Basics of half cell and galvanic cell, types of electrodes, working of PH meter, application in Clinical lab,

1. Potentiometery and Ion selective
electrodes(ISE)
Muhammad Asif Shaheen
Lecturer Pathology
King Edward Medical University Lahore

2. • An electrode is an electrical conductor used to make contact with a
nonmetallic part of a circuit (e.g. a semiconductor, an electrolyte,
a vacuum or air).
• The anode is the electrode where oxidation (loss of electrons)
takes place; in a galvanic cell, it is the negative electrode
• The cathode is the electrode where reduction (gain of electrons)
takes place; in a galvanic cell, it is the positive electrode,
Terms to learn

3. • An ion is a charged atom or molecule. It is charged because the number
of electrons do not equal the number of protons in the atom or molecule.
An atom can acquire a positive charge or a negative charge depending
on whether the number of electrons in an atom is greater or less then the
number of protons in the atom.
• When an atom is attracted to another atom because it has an unequal
number of electrons and protons, the atom is called an ION.
• If the atom has more electrons than protons, it is a negative ion, or
ANION. If it has more protons than electrons,it is a positive ion CATION.

4. ANION CATION

5. • Cations (positively-charged ions) and anions (negatively-charged ions)
are formed when a metal loses electrons, and a nonmetal or metal gains
those electrons. The electrostatic attraction between the positives and
negatives brings the particles together and creates an ionic compound.
• An electrochemical cell is a device capable of either
generating electrical energy from chemical reactions or facilitating
chemical reactions through the introduction of electrical energy. A
common example of an electrochemical cell is a standard 1.5 – volt
cell meant for consumer use. This type of device is known as a
single galvanic cell.

6. • A galvanic cell, or voltaic cell, named after Luigi Galvani,
or Alessandro Volta respectively, is an electrochemical cell that
derives electrical energy from spontaneous redox reactions taking
place within the cell. It generally consists of two different metals
connected by a salt bridge, or individual half-cells separated by a
porous membrane.

7. • a half-cell consists of a solid metal (called an electrode) that is
submerged in a solution; the solution contains cations of the
electrode metal and anions to balance the charge of the cations.
Inside an isolated half-cell, there is an oxidation-reduction (redox)
reaction that is in chemical equilibrium.
• A galvanic cell consists of two half-cells, such that the electrode of
one half-cell is composed of metal A, and the electrode of the other
half-cell is composed of metal B

8. • In general, these two metals can react with each other
• In other words, the metal atoms of one half-cell are able to induce
reduction of the metal cations of the other half-cell; conversely
stated, the metal cations of one half-cell are able to oxidize the
metal atoms of the other half-cell. When metal B has a
greater electronegativity than metal A, then metal B tends to steal
electrons from metal A

9. This reaction between the metals can be controlled in a way that
allows for doing useful work:
The electrodes are connected with a metal wire in
order to conduct the electrons that participate in the reaction.
In one half-cell, dissolved metal-B cations combine with the free
electrons that are available at the interface between the solution
and the metal-B electrode; these cations are thereby neutralized,
causing them to precipitate from solution as deposits on the metal-
B electrode, a process known as plating.

10. • This reduction reaction causes the free electrons throughout the metal-B
electrode, the wire, and the metal-A electrode to be pulled into the metal-
B electrode. Consequently, electrons are wrestled away from some of the
atoms of the metal-A electrode, as though the metal-B cations were
reacting directly with them; those metal-A atoms become cations that
dissolve into the surrounding solution.
• As this reaction continues, the half-cell with the metal-A electrode
develops a positively charged solution (because the metal-A cations
dissolve into it), while the other half-cell develops a negatively charged
solution (because the metal-B cations precipitate out of it, leaving behind
the anions); unabated, this imbalance in charge would stop the reaction.

11. • The solutions are connected by a salt bridge or a porous plate in
order to conduct the ions (both the metal-A cations from one
solution, and the anions from the other solution), which balances
the charges of the solutions and thereby allows the reaction
between metal A and metal B to continue without opposition.

13. Half Cell
• A half-cell is a structure that contains a conductive electrode and a
surrounding conductive electrolyte in a layer
• Chemical reactions within this layer pump electric charges between
the electrode and the electrolyte, resulting in a potential
difference between the electrode and the electrolyte.
• The growing potential difference creates an intense electric
field within the double layer, and the potential rises in value until the
field halts the net charge-pumping reactions.
• In applications two dissimilar half-cells are appropriately connected
to constitute a Galvanic cell

14. Potentiometery
• Potentiometery is the measurement of an electrical potential difference
between two electrodes (half-cells) in an electro-chemical cell when the
cell current is zero (gal-vanic cell).
• Potentiometery is widely used clinically for the measurement of pH,
PCO2 and electrolytes (Na+, K+, CI`, Ca2+, Mg2+, Li+) in whole blood,
serum, plasma and urine
• The e/ectromotive force (E or EMF) is defined as the maximum difference
in potential between the two electrodes (right minus left) obtained when
the cell current is zero.

15. TYPES OF ELECTRODES
• Many different types of electrodes are used for potentiometric
applications. They include
1. Redox electrodes
2. Ion-selective membranes (glass and polymer)
3. PCO2 electrodes

16. Redox
• This potential reflects the tendency of the solution to release or
take up electrons; also called the Oxidation-Reduction Potential
(ORP).
• The Redox electrode tip will develop an electrical potential relative
to the reference electrode when both electrodes are immersed in
the same solution and connected to a high-impedance millivolt-
meter

17. • Types of Redox
• Inert Metal Electrodes
• Platinum and gold are examples of inert metals used to record the
redox potential of a redox couple dissolved in an electrolyte solution.
• Metal Electrodes Participating in Redox Reactions
• The silver-silver chloride electrode is an example of a metal
electrode that participates as a member of a redox couple. The
silver-silver chloride electrode consists of a silver wire or rod coated
with AgCI that is immersed in a chloride solution of constant activity;
this sets the half-cell potential.

18. Ion Selective Electrodes
• Membrane potentials are caused by the permeability of certain types of
membranes to selected anions or cations.
• Such membranes are used to fabricate ISEs that selectively interact with
a single ionic species.
• The potential produced at the membrane-sample solution interface is
proportional to the logarithm of the ionic activity or concentration of the
ion in question.
• Measurements with ISEs are simple, often rapid, nondestructive, and
applicable to a wide range of concentrations.

19. • The ion-selective membrane is the "heart of an ISE as it controls the
selectivity of the electrode. lon selective membranes are typically
composed of glass, crystalline, or polymeric materials.
• The chemical composition of the membrane is designed to achieve an
optimal perm-selectivity toward the ion of interest.
• In practice, other ions exhibit finite interaction with membrane sites and
will display some degree of interference for, determination of an analyte
ion.
• In clinical practice, if the interference exceeds an acceptable level, a
correction is required.
• Most ISEs used in clinical practice have sufficient selectivity and do not
require correction for interfering ions.

20. • Types of ISE
• Glass membrane
• Polymer membrane

21. Glass Membrane
• Glass membrane electrodes are employed to measure pH and Na+
,and as an internal transducer for PCO2 sensors.
• The H+ response of thin glass membranes was first demonstrated in
1906 by Cremer in the l930s, practical application of this phenomenon
for measurement of acidity in lemon juice was made possible by the
invention of the pH meter by Arnold Beckman.

22. • Glass electrode membranes are formulated from melts of silicon
and/ or aluminum oxide mixed with oxides of alkaline earth or alkali
metal cations. By varying the glass composition, electrodes with
selectivity for H+, Na+, K+, Li+, Rb+, CS+, Ag+, Tl+ and NH have
been demonstrated.
• However, glass electrodes for H+ and Na+ are today the only types
with sufficient selectivity over interfering ions to allow practical
application in clinical chemistry.

24. • Polymer Membrane Electrodes
• Polymer membrane ISEs are employed for monitoring pH and for
measuring electrolytes, including K+, Na+, CI`, Ca2+, Li+, Mg2+,
and CO (for total CO2 measurements).
• They are the predominant class of potentiometric electrodes used
in modern clinical analysis instruments.

26. pCo2 Electrode
• Electrodes have been developed to measure PCO2 in body fluids.
The first PCO2 electrode, developed in the 1950 by Stow and
Severinghaus, used a glass pH electrode as the internal element in
a potentiometric cell for measurement of the partial pressure of
carbon dioxide.
• This important development paved the way for commercial
availability of the three-channel blood analyzer (pH, PCO2, PO2) to
give the complete picture of the oxygenation and acid-'base status
of blood.

27. • . A thin membrane (~2Omm), permeable to only to gases and water vapor, is in
contact with the sample.
• Membranes of silicone rubber, Teflon, and other polymeric materials are
suitable for this purpose.
• On the opposite side of the membrane is a thin electrolyte layer consisting of a
weak bicarbonate salt (about 5mmol/L) and a chloride salt. A pH electrode and
Ag/AgCl reference electrode are in contact with this solution. The PCO2
electrode is a self-contained potentiometric cell. Carbon dioxide gas from the
sample or calibration matrix diffuses through the membrane and dissolves in
the internal electrolyte layer. Carbonic acid is formed and dissociates, shifting
the pH of the bicarbonate solution in the internal layer:
• CO2 +H2O H2CO3 H+ + HCO (12)
• And The relationship between the sample PCO2 and the signal generated by
the internal pH electrode is logarithmic and governed by the Nernst equation

30. • DIRECT POTENTIOMETRY BY ISE--UNITS OF MEASURE AND
REPORTING FOR CLINICAL APPLICATIONS
• Classical analytical methods such as flame photometry for the
measurement of electrolytes provide the total concentration Cc) of a
given ion in the sample, usually expressed in units of millimoles of ion
per liter of sample (mmol/L).
• Measurement of ions by direct Potentiometery provides yet another
unit of measurement known as activity (a), the concentration of free,
unbound ion in solution. Unlike methods sensitive to ion
concentration, ISEs do not sense the presence of complexed or
electrostatically hindered" ions in the sample.

31. • Physiologically, ionic activity is assumed to be more relevant than
concentration when considering chemical equilibria or biological
processes.
• Practically, however, ionic concentration is the more familiar term in
clinical practice, forming the basis of reference intervals and
medical decision levels for electrolytes.
• Early in the evolution of ISEs as practical tools in clinical chemistry,
it was decided that changing clinical reference intervals to a system
based on activity instead of concentration was impractical and
carried the risk for clinical misinterpretation.

32. PH Meter
• The basic components of a pH meter
• Indicator Electrode
• The pH electrode consists of a silver wire coated with AgCl, immersed
into an internal solution of 0.1 mmol/L HCl, and placed into a tube
containing a special glass membrane tip. This membrane is only
sensitive to hydrogen ions (H).
• Glass membranes that are selectively sensitive to H consist of
specific quantities of lithium, cesium, lanthanum, barium, or aluminum
oxides in silicate.

33. • When the pH electrode is placed into the test solution, movement of H near
the tip of the electrode produces a potential difference between the internal
solution and the test solution, which is measured as pH and read by a
voltmeter.
• The present concept of the selective mechanism that causes formation of
electromotive force at the glass surface is that an ion-exchange process is
involved. Cationic exchange occurs only in the gel layer—there is no
penetration of H through the glass.
• The specially formulated glass continually dissolves from the surface.
Although the glass is constantly dissolving, the process is slow, and the glass
tip generally lasts for several years. pH electrodes are highly selective for H;
however, other cations in high concentration interfere, the most common of
which is sodium.
• Electrode manufacturers should list the concentration of interfering cations
that may cause error in pH determinations.

34. • Reference Electrode
• The reference electrode commonly used is the calomel electrode.
• Calomel, a paste of predominantly mercurous chloride, is in direct contact
with metallic mercury in an electrolyte solution of potassium chloride.
• As long as the electrolyte concentration and the temperature remain
constant, a stable voltage is generated at the interface of the mercury and its
salt.
• A cable connected to the mercury leads to the voltmeter. The filling hole is
needed for adding potassium chloride solution. A tiny opening at the bottom
is required for completion of electric contact between the reference and
indicator electrodes.
• The liquid junction consists of a fiber or ceramic plug that allows a small flow
of electrolyte filling solution.

35. • Construction varies, but all reference electrodes must generate a
stable electrical potential.
• Reference electrodes generally consist of a metal and its salt in
contact with a solution containing the same anion.
• Mercury/mercurous chloride, as in this example, is a frequently used
reference electrode; the disadvantage is that it is slow to reach a new
stable voltage following temperature change and it is unstable above
80°C. Ag/AgCl is another common reference electrode. It can be used
at high temperatures, up to 275°C, and the AgCl-coated Ag wire
makes a more compact electrode than that of mercury

36. • Liquid Junctions
• Electrical connection between the indicator and reference electrodes is
achieved by allowing a slow flow of electrolyte from the tip of the
reference electrode.
• KCl is a commonly used.

37. • Readout Meter
• Electromotive force produced by the reference and indicator
electrodes is in the millivolt range.
• Zero potential for the cell indicates that each electrode half-cell is
generating the same voltage, assuming there is no liquid junction
potential.
• The isopotential is that potential at which a temperature change has
no effect on the response of the electrical cell. Manufacturers
generally achieve this by making midscale (pH 7.0) correspond to 0 V
at all temperatures.
• They use an internal buffer whose pH changes due to temperature
compensate for the changes in the internal and external reference
electrodes.

38. • pH Combination Electrode
• The most commonly used pH electrode has both the indicator and
reference electrodes combined in one small probe, which is
convenient when small samples are tested.
• It consists of an Ag/AgCl internal reference electrode sealed in a
narrow glass cylinder with a pH-sensitive glass tip. The reference
electrode is an Ag/AgCl wire wrapped around the indicator electrode.
• The outer glass envelope is filled with KCl and has a tiny pore near the
tip of the liquid junction. The solution to be measured must completely
cover the glass tip

41. Principle of ISE
• Ion Selective Electrodes (including the most common pH electrode)
work on the basic principal of the galvanic cell .
• By measuring the electric potential generated across a membrane
by "selected" ions, and comparing it to a reference electrode, a net
charge is determined.
• The strength of this charge is directly proportional to the
concentration of the selected ion. The basic formula is given for the
galvanic cell: Ecell = EISE - ERef

42. Advantages of Ion Selective Electrode
(ISE) Technique
• Relatively inexpensive and simple to use and have an extremely wide
range of applications and wide concentration range.
• They are particularly useful in biological/medical applications because they
measure the activity of the ion directly, rather than the concentration.
• ISEs are one of the few techniques which can measure both positive and
negative ions.
• They are unaffected by sample colour or turbidity.
• ISEs can be used in aqueous solutions over a wide temperature range.

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

44. LIMITATION
• Electrodes can be fouled by proteins or other organic solutes.
• Interference by other ions.
• Electrodes are fragile and have limited shelf life.
• Electrodes respond to the activity of uncomplexed ion. So ligands
must be absent.

45. APPLICATIONS
• 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.

46. • 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.
• 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.
• Ca in dairy products and beer.

47. Thanks