Potentiometery and Ion selective electrodes(ISE)

Muhammad Asif Shaheen

Lecturer Pathology

King Edward Medical University Lahore

• 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

• An ion is a charged atom or molecule. It is charged because the numberof electrons do not equal the number of protons in the atom or molecule.An atom can acquire a positive charge or a negative charge dependingon whether the number of electrons in an atom is greater or less then thenumber of protons in the atom.

• When an atom is attracted to another atom because it has an unequalnumber of electrons and protons, the atom is called an ION.

• If the atom has more electrons than protons, it is a negative ion, orANION. If it has more protons than electrons,it is a positive ion CATION.

ANION CATION

• Cations (positively-charged ions) and anions (negatively-charged ions)are formed when a metal loses electrons, and a nonmetal or metal gainsthose electrons. The electrostatic attraction between the positives andnegatives brings the particles together and creates an ionic compound.

• An electrochemical cell is a device capable of eithergenerating electrical energy from chemical reactions or facilitatingchemical reactions through the introduction of electrical energy. Acommon example of an electrochemical cell is a standard 1.5 – voltcell meant for consumer use. This type of device is known as asingle galvanic cell.

• 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.

• 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

• 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

This reaction between the metals can be controlled in a way thatallows 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.

• This reduction reaction causes the free electrons throughout the metal-Belectrode, the wire, and the metal-A electrode to be pulled into the metal-B electrode. Consequently, electrons are wrestled away from some of theatoms of the metal-A electrode, as though the metal-B cations werereacting directly with them; those metal-A atoms become cations thatdissolve into the surrounding solution.

• As this reaction continues, the half-cell with the metal-A electrodedevelops a positively charged solution (because the metal-A cationsdissolve into it), while the other half-cell develops a negatively chargedsolution (because the metal-B cations precipitate out of it, leaving behindthe anions); unabated, this imbalance in charge would stop the reaction.

• 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.

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

Potentiometery

• Potentiometery is the measurement of an electrical potential differencebetween two electrodes (half-cells) in an electro-chemical cell when thecell 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 differencein potential between the two electrodes (right minus left) obtained whenthe cell current is zero.

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

Redox

• This potential reflects the tendency of the solution to release ortake up electrons; also called the Oxidation-Reduction Potential(ORP).

• The Redox electrode tip will develop an electrical potential relativeto the reference electrode when both electrodes are immersed inthe same solution and connected to a high-impedance millivolt-meter

• 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.

Ion Selective Electrodes

• Membrane potentials are caused by the permeability of certain types ofmembranes to selected anions or cations.

• Such membranes are used to fabricate ISEs that selectively interact witha single ionic species.

• The potential produced at the membrane-sample solution interface isproportional to the logarithm of the ionic activity or concentration of theion in question.

• Measurements with ISEs are simple, often rapid, nondestructive, andapplicable to a wide range of concentrations.

• The ion-selective membrane is the "heart of an ISE as it controls theselectivity of the electrode. lon selective membranes are typicallycomposed of glass, crystalline, or polymeric materials.

• The chemical composition of the membrane is designed to achieve anoptimal perm-selectivity toward the ion of interest.

• In practice, other ions exhibit finite interaction with membrane sites andwill display some degree of interference for, determination of an analyteion.

• In clinical practice, if the interference exceeds an acceptable level, acorrection is required.

• Most ISEs used in clinical practice have sufficient selectivity and do notrequire correction for interfering ions.

• Types of ISE

• Glass membrane

• Polymer membrane

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.

• 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.

• 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.

pCo2 Electrode

• Electrodes have been developed to measure PCO2 in body fluids.The first PCO2 electrode, developed in the 1950 by Stow andSeveringhaus, used a glass pH electrode as the internal element ina potentiometric cell for measurement of the partial pressure ofcarbon dioxide.

• This important development paved the way for commercialavailability of the three-channel blood analyzer (pH, PCO2, PO2) togive the complete picture of the oxygenation and acid-'base statusof blood.

• . A thin membrane (~2Omm), permeable to only to gases and water vapor, is incontact with the sample.

• Membranes of silicone rubber, Teflon, and other polymeric materials aresuitable for this purpose.

• On the opposite side of the membrane is a thin electrolyte layer consisting of aweak bicarbonate salt (about 5mmol/L) and a chloride salt. A pH electrode andAg/AgCl reference electrode are in contact with this solution. The PCO2electrode is a self-contained potentiometric cell. Carbon dioxide gas from thesample or calibration matrix diffuses through the membrane and dissolves inthe internal electrolyte layer. Carbonic acid is formed and dissociates, shiftingthe 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 bythe internal pH electrode is logarithmic and governed by the Nernst equation

• 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.

• Physiologically, ionic activity is assumed to be more relevant thanconcentration when considering chemical equilibria or biologicalprocesses.

• Practically, however, ionic concentration is the more familiar term inclinical practice, forming the basis of reference intervals andmedical 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 systembased on activity instead of concentration was impractical andcarried the risk for clinical misinterpretation.

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.

• When the pH electrode is placed into the test solution, movement of H nearthe tip of the electrode produces a potential difference between the internalsolution and the test solution, which is measured as pH and read by avoltmeter.

• The present concept of the selective mechanism that causes formation ofelectromotive force at the glass surface is that an ion-exchange process isinvolved. Cationic exchange occurs only in the gel layer—there is nopenetration 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 glasstip generally lasts for several years. pH electrodes are highly selective for H;however, other cations in high concentration interfere, the most common ofwhich is sodium.

• Electrode manufacturers should list the concentration of interfering cationsthat may cause error in pH determinations.

• Reference Electrode

• The reference electrode commonly used is the calomel electrode.

• Calomel, a paste of predominantly mercurous chloride, is in direct contactwith metallic mercury in an electrolyte solution of potassium chloride.

• As long as the electrolyte concentration and the temperature remainconstant, a stable voltage is generated at the interface of the mercury and itssalt.

• A cable connected to the mercury leads to the voltmeter. The filling hole isneeded for adding potassium chloride solution. A tiny opening at the bottomis required for completion of electric contact between the reference andindicator electrodes.

• The liquid junction consists of a fiber or ceramic plug that allows a small flowof electrolyte filling solution.

• 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

• 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.

• Readout Meter

• Electromotive force produced by the reference and indicatorelectrodes is in the millivolt range.

• Zero potential for the cell indicates that each electrode half-cell isgenerating the same voltage, assuming there is no liquid junctionpotential.

• The isopotential is that potential at which a temperature change hasno effect on the response of the electrical cell. Manufacturersgenerally achieve this by making midscale (pH 7.0) correspond to 0 Vat all temperatures.

• They use an internal buffer whose pH changes due to temperaturecompensate for the changes in the internal and external referenceelectrodes.

• 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

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

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.

Non-destructive: no consumption of analyte.

Non-contaminating.

Short response time: in sec. or min. useful in industrial applications

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.

APPLICATIONS

• Ion-selective electrodes are used in a wide variety of applications fordetermining the concentrations of various ions in aqueous solutions.The following is a list of some of the main areas in which ISEs havebeen used.

• Pollution Monitoring: CN, F, S, Cl, NO3 etc., in effluents, and naturalwaters.

• 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, brewingsolutions.

• F in drinking water and other drinks.

• 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 combustionproducts.

• 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.

Thanks