potentiomentry and ion selective electrodes

1. POTENTIOMETRY
&
ION SELECTIVE ELECTRODE
DR. Animikh Ray
Dept of biochemistry
FMMC

2. 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.

3. Potentiometry
 In potentiometry, the potential of an electrochemical cell is
measured under static conditions because no current flows
while measuring a solution’s potential.
 Use of electrodes to measure voltage that provide chemical
information.
 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.

5. Potentiometric Analysis
• Based on potential
measurement of
electrochemical cells without
any appreciable current
• The use of electrodes to
measure voltages from
chemical reactions

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

7.  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.

8.  System Components
 Liquid Junction
 Reference electrode
 Indicator or measuring electrode
 Readout device (Potentiometer)

9.  Liquid junction – also known as a salt bridge 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.

10. Salt bridge
• Prevents mixing up of analyte
components

11. Reference electrode
• Half-cell with known potential (Eref)
• Left hand electrode (by convention)
• Easily assembled
• Rugged
• Insensitive to analyte concentration
▫ Reversible and obeys Nernst equation
▫ Constant potential
▫ Returns to original potential

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.
o 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. Saturated Calomel Electrode
Hg2Cl2(s)+2e-2Hg(l)+2Cl-(aq)
•Aka SCE
•Easy to prepare
•Easy to maintain
•0.2444 V at 25C
•Dependent on temp

14. Saturated calomel electrode

15. SHE
• Hydrogen Gas Electrode
• Pt (H2 (1 atm), H+ (1M)

16. Silver-silver chloride electrode
Standard hydrogen electrode (SHE)

17.  Indicator Electrode- also called the measuring
electrode
It is immersed in a solution of the analyte, develops
a potential, Eind that depends on the activity of the
analyte.
Is selective in its response
It is the other electrochemical half-cell that responds
to changes in the activity of a particular analyte
species in a solution.
Example:
Ion-Selective Electrodes

18. Indicator or measuring electrode
• The potential of this electrode is proportional to the
concentration of analyte.
• Two classes of indicator electrodes are used in
potentiometry:
o metallic electrodes
• Electrodes of the first kind
• Electrode of the second kind
• Redox electrode
o membrane electrodes (ion-selective electrodes)
• glass pH electrode

19. Metallic electrodes
Electrodes of the first kind
• A metal in contact with a solution containing its cation.
• The potential is a function of concentration of Mn+ in a
Mn+ / M. The most common ones:
o Silver electrode (dipping in a solution of AgNO3)
• Ag+ + e ↔ Ag
o Copper electrode
• Cu+2 + 2e ↔ Cu
o Zn electrode
• Zn+2 + 2e ↔ Zn

20. Electrodes of the First Kind
• Pure metal electrode in direct equilibrium with its cation
• Metal is in contact with a solution containing its cation.
M+n(aq) + ne-  M(s)

21. Disadvantages of First Kind Electrodes
• Not very selective
oAg+ interferes with Cu+2
• May be pH dependent
oZn and Cd dissolve in acidic solutions
• Easily oxidized (deaeration required)
• Non-reproducible response

22. Electrode of the second kind
• A metal wire coated with one of its salts precipitate.
• Respond to changes in ion activity through
formation of complex.
• A common example is silver electrode and AgCl as
its salt precipitate.
• This kind of electrode can be used to measure the
activity of chloride ion in a solution.

23. Redox electrode
• An inert metal is in contact with a solution containing
the soluble oxidized and reduced forms of the redox
half-reaction.
• The inert metal is usually is platinum (Pt).
• The potential of such an inert electrode is
determined by the ratio of the reduced and oxidized
species in the half-reaction.
• A very important example of this type is the
hydrogen electrode.

24. Ion selective electrode
 An ion-selective electrode (ISE), also known as a
specific ion electrode (SIE), is a transducer (or sensor)
that converts the activity of a specific ion dissolved in a
solution into an electrical potential, which can be
measured by a voltmeter or pH meter.
 indicator electrode based on determination of cations
or anions by the selective absorption of these ions to a
membrane surface.

25. Introduction
• Ion selective electrode (ISE) is an analytical
technique used to determine the activity of ions
in aqueous solution by measuring the electrical
potential.
• Specific ion dissolved in a solution create an
electrical potential, which can be measured by a
voltmeter or pH meter.
• The strength of this charge is directly
proportional to the concentration of the selected
ion.

27. Principle
• ISE consists of a thin membrane
• Only specific ion can be diffuse.
• By measuring the electric potential
generated across a membrane by
“selected” ions, and comparing it with
reference electrode.
• And net charge is determined.

28. Types of
ISE
• Glass membrane
• Solid state electrode
• Liquid based electrode
• Compound electrode

29. Glass Membrane Electrode
• This method uses the electrical potential
of pH-sensitive electrodes as a
measurement signal.
• The glass electrode is the most
commonly used sensor.

30. Solid State Electrode
• Electrode body of
Inorganic crystalline
polymer.
• E.g. Special Epoxide
Resin with excellent
mechanical
properties.
• High temperature
stability.

31. Liquid based
electrode
• Formed by a very thin layer
of an organic liquid.
• Membrane is like jelly
• Impermeable to water
• only to allow to pass certain
ion.
• Organic material
- Carbon tetrachloride
- Benzene
- Mesitylene

32. Compound
electrode
• Electrode have membrane of multiple type

33. Sample
Collection
• Serum
• Collected in heparin bulb
• Plain
• EDTAcan not be use for doing electrolyte
• EDTAis chelating agent & anti-coagulant.
• It chelat with all ions of blood
• So interfere with concentration of ions
• Urine
• Collected in plain vacuette
13

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

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

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

37.  Electrolyte
 Sodium
 Potassium
 Calcium
 Lithium
 Iodine
 Magnesium
 Chloride
 Fluoride
 Glucose
 Urea
 Arterial Blood Gas
Analysis
 pO2
 pCO2
 pH
 HCO3-
Application of
ISE
27

38. Routinely measured
electrolytes
Sodium
– (90%)Major cation
– Extracellular fluid outside cells
Normal values
– Serum = 135-145 mEq/L
– Urine (24 hr ) = 40-220 mEq/L
Functions
– Influence on regulation of body water
– Osmotic activity
– Central - Neuromuscular activity
15

39. Hyponatremia
16
 Hyponatremia <135 mEq/L
– Increased Na+ loss
– Causes
•Diabetes mellitus
•Diabetic Ketoacidosis
–- Because of diuresis
•Severe diarrhea & Severe Vomiting

40. Hypernatremia
• Excess water loss resulting in
dehydration (relative increase)
– Dehydration from inadequate
intake
– Dehydration due severe diarrhea
– Diabetes insipidus
– Burns
water
17

41. Potassium (K)
• (2%)major cation
• Intracellular fluid inside cell
Normal value
• Serum- 3.5-5.3 mEq/L
• Urine- 25-125 mEq/L
Function
Heart muscle contraction
Increase or Decrease K+ =Arrhythmiasis
18

42. • Hypokalemia = a low level
of potassium (K+) in the blood serum.
• Diarrhea
Medications like furosemide (diuretic)
• Dialysis
• Diabetes insipidus
• Hyperaldosteronism
Hypokalemi
a

43. Hyperkale
mia
• Increased K concentration
• Causes
– Acute Renal failure
– Chronic Renal failure
– Acidosis (Diabetes mellitus )
• H+ competes with K+ to get into cells & to be
excreted by kidneys
• Decreased insulin promotes cellular K loss
• Hyperosomolar plasma (from ↑ glucose) pulls
H2O and potassium into the plasma .
20

45. Chloride ( Cl - )
Chloride
Major cation
Extracellular fluid
Normal value
– Serum – 100 -110 mEq/L
– 24 hour urine – 110-250 mEq/L
varies with intake
– CSF – 120-132 mEq/L
22

46. Hypochloremia
Same as Hyponatremia
• congestive heart failure
• Severe diarrhea
• Severe vomiting
• drugs such as
•Laxatives
•diuretics
•corticosteroids
•Bicarbonates.
23

47. Hyperchlo
remia
• Same as Hypernatremia
• Increased serum Cl
– dehydration
– renal tubular disease
– metabolic acidosis
24

48. 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.

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

50. 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.