have negative slopes. A positive slope can also occur. The characteristic curve is then rotated by approx. 90. This is the case for all positively charged ions, which are also called cations. 2. The measuring points for concentrations 0.1 and 0.01 mg/l bromide no longer lie on the straight line. Thus, the characteristic curve has a nonlinear part. All ion-selective electrodes exhibit this phenomenon. The reason for this behaviour is found in the electrode structure. The ionselective membrane usually already contains the ion that is to be measured. Unfortunately, there are no perfectly insoluble substances in this field. Thus, a small quantity dissolves out of the membrane and introduces into the solution precisely that ion which is to be measured. If the measured concentrations are low, the magnitude of the dissolved ion concentration is significant. There is also another bothersome phenomenon found within this measuring range. The stabilisation period of the voltage can become extremely long. Frequently, one cannot be sure at what point the final value has actually been reached. 3. The characteristic curve can be used to determine the concentration of an unknown sample. The sample is usually prepared in the same manner as the solution for recording the characteristic curve. In this particular case, 2 ml conducting salt solution are added to 100 ml sample. The voltage of the solution is determined during stirring. The voltage reading is compared to the values in the plot of the characteristic curve and the corresponding concentration is found. This is done by drawing a line parallel to the concentration axis from the voltage value to the characteristic curve. At the point of intersection, a perpendicular is dropped to the concentration axis. The point of intersection with the concentration axis represents the concentration value we are looking for. To be precise, this value is actually the concentration of the measuring solution rather than of the sample itself. The sample was diluted. By performing a simple calculation, we can arrive at the concentration of the sample. For example, let us assume we have a measured value of 25 mg/l bromide. Our measuring solution had 102 ml and the sample had 100 ml. Thus, the absolute bromide quantity in the plastic beaker is 25 mg/l * 0.102 l = 2.55 mg bromide. A 100 ml sample therefore contains a concentration of 2.55 mg/0.1 l = 25.5 mg/l bromide. Even if we haven t yet done so practically, we have at least theoretically carried out the first analysis using the ion-selective measuring technique. Incidentally, when this measuring technique was first developed, it was used in just this way. Only later did some companies build the first ion meters, which eliminated the need for recording measuring points. Today, modern ion meters are highly advanced and simplify the task for the user by offering many additional functions. However, even these devices require that at least two points of the characteristic curve be recorded, i.e. two solutions of know concentrations are needed. The user must enter these concentrations into the device. The device can then measure the corresponding voltages and determine the slope. The slope is generally indicated on the display. This entire procedure is referred to as calibration. In this particular case, it is also known as a two-point calibration. Most users employ this method. You should proceed in such a way that the expected sample concentration lies between the two calibration concentrations. This ensures that measurement imprecisions will have the lowest impact on the results. Page 8 of 39

When using new ion-selective electrodes for the first time, voltage stabilisation may be very sluggish at first. This does not necessarily mean that the electrode is defective. Rather, the electrode first requires conditioning, meaning that it must be kept in a diluted standard solution for several hours. This builds a so-called solvent layer on the membrane surface that accelerates voltage stabilisation. Of course, there is a lot more to say on many of the topics discussed here. There is also much to add from the point of view of physics and chemistry. Up to now we have refrained from touching on these topics to give the reader a chance to understand the basics of ion-selective measurement without the distraction of extraneous information. On the following pages, we now make up for this lapse. 1.2 Reference electrode design and principle of operation The most important property of a reference electrode has already been mentioned. At a fixed temperature, it guarantees that the voltage between the measuring device connection and the electrolyte outlet remains constant. The electrolyte will be referred to as the diaphragm from here on in. A constant voltage is indispensable for being able to measure voltage changes at the ionselective electrode. The function of the diaphragm is to provide the necessary electrical contact with the measuring solution. This is established by an electrolyte passing through the diaphragm. There are several types of diaphragms, including the ground diaphragm, ceramic diaphragm, pinhole diaphragm and platinum wire diaphragm. Ion-selective measurement usually employs the ground diaphragm. It has a relatively high flow rate and is easy to clean. It is also the diaphragm type that most reliably provides the defined electrical conditions with respect to the measuring solution. This issue is further discussed under Diffusion Potential or Diffusion Voltage. Since electrolyte continuously flows out of the reference electrode, it must be ensured that there is always an adequate quantity of electrolyte inside the reference electrode. When performing measurements, the electrolyte level within the reference electrode should be about 1 cm higher than the level of the measuring solution. How does the reference electrode generate a constant voltage? In the silver-silver chloride reference electrode, a silver wire coated with silver chloride plays a very important role. It is usually immersed in a chloride solution. Depending on the chloride concentration, a constant silver ion concentration is generated in the chloride solution (in a chemistry book, you will find this concept under solubility product). In accordance with the Nernst equation, the silver concentration generates a constant voltage between the chloride solution and the silver wire. The processes that take place at the silver wire have already been described in the introduction. This is the design of a simple reference electrode. For many applications, it does a good job of accomplishing the required tasks. However, what happens if we would like to measure small chloride ion concentrations? Then the ions that flow through the diaphragm into the solution are precisely those that we want to measure. Obviously, this does not allow for accurate measurement. In practice, therefore, the reference electrode consists of two parts. The actual reference electrode is contained in an internal chamber. Beyond it there is an additional diaphragm that leads to the outer chamber, which in turn comes into contact with the measuring solution through the ground diaphragm. The internal chamber has a chloride as an electrolyte (internal electrolyte, usually coloured), while the outer chamber contains an electrolyte of your choice (outer electrolyte). In this way we can select the electrolyte that flows into the measuring solution to match the measuring task at hand. The outer chamber is also called the bridge or electrolyte bridge, and the electrolyte is also referred to as the bridge electrolyte. Page 9 of 39

As already indicated, the voltage between the device connection and the ground diaphragm is only constant at a fixed temperature. A change in temperature brings with it a change in the voltage. This change is described by the Nernst equation and by the fact that the solubility of the silver chloride on the silver wire is temperature-dependent. Discharge electrode Reference electrolyte solution Internal bridge Bridge electrolyte solution Outer bridge Figure 2: Reference electrode Unfortunately, all voltages that contribute to the electrode voltage are temperature dependent. We have already noted the temperature-dependence of the electrode slope of the measuring electrode. To keep measurement errors caused by temperature fluctuations to a minimum, there is a basic rule that says the calibration temperature and the measuring temperature may not differ by more than 2 C. In modern WTW ion meters, the display blinks if this limit is exceeded. In addition to the silver-silver chloride electrode described here, which is the electrode type most commonly used today, other types of reference electrodes are also available. They differ in the nature of the electrochemical reactions on which they are based. The most familiar types are the calomel electrode and the thalamide electrode. The normal hydrogen electrode occupies a special position among electrodes. It is primarily employed in scientific research. It consists of a platinum plate that is coated with platinum black and immersed in hydrochloric acid with a concentration of 1 mol/l. Pure hydrogen gas at a pressure of p = 1 bar is bubbled over the electrode. By definition, this type of reference electrode has a potential of 0 mv at all temperatures. Because it is difficult to handle, it is not used for standard field work. Instead, either of the following electrodes are used: the silver-silver chloride electrode, which at 3 mol/l potassium chloride and a temperature of 25 C has a voltage of 208 mv with respect to the normal hydrogen electrode; or the calomel electrode, which at a temperature of 25 and a saturated potassium chloride solution exhibits a voltage of 244 mv with respect to the normal hydrogen electrode. Page 10 of 39

If a voltage measured against a particular reference electrode is to be converted to the corresponding value for the normal hydrogen electrode, the following simple conversion equation is used: U = U + U H gem U H = Voltage that would be measured using the normal hydrogen electrode U Gem = Measured voltage with respect to the reference electrode actually in use U Ref = Voltage of the reference electrode actually in use with respect to the normal hydrogen electrode Re f It is of no relevance which type of reference electrode is actually used for the ion-selective measurement. It has no effect on the nature of the characteristic curve. The use of different reference electrode merely causes the characteristic curve to shift to larger or smaller voltage values. 1.3 Concentration and activity We have already established that for low ion concentrations, the Nernst equation yields the correct voltage values. (This of course does not apply to concentrations in the nonlinear region of the characteristic curve). The calculated values closely coincide with the measured values. When measuring higher concentrations, however, deviations begin to appear. The ion solutions then behave as though a part of the ions were no longer present. There is an apparent loss in ions. This is a common phenomenon in chemistry. It is related to the fact that ions in high concentrations impede each other. To eliminate this source of error, a new entity, namely the activity, was introduced. The activity represents that portion of ions that are free to act. Mathematically, the activity is defined by the following equation: a i = f c i i a i = Activity of ion i c i = Concentration of ion i f i = Activity coefficient of ion i The activity coefficient is a function of the concentration and approaches the value of f i = 1 as the solution becomes more and more dilute. The activity then becomes equal to the concentration. With an increase in the concentration, the activity coefficient falls to values < 1 and the activity becomes smaller than the concentration. It has been shown that the activity coefficient depends not only on the concentration, but also on the ionic charge. A parameter that takes this into account is the ionic strength J. It is defined by the following equation: c i = Concentration of ion i z i = Charge of ion i J = 0,5 c i z i 2 A continuation of this line of thought produced a relationship between the activity coefficient f i and the ionic strength J. log f i = A J A = Constant Page 11 of 39

This equation is a mathematical expression for what we already know, namely that the activity coefficient depends on the ionic strength, which is a function of the concentration and charge of all ions in a solution. The addition of a conducting salt to the calibration and measuring solutions generates the same activity coefficient for all solutions. Under these conditions, the activity, which correctly appears in the Nernst equation, is proportional to the concentration. Therefore, we can then work with this entity as well. The difference between the activity and the concentration is cancelled out during calibration. Voltage Concentration Voltage Concentration Figure 3: Relationship between the concentration and the measured signal without and with the addition of a conducting salt Page 12 of 39

1.4 Diffusion voltage, diffusion potential Two different electrolyte solutions come into contact at the ground diaphragm of the reference electrode. They are the bridge solution and the measuring solution. Generally, they will differ in their composition and concentration. A voltage will usually be generated at bridges of this type. The voltage is generated from the difference in the diffusion speeds of the different ion species. For example, if the cations of one solution diffuse more rapidly into the other solution than the anions, a positively charged liquid front will build up, which is followed by a negative front. Separated electrical charges always give rise to an electrical voltage. Unfortunately, this voltage lies within the measuring circuit and is included in the measurements. The diffusion voltage, or the diffusion potential as this voltage is also called, can rise to values of 30 mv under certain conditions. In particular, this may be assumed to be the case when solutions with very different diffusion properties, such as acids and neutral salt solutions, border on each other. Considering that a measurement error of 1 mv results in a 4% error in the measurement results for monovalent ions and in an 8% error for divalent ions, it is clear that the diffusion voltage is the weakest link in the measurement circuit. Luckily, errors such as these can be prevented. For example, the salt concentration of the measuring solution can be raised. This is accomplished by adding salt to the measuring solution as described in the previous section, which already counts as a sample preparation. However, still more can be done. The generation of a diffusion voltage can be limited by the selection of suitable salts. If the cations and anions wander at roughly the same speeds, the ensuing diffusion voltage will only be small. This is the case for potassium chloride, potassium nitrate, ammonium chloride and, to a limited degree, for sodium nitrate and sodium chloride. The maintenance of identical calibration and measurement conditions is extremely important under this aspect. In technical jargon, the solutions used to add salt to the calibration and measuring solutions are known as ISA solutions. The abbreviation ISA stands for Ionic Strength Adjustor. Reference electrode Measuring medium K + Cl - Cl - Cl - K + K + K + K + Cl- K + ClK - + Cl - K + Cl - + K Cl - Cl - Figure 4: Balanced diffusion at the reference electrode diaphragm Page 13 of 39

1.5 Stirrer When working with ion-selective electrodes, we generally use a magnetic stirrer that allows uniform and comparable stirring of all solutions. Stirring has a stabilising effect on the diffusion voltage, that is, it brings about the constant values we require. Since this also accelerates the transportation processes at the measuring electrode and at the diaphragm, the voltage stabilisation periods are shortened as well. A disadvantage is that the stirrer can warm up the solution. An insulated plate between the stirrer and the vessel can effectively suppress this effect. 1.6 Selectivity Up to now, we have assumed that every ion-selective electrode responds only to one type of ion and remains unaffected by all other species. Unfortunately, this is not the case. For almost all electrodes, there are ions that we do not wish to measure, but that still affect the electrode signal. These ions are often called interfering ions and one speaks of interference error. Building on the Nernst equation, Nikolsky developed an equation that carries his name and allows a mathematical treatment of the interference effect. n i n s a i a j K is RT U = U 2,303 nif a j ai Charge of ion to be measured Charge of interfering ion species Activity of the ion to be measured Activity of the interfering ion species Sum of all interfering ion species Selectivity constant (measured ion interfering ion) n /n j i s 0 + log (ai + Kis a ) Strictly speaking, the selectivity constant K is is not really a constant since its value is a function of the ionic strength. Moreover, K is is also dependent on the method of determination. This especially applies to PVC membrane electrodes. In contrast, the selectivity constant of pure solid-state membrane electrodes is clearly a function of the solubility products. The Nikolsky equation shows that the interference effect increases with a rise in the interfering ion concentration and the selectivity constant. Therefore, it is advantageous to have a small selectivity constant. A small selectivity constant of 0.001 means that the interfering ion, if it has the same concentration as the measured ion, causes only 1/1000 of the resulting voltage. However, it should be mentioned that some manufacturers of ion-selective electrodes specify the selectivity constant as the reciprocal value. This is always the case when the formulation more sensitive than is used. In practice, selectivity constants > 1 may also be encountered. This means: The interfering ion plays a greater role in generating the voltage than the measured ion itself. This is often the case with calcium electrodes, which are three times as sensitive to zinc ions than to calcium ions. As long as the calcium ions are not measured in the presence of zinc ions, as is usually the case, this fact will be of no importance. Page 14 of 39

2 Description of ion-selective electrodes Ion-selective electrodes can be categorised into three groups according to their design: 1. Solid-state electrodes 2. PVC membrane electrodes (matrix electrodes) 3. Gas-sensing electrodes 2.1 Solid-state electrodes The ion-selective membrane in these electrodes is a solid substance that is in direct contact with the measuring solution. Substances suitable as membrane material possess a very small yet detectable electrical conductivity, have ion-selective properties and are sparingly soluble in water, i.e. the solubility product is extremely small. Lanthanum fluoride is used for fluoride measurement since it dissolves in water to only 19.5 g/l. This minute amount is dissociated, i.e. it is completely broken down into ions. LaF 3 La 3 + + 3 F The ions are in equilibrium with the solid membrane and generate a voltage at the membrane that is a function of the fluoride or lanthanum concentration. The membrane has a glassy appearance and is very hard. The silver electrode is of similar design. A membrane of silver sulphide, Ag 2 S, is used. The solubility of this compound is almost a factor 10-10 smaller than the solubility of the lanthanum fluoride. The dissociation is described by the following reaction: Ag + 2 2S 2 Ag + S Just as the fluoride electrode responds to lanthanum and fluoride, the silver electrode responds to silver and sulphide. Silver sulphide is very suitable for creating electrodes of the third type. These are electrodes whose membrane consists of a combination of silver sulphide, Ag 2 S, and either a sulphide of a heavy metal, MeS, or a silver halide, AgX. Both types of compounds are sparingly soluble in water. Electrodes with this membrane design respond to either the heavy metal or the halite ion. The scope of identifiable ions is thus usefully extended. The method of operation of these electrodes can be explained by looking at the equilibrium equations. In the electrodes that respond to heavy metals, the components Ag 2 S and MeS are present in the membrane. Both exist in the solution in a dissociated state: Ag + 2 2S 2 Ag + S MeS Me 2 + + S 2 The second equilibrium is shifted to the left by the addition of the heavy metal ion, Me 2+, i.e. insoluble heavy metal sulphide is formed and the concentration of the sulphide ions, S 2-, decreases. This process affects the first equilibrium shown here. As the sulphide concentration decreases, the equilibrium is shifted to the right. This means: Silver sulphide, Ag 2 S, dissolves and the concentration of the voltage-determining silver ions, Ag +, increases. The sensitivity of the Ag 2 S/AgX mixtures to the halites has a similar explanation. Examples of electrodes with a heavy metal sulphide are the electrodes for lead, cadmium and copper; examples of electrodes with a silver halite are the electrodes for chloride, bromide, iodide and rhodanide. Page 15 of 39

With regard to the copper electrode, it should be added that its ion-selective membrane is photosensitive. The electrode voltage varies depending on the intensity of illumination in the laboratory. This may no longer be tolerated by the acceptance criterion of modern ion meters. Also, the use of a copper chloride standard for the calibration should be avoided because it gives rise to the following equilibrium: 2+ Ag2 S + Cu + 2 Cl 2 AgCl + CuS The silver chloride on the membrane of the electrode is sparingly soluble and slowly converts the copper electrode into a chloride electrode. It delivers an interference voltage that is a function of the chloride concentration. When measuring the sample, small chloride concentrations can be tolerated to a certain degree. This depends on the concentration conditions. More information is provided in the operating manuals of the copper electrode. The cyanide electrode represents a special case. It is actually a standard iodide electrode whose membrane consists of silver sulphide, Ag 2 S, and silver iodide, AgJ. The cyanide ion is a good complex builder for heavy metals and slowly dissolves the silver iodide of the electrode: AgJ + CN AgCN + J or AgJ + 2 CN Ag(CN) 2 + J The iodide electrode responds to the liberated iodide. Electrodes that slowly dissolve in the measured medium are also called corrosion electrodes. Care should be taken not to leave them in the measuring solution for longer than necessary. The iodide electrode, however, can do still more. The thiosulphate ion, like the cyanide ion, is capable of forming stable complexes with silver ions, which makes it possible to measure the thiosulphate. The iodide electrode can even be used to measure mercury silver ions within the concentration range of 10-4 to 10-8 mol/l. Mercury-II salts also dissolve the silver iodide. However, in this reaction it is the silver ion rather than the iodide ion that is liberated. This is of secondary importance for the measurement. 2+ AgJ + Hg HgJ + Ag + + The last solid-state electrode to be discussed here is the sodium electrode. On the surface, it has a different design than the electrodes discussed up to now and is very similar to the ph electrodes. This is not a coincidence. After all, it was developed on the basis of the ph electrode. The electrode voltage of ph electrodes exhibits deviations from the linear characteristic that are called alkaline error and acid error. The alkaline error can be increased by selecting a suitable glass type. If an electrode of this type is placed in a solution with a high constant ph value, it develops a sensitivity for sodium ions. To establish and maintain the ph value, buffer solutions based on organic amines are required. A sodium electrode can determine sodium concentrations as low as 0.02 to 0.002 mg/l. However, preparation of the buffer requires very good water, i.e. water that is low in sodium, and very clean laboratory procedures since traces of sodium are widespread. Page 16 of 39

2.2 PVC membrane electrodes (matrix electrodes) There are many ions for which there are no known membrane substances that are sufficiently sparingly soluble to be able to use them in solid-state electrodes. However, there are substances that are in part organic which are able to form complexes with the ions to be measured or that have ion exchange properties. Substances such as these are dissolved in organic solvents. The substance and solvent should be sparingly soluble in water. The ion to be measured is capable of migrating in both directions across the border area between the organic solvent and the water. This builds up a voltage jump. To create electrodes that are practical to use, the organic solvent is polymerised into a PVC membrane together with the ion-selective organic substance. Many manufacturers build PVC membrane electrodes that are similar in design to the solid-state electrodes, the only difference being that the PVC membrane is found at the end of the head which is screwed onto the other part of the electrode. Because the PVC membrane is exhausted after ¼ to 1 year at the latest, the availability of a separate, exchangeable measuring head lowers the costs of measuring with these electrodes. The lifetime of the membrane largely depends on the frequency of use. 2.3 Gas-sensing electrodes These sensors are actually not counted among the ion-selective electrodes, since they are designed to respond to gasses. However, they can be used to determine ion concentrations as well, and we will discuss them here briefly in this capacity. As an example, we will single out the ammonia electrode, which is used to determine ammonium concentrations. From a chemical point of view, the following equilibrium exists between the ammonium ion, NH 4 +, and ammonia, NH 3 : + + NH3 + H2O NH4 OH This equilibrium indicates that the ammonium ion dominates in a strongly acidic solution, while ammonia gas dominates in a strongly alkaline solution. Each form can be converted to the other by shifting the ph. There is an intermediate range in which the ammonium ion exists alongside the ammonia in a particular ratio. At a ph of 9.25, the ratio is 1:1. At a ph of 12.25, only 1/1000 exists in the form of the ammonium ion, i.e. practically all ammonium ions will have been converted into ammonia. How can the ammonia be measured in order to determine the concentration of ammonium ions? In principle, the ammonia, which is dissolved in the alkaline solution, is allowed to diffuse through a Teflon membrane into a chamber that contains an ammonium chloride solution with a relatively high concentration. The equilibrium shown above now takes effect, i.e. the ammonia reacts with water and produces ammonium ions and OH - ions. The formulation of the equilibrium using the mass action law yields the following expression: a OH a = K a NH3 + NH4 a OH - K a NH3 + a NH4 Activity of the OH - ions Equilibrium constant Activity of the ammonia Activity of the ammonium ion Page 17 of 39

If we assume that the activity of the ammonium ion is constant, we have the following relationship: a OH = K a NH 3 This equation shows: When in equilibrium, the activity of the OH - ions is proportional to the activity of the ammonia. Due to the diffusion equilibrium, this also applies to the activity of the ammonium ions in the original sample. Thus, by measuring the voltage at a ph electrode, we can determine the ammonium concentration. The electrode design differs only marginally from the ion-selective measuring apparatus. We connect the ammonia electrode to our ion meter, calibrate the device with two ammonium standards that have been made alkaline and directly measure the ammonium concentration in the sample that has been made alkaline. To prevent ammonia from escaping from the solution, which would lead to a result that lies below the actual value, it is advisable to calibrate and measure in a closed vessel. If a closed vessel is used, the ammonia electrode does not even have to be dipped into the alkaline solution. Measurements can be made in the gaseous phase because it, too, builds up an equilibrium. Electrodes that work on the same principle as the ammonia electrode described above have been built for other systems as well. The most important ones are: Carbon dioxide/hydrogen carbonate Sulphur dioxide/hydrogen sulphite Hydrocyanic acid/cyanide Hydrofluoric acid /fluoride Hydrogen sulphide/hydrogen sulphide Naturally, the sample preparation differs for each system and the solutions for establishing a defined ph value have different compositions. 3 Calibration The user can select from three different calibration procedures depending on the design and features of the ion meter. 1. One-point calibration 2. Two-point calibration 3. Multi-point calibration 3.1 One-point calibration A prerequisite to using this procedure is that you must have available to you an ion meter on which the slope of the characteristic curve can be adjusted and the characteristic curve itself can be shifted horizontally. The theoretical slope is usually used. The equipment is calibrated using a single calibration standard. The electrode is dipped into the standard. The characteristic curve on the ion meter is now shifted until the reference value of the calibration standard is reached. This procedure is often the method of choice for production control purposes. The concentration of the calibration standard should be very close to that of the product. Under these conditions, a slope that has not been precisely set to the actual value has little bearing on the results. The advantage of this method is the speed with which both the procedure and a calibration check can be performed. Also, it avoids the expense of having to have a second calibration solution. Page 18 of 39

3.2 Two-point calibration The two-point calibration precisely defines the characteristic curve of the electrode by recording two measuring points. At the same time it calculates the slope. Two calibration standards are needed to record the two measuring points. Their concentration must be entered into the device. The standards are generally prepared by the user since solutions with a low concentration have a limited shelf life and also because the sample preparation must be taken into account when calibrating. The concentrations of the calibration standard should differ by a factor of 10. A smaller difference will have a negative effect the accuracy of the calibration. With a larger difference you run the risk of going well into the nonlinear range of the characteristic curve. It is advisable to define the calibration standards in such a way that the expected measured value lies between the two calibration points. Since the characteristic curve is a function of the temperature, the two calibration standards must be at the same temperature. The temperature during measurement should deviate as little as possible from this temperature. A deviation of 2 C is considered to be the absolute limit. E 2 E 1 Figure 5: Example of a two-point calibration c 1 c 2 3.3 Multi-point calibration The purpose of the multi-point calibration is to determine where the characteristic curve deviates from a straight line during calibration in order to achieve a higher accuracy during measurement. This may not be of very great significance for a three-point calibration, since the user can select the required calibration points for the two-point calibration. However, if 5 or 6 calibration points are available, the nonlinear part of the characteristic curve can be reproduced, making this range available for measurement as well. The two-point calibration would fail in this region of the curve. In practice, the multi-point calibration reproduces the entire characteristic curve between the calibration points. One method is to use the polygon procedure where the ion meter uses straight lines to connect the calibration points. However, a more accurate method is the polynomial procedure. Here the device calculates the characteristic curve with a high degree of accuracy using polynomials of the fourth order. Be aware that the polynomial is only defined between the first and last standard. Measuring samples outside of this range can lead to errors. This situation can be avoided by selecting a more suitable calibration standard. Many ion meters output a warning when the measurement range limit is exceeded. When taking measurements, you should also know that the electrode voltages adjust very slowly in the nonlinear range of the characteristic curve. Page 19 of 39

6-point calibration with a nonlinear region Voltage 200 150 170.3 152.8 100 101.9 50 43.3 0-15.8 0.01 0.1 1 10 100 1000-50 -73.7-100 Concentration c Figure 6: Example of a 6-point calibration with a nonlinear region 4 Sample preparation Samples are prepared with the aim of ensuring that the conditions for measurement are optimal. After preparation, the sample to be measured should be in a condition where the ions are in a free state. Any factors that arise from the measuring solution and can affect the electrode must be constant at all times. This applies to both the calibration and the measuring solutions. Any interfering ions should be removed. The most important method of sample preparation has already been presented. It is the use of ISA solutions. In the calibration and measuring solutions, the concentration of ISA solutions with a value of 1.1 should be 0.1 mol/l. The ISA solutions ensure that the electrical conductivity of the solutions is adequate, the diffusion voltage remains constant and the activity coefficient of the measured ion also remains constant. If the sample already has this concentration of salts, then adding ISA solution would be superfluous. Solutions with a salt content greater than 1 mol/l are diluted as required. Other methods of sample preparation are available and are described here in detail: 4.1 ph value It is necessary to establish specific ph ranges if the ion to be measured reacts with the H + or OH - ions at lower or higher ph values. The reaction products cannot be detected by the ion-selective electrode. For example, if ph values are too high, the ions of the heavy metals silver, copper, lead and cadmium are precipitated as hydroxides. Likewise, the measurement of fluoride is disturbed because the OH - ions cause the lanthanum fluoride in the membrane to be converted into lanthanum hydroxide. Especially at low concentrations, the liberated fluoride ions simulate a higher fluoride concentration. On the other hand, incorrect measured values can also result from solutions that are too acidic. For example, this is the case for the fluoride, cyanide, sulphide and thiocyanate anions. From these anions, H + is taken up to form undissociated acids that cannot no longer be detected. As a corrective measure, it is advisable to substitute the ISA solution with a suitable buffer. The specifications in the application reports provided by the companies should be observed. Page 20 of 39

4.2 Interfering ions As mentioned above, there is a very large number of different interfering ions, which is why only the most important examples will be discussed here. The procedures available for eliminating these interfering ions are precipitation, complexing and cooking. 4.2.1 Precipitation High concentrations of lead interfere with cadmium measurement. Since lead sulphate is sparingly soluble, an ISA solution of 5 mol/l sodium sulphate is used here. Iron-II ions interfere with the measurement of lead, cadmium and copper. Several milliliters of hydrogen peroxide are added to raise the ph value to 5-6. By adding ammonia solution drop by drop, iron-iii hydroxide precipitates and can be filtered out. Calcium measurement is disturbed by the divalent ions of iron, copper, magnesium, nickel and zinc. These ions can be precipitated at a ph value of 8. The calcium can then be directly measured at this ph value. Carbonates are destroyed prior to measurement (see below). Nitrate measurement is disturbed by iodide, cyanide, bromide, chloride, hydrogen sulphide and phosphate ions. Here, an ISA solution containing silver sulphate is used. The interfering ions are precipitated. Specifically for nitrate measurement in the presence of chloride, WTW sells a nitrate TISAB solution that contains silver ions (TISAB: total ionic strength adjustment buffer). If carbon acids are also interfering with the nitrate electrode, they can be precipitated by the addition of an ISA solution with 0.1 mol/l aluminium sulphate. 4.2.2 Complexing Often the ions to be determined are bound in complexes or are themselves complexing agents. Fluoride is an example for the latter case. Many metal ion species can form complexes with fluoride, which then can no longer be identified by an ion-selective electrode. Therefore, a TISAB solution is added to the sample solutions. The CDTA (Titriplex IV) contained in the TISAB solutions enters into a more stable complex with the metal ion than with the fluoride. The fluoride is displaced from the original complex and can then be measured. A ph value of 5-6 is used. This procedure has proved effective in most cases. However, for aluminium and iron fluoride complexes, it is often more advantageous to use TIRON (disodium catechol disulphonate) at a ph value of 6-7. The other case also occurs, namely that the ion to be determined is made unavailable for measurement by a complexing agent. To bring the ion into a form that is measurable, the complexing agent must be destroyed. For example, for the many different cyanide complexes, the metal ion can generally be liberated by treatment with acid. Dilute sulphuric acid will be sufficient for complexes that are easily dissociated. Those that are more difficult to dissociate must be treated with concentrated sulphuric acid. Caution: extremely toxic hydrocyanic acid is produced when cyanide complexes are dissociated. The dissociation reaction must therefore be performed under a good fume hood. 4.2.3 Cooking Certain interfering ions can be removed by cooking. These include the cyanide ion. The sample is made acidic and heated. Caution: In an acidic environment, the cyanide ions can form extremely toxic hydrocyanic acid. Therefore, removal of the cyanide must take place under a fume hood. Carbonates can also be removed by this method. A different procedure is used if large quantities of ammonium ion are an interference factor. Here, the ph value is raised to approximately 13. The heat causes the gaseous ammonia generated from the ammonium ion to escape. It has a biting odour (fume hood). Page 21 of 39