1. POTENTIOMETRY

2. POTENTIOMETRY - Based on static (zero-current) measurements
- Used to obtain information on the composition of an analyte
- Potential between two electrodes is measured
Applications
- Environmental monitoring
- Clinical diagnostics (blood testing, electrolytes in blood)
- Control of reaction processes

3. Potentiometry
Potentiometry is a kind of electro analytical techniques in that the potential difference between two electrodes of electrochemical cell is measured under condition of zero current.
Concentration of ions in solution is calculated from the measured potential difference between the two electrodes immersed in solution under condition of zero current.
This type of system includes at least two electrodes, identified as an indicator electrode (right half-cell) and a reference electrode (left half-cell ) which act as the cathode and anode respectively.

4. Each electrode is in contact with either the sample (in the case of the “indicator electrode”) or a reference solution ( in the case of the “reference electrode”).
Electrodes + solution = electrochemical cell
Each electrode represent half cell reaction with correspondance half cell potential.

6. Potentiometer
A device for measuring the potential of an electrochemical cell without drawing a current or altering the cell’s composition.
The potential is measured under static conditions.
Because no current, or only a negligible current flows while measuring a solution’s potential, its composition remains unchanged.
For this reason, potentiometry is a useful quantitative method.

7. Two electrodes connected to Potentiometer to measure potential difference
Indicator electrode (Eind) – potential respond to change according to conc. of ions
Reference electrode (Eref) – half cell potential does not change.
Ecell = Eind ─ Eref + Elj

8. Example 1
Schematic diagram of an electrochemical cell of potentiometric measurement

9. System Components Liquid Junction Reference electrode
Indicator or measuring electrode
Readout device (Potentiometer)

10. Liquid junction – also known as a salt bridge are required to complete the circuit between the electrodes.
Functions:
It allows electrical contact between the two solutions.
It prevents the mixing of the electrode solutions.
It maintains the electrical neutrality in each half cell as ions flow into and out of the salt bridge.

11. Always treated as the left-hand electrode
Reference Electrode- is an electrochemical half-cell that is used as a fixed reference for the measurement of cell potentials.
A half-cell with an accurately known electrode potential, Eref, that is independent of the concentration of the analyte or any other ions in the solution
Always treated as the left-hand electrode
Examples:
Normal hydrogen electrode
Saturated calomel electrode
Ag-AgCl electrode

12. Reference Electrodes
Calomel electrode- composed of mercury/mercurous chloride; It is dependable but large, bulky, and affected by temperature.
Silver/silver chloride- Widely used because simple, inexpensive, very stable and non-toxic.
reference electrodes are more compact -- overall better & faster
Normal Hydrogen Electrode- consists of a platinized platinum electrode in HCl solution with hydrogen at atmospheric pressure bubbled over the platinum surface.
-determination of pH of the solution.

13. General Principles
Reference electrode | salt bridge | analyte solution | indicator electrode
Eref Ej Eind
Ecell = Eind – Eref + Ej
Reference cell :
a half cell having a known electrode potential
Indicator electrode:
has a potential that varies in a known way
with variations in the concentration of an analyte
A cell for potentiometric determinations.

14. Reference electrode 1) Saturated calomel electrode (S.C.E.)
: maintains a fixed potential :
a half cell having a known electrode potential
1) Saturated calomel electrode (S.C.E.)
Hg(l) | Hg2Cl2 (sat’d), KCl (sat’d) | |
electrode reaction in calomel hal-cell
Hg2Cl2 (s) + 2e = 2Hg(l) + 2Cl–
Eo = V
E = Eo – ( /2) log[Cl–]2 = V
Temperature dependent
A calomel electrode saturated with KCl is called a saturated calomel electrode, abbreviated S.C.E.
Advantage : using saturated KCl is that [Cl-] does not change if some liquid evaporates.

15. Hg2Cl2  Hg22+ + 2Cl– Ksp = 1.8 ×10–18 Saturated KCl = 4.6 M KCl
The crystal structure of calomel(Hg2Cl2), which has limited solubility in water (Ksp = 1.8 ×10–18).
Hg2Cl2  Hg Cl–
Ksp = 1.8 ×10–18
Saturated KCl = 4.6 M KCl

16. Fig. 21-2. Diagram of a typical commercial saturated calomel electrode.
Fig A saturated calomel electrode made from materials readily available in any laboratory.

17. 2) Silver-silver chloride electrode
Ag(s) | AgCl (sat’d), KCl (xM) | |
AgCl(s) + e = Ag(s) + Cl–
Eo = V
E = Eo – ( /1) log [Cl–]
E (saturated KCl) = V (25oC)

18. 3) standard hydrogen electrode (SHE)
The most fundamental reference electrode in electrochemistry. "By definition" its equilibrium potential is considered zero at any temperature, because this electrode was chosen as an arbitrary zero point for electrode potentials. A zero point is needed since the potential of a single electrode cannot be measured, only the difference of two electrode potentials is measurable. All electrode potentials are expressed on this "hydrogen scale." It is a hydrogen electrode with an electrolyte containing unit concentration of hydrogen ions and saturated with hydrogen gas at unit atmosphere pressure. This electrode can be somewhat inconvenient to use because of the need to supply hydrogen gas. Therefore, other reference electrodes (e.g., calomel or silver/silver chloride) are often used instead, but the measured electrode potentials can be converted to the "hydrogen scale." Also called "normal hydrogen electrode." Strictly speaking, one must use unit activity rather than concentration of hydrogen ions and unit fugacity rather than unit pressure of hydrogen gas.
Pt | H2(g, 1.0 atm)|H+(aq, A= 1.0M)
½ H2(g, 1.0 atm) = H+(aq, A= 1.0M) + e Eo = V

19. Voltage Conversions Between Different Reference Scales
If an electrode has a potential of – 0.461V with respect to a calomel electrode, what is the potential with respect to a silver-silver chloride electrode? What would be the potential with respect to the standard hydrogen electrode? 20

20. Liquid-junction potential
A potential difference between two solutions of different compositions separated by a membrane type separator. The simplest example is the case of two solutions containing the same salt in different concentrations. The salt will diffuse from the higher concentration side to the lower concentration side. However, the diffusion rate of the cation and the anion of the salt will very seldom be exactly the same (see mobility). Let us assume for this example that the cations move faster; consequently, an excess positive charge will accumulate on the low concentration side, while an excess negative charge will accumulate on the high concentration side of the junction due to the slow moving anions. This sets up a potential difference that will start an electromigration of the ions that will increase the net flux of the anions and decrease the net flux of the cations. In steady-sate conditions, the two ions will move at the same speed and a potential difference will be created between the two solutions. This "steady-sate" potential difference seems constant, but this is misleading because it slowly changes as the concentrations between the two solutions equalize. The diffusion process will "eventually" result in equal concentrations of the salt in the two solutions separated by the membrane, and the liquid-junction potential will vanish. For a simple case, the value of the liquid junction potential can be calculated by the so called "Henderson" equation.

21. Junction potential :
a small potential that exists at the interface between two electrolyte solutions that differ in composition.
Development of the junction potential caused by unequal mobilities of ions.
Mobilties of ions in water at 25oC:
Na+ : 5.19 × 10 –8 m2/sV K+ : 7.62 × 10 –8 Cl– : 7.91× 10 –8

22. Fig Schematic representation of a liquid junction showing the source of the junction potential, Ej. The length of the arrows corresponds to the relative mobilities of the ions.
Fig 21-4 Diagram of a silver/silver chloride electrode showing the parts of the electrode that produce the reference electrode potential Eref and the juction potential Ej

24. Liquid junction potential
Cells without liquid junction
Pt/H2(g), HCl/AgCl/Ag
Rare to have this type of cell
Cells with liquid junction
Glass frit
Salt bridge
Develop a potential by differential migration rates of the cation and anion.
Junction potential
HCl(0.1)/HCl(0.01) Ej = 40 mV (H+ faster than Cl– )
KCl(0.1)/KCl(0.01) Ej = –1.0 mV (K+ slower than Cl– )
Usually experimentally determine instrument response

25. Indicator electrodes
Metallic indicator electrode responds to analyte activity.
Electrode of the first type
Direct equilibrium with analyte
Ag for Ag+, Au for Au3+, etc
Potential described by Nernst equation.
As [M] , E
Note potential linearly related to log of the concentration !
Remember - indicator BY DEFINITION cathode
measurement theoretically under zero-current (steady state)
Electrode of the second type
Indirect equilibrium with analyte
M/MX/X–
Silver/Silver chloride for chloride
also Nernstian response
as [X–] , E
Inert Metallic electrode for Redox systems
Provides a surface for the electrochemistry to occur
Pt, Au, Pd, C
Xn+(aq) + ne = X(s)
Eind = Eo – (0.5916/n) log (1/[Xn+])
A plot of Equation 21-3 for an electrode of the first kind.
AgCl(s) + e = Ag(s) + Cl–(aq)
Eind = Eo – log [Cl–]
A plot of Equation 21-4 for an electrode of the second kind for Cl–.

26. Indicator electrodes
Indicator electrodes for potentiometric measurements are of two basic types, namely, metallic and membrane.
1) Metallic indicator electrodes :
develop a potential that is determined by the equilibrium position of redox half-reaction at the electrode surface.
First-order electrodes for cations :
A first order electrode is comprised of a metal immersed in a solution of its ions, such as silver wire dipping into a silver nitrate solution. Only a few metals such as silver, copper, mercury, lead , zinc, bismuth, cadmium and tin exhibit reversible half- reactions with their ions and are suitable for use as first order electrodes. Other metals, including iron, nickel, cobalt, tungsten, and chromium, develop nonreproducible potentials that are influenced by impurities and crystal irregularities in the solid and by oxide coatings on their surfaces. This nonreproducible behavior makes them unsatisfactory as first-order electrodes.

27. Example of first order metallic indicator electrodes
Use of Ag and calomel electrode to measure [Ag+]
The outer compartment of the electrode is filled with KNO3, so there is no direct contact between Cl- in the inner compartment and Ag+ in the beaker
The reaction at the Ag indicator electrode is
The calomel reference half-cell reaction is
Ag+ + e- ⇌ Ag(s) Eo+ = V
Hg2Cl2(s) + 2e- ⇌ 2Hg(l) + 2 Cl - E_ = V
The reference potential is fixed at 0.241V
because the reference cell is saturated with KCl.
The Nernst equation for the entire cell is therefore
Potential of Ag | Ag+ Indicator electrode
Potential of S.C.E reference electrode
[Ag+]
E = E+ - E– = { log ( )} - {0.241} 1
E = – log (1/[Ag+])
Ideally, the voltage changes by 59.16mV (at 25℃)
for each factor-of-10 change in [Ag+] 28

28. Second-order electrodes for anions
A metal electrode can sometimes be indirectly responsive to the concentration of an anion that forms a precipitate or complex ion with cations of the metal.
Ex Silver electrode
The potential of a silver electrode will accurately reflect the concentration of
iodide ion in a solution that is saturated with silver iodide.
AgI(s) + e = Ag(s) + I– Eo = – 0.151V
E = – – ( /1) log [I–]
= – ( /1)pI
2. Mercury electrode for measuring the concentration of the EDTA anion Y4–.
Mercury electrode responds in the presence of a small concentration of the
stable EDTA complex of mercury(II).
HgY2– + 2e = Hg(l) + Y4– Eo = 0.21V
E = – ( /2) log ([Y4–] /[HgY2–])
K = 0.21 – ( /2) log (1 /[HgY2–])
E = K – ( /2) log [Y4–] = K +( / 2) pY

29. Inert electrodes
Chemically inert conductors such as gold, platinum, or carbon that do not participate, directly, in the redox process are called inert electrodes. The potential developed at an inert electrode depends on the nature and concent-ration of the various redox reagents in the solution.
Ag(s) | AgCl[sat’d], KCl[xM] | | Fe2+,Fe3+) | Pt
Fe3++e = Fe2+ Eo = V
Ecell = Eindicator – Ereference
= {0.770 – ( /1) log [Fe2+]/[Fe3+]} – {0.222 – ( /1) log [Cl–]}

30. Ion-Selective Electrodes (ISE)
Introduction
An Ion-Selective Electrode (ISE) produces a potential that is proportional to the concentration of an analyte. Making measurements with an ISE is therefore a form of potentiometry. The most common ISE is the pH electrode, which contains a thin glass membrane that responds to the H+ concentration in a solution.
Theory
The potential difference across an ion-sensitive membrane is:
E = K – (2.303RT/nF)log(a)
where K is a constant to account for all other potentials, R is the gas constant, T is temperature, n is the number of electrons transferred, F is Faraday's constant, and a is the activity of the analyte ion. A plot of measured potential versus log(a) will therefore give a straight line.
ISEs are susceptible to several interferences. Samples and standards are therefore diluted 1:1 with total ionic strength adjuster and buffer (TISAB). The TISAB consists of 1 M NaCl to adjust the ionic strength, acetic acid/acetate buffer to control pH, and a metal complexing agent.

31. ISE (ion selective electrode)
Any electrode that preferentially responds to one ion species.
1. Glass membrane electrode : H+
2. Liquid membrane electrodes
3. Solid state and precipitate electrodes
4. Gas sensing electrodes
5. Enzyme electrodes
Selectivity coefficient
kX,Y = (response to Y) / (response to X)
General behavior of ISE
E = constant  ( /nX) log [AX +(kX,Y AY nX nY )]

32. Instrumentation
ISEs consist of the ion-selective membrane, an internal reference electrode, an external reference electrode, and a voltmeter. A typical meter is shown in the document on the pH meter. Commercial ISEs often combine the two electrodes into one unit that are then attached to a pH meter.
Schematic of an ISE measurement
Picture of a commercial fluoride ISE

33. Membrane indicator electrodes
 Glass membrane pH electrodes
The internal element consists of silver-silver chloride electrode immersed in a pH 7 buffer saturated with silver chloride. The thin, ion-selective glass membrane is fused to the bottom of a sturdy, nonresponsive glass tube so that the entire membrane can be submerged during measurements. When placed in a solution containing hydrogen ions, this electrode can be represented by the half-cell :
Ag(s) | AgCl[sat’d], Cl–(inside), H+(inside) | glass membrane | H+(outside)
E = Eo – ( /1) log [Cl–] + ( /1) log ([H+(outside)]/[H+(inside)])
E = Q + ( /1) log [H+(outside)]

34. Typical electrode system for measuring pH
Typical electrode system for measuring pH. (a) Glass electrode (indicator) and saturated calomel electrode (reference) immersed in a solution of unknown pH. (b) Combination probe consisting of both an indicator glass electrode and a silver/silver chloride reference. A second silver/silver chloride electrode serves as the internal reference for the glass electrode. The two electrodes are arranged concentrically with the internal reference in the center and the external reference outside. The reference makes contact with the analyte solution through the glass frit or other suitable porous medium. Combination probes are the most common configuration of glass electrode and reference for measuring pH.

35. pH Meters

36. / pH meter : A glass combination electrode E = K – b (0.05916) log ( A
out )
K : Asymmetry potential
b : electromotive efficiency ( close to 1.00)
A : Activity of hydrogen ion

37. pH Measurement with a Glass Electrode
Glass electrode is most common ion-selective electrode
Combination electrode incorporates both glass and reference electrode in one body
Ag(s)|AgCl(s)|Cl-(aq)||H+(aq,outside) H+(aq,inside),Cl-(aq)|AgCl(s)|Ag(s)
Outer reference
electrode
[H+] outside
(analyte solution)
[H+] inside
Inner reference
electrode
Glass membrane
Selectively binds H+
Electric potential is generated by [H+] difference across glass membrane

38. Theory of the glass membrane potential:
The pH electrode functions as a result of ion exchange on the surface of a hydrated layer. The membrane of a pH glass electrode consists of chemically bonded Na2O and SiO2. For the electrode to become operative, it must be soaked in water. During this process, the outer surface of the membrane becomes hydrated. When it is so, the sodium ions are exchanged for protons in the solution:
SiO─ Na+ (solid) + H+ (solution) ↔ SiO─H+ (solid) + Na+ (solution)
The protons are free to move and exchange with other ions.
Charge is slowly carried by migration of Na+ across glass membrane
Potential is determined by external [H+]

39. Composition of glass membranes 70% SiO2 30% CaO, BaO, Li2O, Na2O,
and/or Al2O3
Ion exchange process at glass membrane-solution interface:
Gl– + H+ = H+Gl–
(a) Cross-sectional view of a silicate glass structure. In addition to the three Si│O bonds shown, each silicon is bonded to an additional oxygen atom, either above or below the plane of the paper. (b) Model showing three-dimensional structure of amorphous silica with Na+ ion (large dark blue) and several H+ ions small dark blue incorporated.

40. GLASS MEMBRANE ELECTRODE
• Responsive to univalent cations ( H+ , Na+)
• Glass electrodes available for Na+, K+, NH4+, Li+, Ag+(cations only) by varying glass composition
• The selectivity for this cation by varying the composition of a thin ion sensitive glass membrane.
Glass membrane manufactured from SiO2 with negatively charged oxygen atom.
• Inside the glass bulb, a dilute HCl solution and silver wire coated with a layer of silver chloride.
The electrode is immersed in the solution and pH is measured
• Example: pH electrode

41. Potential of the glass electrode
The potential difference across the glass pH electrode depends on the activity of H+ on each side of the glass membrane.
Em = V1 – V2 = (RT/F)lna1 – (RT/F)lna2
= Easym log(a1 / a2)
if A1 = constant,
Em = K loga1
= K – pH
Standardization
at pH=7.00 , E = 0 V.
pH 4.00, E= mV/pH unit
Potential profile across a glass membrane from the analyte solution to the internal reference solution. The reference electrode potentials are not shown.

42. Glass pH electrode
Advantages over other electrodes for pH measurements:
Its potential is essentially not affected by the presence of oxidizing or reducing agents.
It operates over a wide pH range.
It responds fast and functions well in physiological systems.
Selective for monovalent cations only because polyvalent ions can not penetrate the surface of membrane.

43. Calibrating a glass electrode 
Always keep the electrodes in distilled water, saturated KCl solution(3.7M) or buffer when not in use.
2. Power ON. Switch to “STANDBY” : allow to warm for 30 min.
3. Rinse the electrode thoroughly with distilled water and then with pH 7.00 buffer solution.
Blot with clean tissue.
4. Determine the temperature of the buffer solution with a thermometer.
Adjust “TEMPERATURE” knob on the unit to the temperature.
5. Place the electrode in pH 7.00 (isopotential point) buffer solution.
Rotate the selector switch to “pH”. Wait for a stable display.
By using “CALIBRATION” knob, set the meter to the pH value of the buffer at its measured
temperature. Switch to “STANDBY”
6. Rinse the electrode thoroughly with distilled water and then with pH 4.00 buffer solution.
7. Place the electrode in pH buffer solution. Rotate the selector switch to “pH”.
Wait for a stable display. Using “SLOPE” knob, set the meter to the pH value of the buffer at its
measured temperature. Switch to “STANDBY”

44. Calibration of the Meters with pH 7 and pH 2 Buffers
1. Select the pH Mode and set the temperature control knob to 25°C. Adjust the cal 2 knob to read 100%.
2. Rinse the electrode with deionized water and blot dry using a piece of tissue (Shurwipes or Kimwipes are available in the labs).
3. Place the electrode in the solution of pH 7 buffer, allow the display to stabilize and, then, set the display to read 7 by adjusting cal 1. Remove the electrode from the buffer.
4. Rinse the electrode with deionized water and blot dry using a piece of tissue (Shurwipes or Kimwipes are available in the labs).
5. Place the electrode in the solution of pH 2 buffer, allow the display to stabilize and, then, set the display to read 2 by adjusting cal 2. Remove the electrode from the buffer.
6. Rinse the electrode with deionized water and blot dry using a piece of tissue (Shurwipes or Kimwipes, as before).
NOTE - Buffer solution are made available to you in individually labeled 2 oz. bottles. The buffers are to be used in these containers, only! Do not pour them into other containers at any time. After use, cap the bottles so that the buffers can be re-used.

45. Measuring pH
1. Make sure that the meter is set to the pH Mode and adjust the temperature to 25°C.
2. Place the electrode in the sample to be tested.
3. The pH of the solution appears in the display.
NOTE: Allow the display to stabilize before taking your reading!
4. Rinse the pH electrode and place it back in the storage solution.

46. Errors that affect pH measurements with glass electrode
1. The alkaline(sodium) error : low readings at pH values greater than 9
2. The acid error : somewhat high when the pH is less than about 0.5
3. Dehydration may cause erratic electrode performance.
4. Variation in junction potential : ~ 0.01 pH unit
5. Error in the pH of the standard buffer :  0.01 pH unit
Cleaning glass electrode :
1. Washing with 6M HCl
w/w% aqueous ammonium bifluoride (NH4HF2)

47. LIQUID MEMBRANE ELECTRODE
Liquid membrane is a type of ISE based on water- immiscible liquid substances produced in a polymeric membrane used for direct potentiometric measurement.
Used for direct measurement of several polyvalent cations (Ca ion) as well as a certain anions.
Inner compartment of electrode contains reference electrode & aqueous reference solution.
Outer compartment – organic liquid ion exchanger

48. •The polymeric membrane made of PVC to separate the test solution from its inner compartment which contains standard solution of the target ion. •The filling solution contains a chloride salt for establishing the potential of the internal Ag/AgCl wire electrode.

49. Liquid ISE
Ca ISE
Calcium didecylphosphate dissolved in dioctylphenylphosphonate
[(CH3(CH3)8CH2O)2PO2]2Ca  2[(CH3(CH3)8CH2O)2PO2]– + Ca2+
Diagram of a liquid-membrane electrode for Ca2+.

51. Comparison of a liquid-membrane calcium ion electrode with a glass pH electrode.

53. Photograph of a potassium liquid-ion exchanger microelectrode with 125 m of ion exchanger inside the tip. The magnification of the original photo was 400×.

54. A homemade liquid-membrane electrode.

55. Solid state crystalline membrane electrode
Migration of F– through LaF3 doped with EuF2.

57. Example: Fluoride (F-) electrode
Internal ref electrode
Ag/AgCl
Filling soln.
Aqueous NaCl + NaF
Membrane
LaF3 crystal disc
Applications
Electroplating industry, water treatment (fluoridation), toothpaste

58. GAS SENSING ELECTRODE
Available for the measurement of ammonia, carbon dioxide and nitrogen oxide.
This type of electrode consist of permeable membrane and an internal buffer solution.
The pH of the buffer changes as the gas react with it.
The change is detected by a combination pH sensor.
This type of electrode does not require an external reference electrode.

59. pco2 electrode Measurement of PCO2 in routine blood gases
A modified pH electrode with a CO2 permeable membrane covering the glass membrane surface
A bicarbonate buffer separates the membranes
Change in pH is proportional to the concentration of dissolved CO2 in the blood

60. Applications of ion selective electrodes
Ion-selective electrodes are used in a wide variety of applications for determining the concentrations of various ions in aqueous solutions. The following is a list of some of the main areas in which ISEs have been used.
Pollution Monitoring: CN, F, S, Cl, NO3 etc., in effluents, and natural waters.
Agriculture: NO3, Cl, NH4, K, Ca, I, CN in soils, plant material, fertilisers and feedstuffs.
Food Processing: NO3, NO2 in meat preservatives.
Salt content of meat, fish, dairy products, fruit juices, brewing solutions.
F in drinking water and other drinks.
Ca in dairy products and beer.
K in fruit juices and wine making.
Corrosive effect of NO3 in canned foods.
Detergent Manufacture: Ca, Ba, F for studying effects on water quality.
Paper Manufacture: S and Cl in pulping and recovery-cycle liquors.
Explosives: F, Cl, NO3 in explosive materials and combustion products.
Electroplating: F and Cl in etching baths; S in anodising baths.
Biomedical Laboratories: Ca, K, Cl in body fluids (blood, plasma, serum, sweat).
F in skeletal and dental studies.
Education and Research: Wide range of applications.

61. Application of Potentiometric Measurement
Clinical Chemistry
Ion-selective electrodes are important sensors for clinical samples because of their selectivity for analytes.
The most common analytes are electrolytes, such as Na+, K+, Ca2+,H+, and Cl-, and dissolved gases such as CO2.
Environmental Chemistry
For the analysis of of CN-, F-, NH3, and NO3- in water and wastewater.

62. Potentiometric Titrations
pH electrode used to monitor the change in pH during the titration.
For determining the equivalence point of an acid–base titration.
Possible for acid–base, redox, and precipitation titrations, as well as for titrations in aqueous and nonaqueous solvents.
Agriculture
NO3, NH4, Cl, K, Ca, I, CN in soils, plant material, fertilizers.
Detergent Manufacture
Ca, Ba, F for studying effects on water quality

63. Food Processing NO3, NO2 in meat preservatives
Salt content of meat, fish, dairy products, fruit juices, brewing solutions.
F in drinking water and other drinks.
Ca in dairy products and beer.
K in fruit juices and wine making.
Corrosive effect of NO3 in canned foods.

64. advantages
Relatively inexpensive and simple to use and have an extremely wide range of applications and wide concentration range.
Under the most favourable conditions, when measuring ions in relatively dilute aqueous solutions and where interfering ions are not a problem, they can be used very rapidly and easily.
ISEs can measure both positive and negative ions.
They are unaffected by sample colour or turbidity.

65. Non-destructive: no consumption of analyte.
Non-contaminating.
Short response time: in sec. or min. useful in industrial applications.

66. LIMITATION Precision is rarely better than 1%.
• Electrodes can be affected by proteins or other organic solutes.
• Interference by other ions.
• Electrodes are fragile and have limited shelf life.