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Daniele Naviglio » 21.Potentiometry


Potentiometry generally refers to all the analytical methods, based on the measurement of electrochemistry potential of a galvanic cell in the ‘absence of current’ state. Instrumentation typically used in potentiometry includes the base electrode, with a known potential, constant over time and indipendent of the composition of the solution containing the analyte in which it is immersed, and an indicator (or working) electrode, whose response depends on the concentration of the analyte, and finally an instrument for measuring potential, which may be a Poggendorf potentiometer or a modern electronic voltmeter. Potentiometry together with conducibility, coulometry, electrogravimetry, polarography, amperometry, etc. are part of electrochemical techniques or electroanalyticals. Electrochemistry is the branch of chemistry that studies the processes involving the transfer of electrons: the reactions of oxide reduction (commonly known as redox). Thus it deals with the chemical transformations produced by the passage of electricity in determinate chemical systems, and the production/storage of electricity by means of chemical transformations. In addition to its numerous, state-of-the-art electro-chemical applications, it provides the means to investigate phenomena like metal corrosion (refining), and most biochemical reactions which govern the functions of living, organisms such as photosyntesis, cellular respiration, the transmission of nerve impulses, etc. On the other hand, electrosynthesis permits the use of electrical energy for synthetic purposes, electroanalytical chemistry permits the exploitation of electrochemical principles applied to chemical analysis. Electrochemistry also studies all phenomena and possible applications of the conduction of electricity on the part of electrolytes.

Source: Wikipedia.

Electrochemical cells

Reference has been made above to oxide reduction reactions; a unique aspect of such  reactions is the transfer of electrons, thus an identical final reaction may frequently be conducted in an electrochemical cell, in which the oxidizing agent and reducing agent are held physically separate. A salt bridge isolates the reactants but promotes an electrical contact between the two semi cells. An external metallic conductor connnects the two metals.

Diagram of an electrochemical cell

In this cell the metallic copper is reduced, zinc ions are oxidized, and the electrons flow through the external circuit towards the copper electrode. The voltmeter measures the potential difference between the two metals in each instance, corresponding to the measurement of tendency of the net cell reaction to reach the equilibrium point. When the reaction proceeds, this tendency, and thus the potential, decreases continuously down to zero, the point at which equilibrium of the complex reaction is reached:

Zn + Cu2+ → Zn2+ + 2Cu

Diagram of an electrochemical cell. Source: Moterma

Diagram of an electrochemical cell. Source: Moterma

The pH meter

The pH meter is a cell consisting of a glass indicator electrode and a satured calomel reference electrode, immersed in the solution whose pH value needs to be evaluated.

Reference electrode→ Is a half-cell with a precisely known Erif electrode potential, independent of the analyte concentration, or of any other ion in the solution being examined. Conventionally the reference electrode is always treated as anode.

Indicator electrode→ Immersed in the analyte solution, develops a potential, Eind, which depends on the analyte activity

The salt bridge→ Prevent the analyte solution components from mixing with those of the reference electrode.

Diagram of a pH meter. Source: Unisdo

Diagram of a pH meter. Source: Unisdo

PH meter structure

Indicator electrode: consists of a thin glass membrane, sensitive to pH, soldered to the ends of a reinforced glass tube. This latter contains a small volume of cloridric acid, diluted and satured with silver chloride. In this solution a silver thread forms a silver/silver chloride reference electrode, which is linked to one of the terminals of a potential measuring instrument.

Reference electrode: the calomel electrode (SCE) has a known potential, which is constant and completely indifferent to the composition of the analyte solution.

Reference electrodes

The ideal reference electrode:

  • must have a known potential which is constant, and completely insensitive to the composition of the solution containing the analyte;
  • must be robust and easy to assemble;
  • must maintain a constant potential, even after small waves of current through the cell.

Indicator electrodes

The ideal indicator electrode must respond in a way which is rapid, and repeatable, to variations in concentration of single analyte ions, or a group of ions. The indicator electrodes are chosen on the basis of the titration to be carried out.

  • titration with formation of a complex or precipitate: the electrode must consist of the same element as the cation in the reaction;
  • acid-base titration: hydrogen or “glass” electrodes must be used;
  • Redox titration: the materials used must be inert to chemical reaction, eg. platinum or gold.
Diagram of an electrochemical cell. Source: Uniba

Diagram of an electrochemical cell. Source: Uniba


The potential of a cell is given by the equation:

Ecella = Eind – Erif + Ej
Eind: contains the required information regarding the concentration of the analyte;
Erif: is the potential of the reference electrode;
Ej:  is the potential which develops through the liquid junction of the salt bridge.

The potential Erif is measured by silver/silver chloride and calomel reference electrodes;
Whereas Ej is the potential junction of salt bridge.
Eind,also known as interphase, is the most important and varies with the pH of the analyte solution. It corresponds to the difference between the two potentials which develops on the two surfaces of the glass membrane of the pH-meter probe. This difference is linked to the differing concentrations of hydrogen ions, on the inside and outside the membrane.

Basic parts of the pH meter

A typical pH-meter consists of a probe (glass electrode) connected to an electronic device that records the signal from the probe, calculates the corresponding pH value and displays it. Frequently, two probes are immersed in the solution: in addition to the electrode, a temperature probe is also immersed, in order to correct the electrode reading according to the actual sample temperature. The meter circuit is simply a voltmeter that displays measurements in pH units instead of volts. Before use, the pH-meter should be calibrated, with at least two standard buffer solutions that span the range of pH values to be measured. For general purposes buffers at pH 7,01 and pH 10 are acceptable; for more precise measurements, a three buffer solution pH 4,01 calibration is preferred. Once calibration is complete, the electrode must be rinsed with distilled water, dried and immersed in the sample. A glass electrode is generally kept immersed in a solution of pH 3 to prevent the glass membrane from drying out; distilled water is best avoided for this purpose, as the hydrogen ions inside the electrode may be extracted by osmosis.

pH-meter. Source: Geass

pH-meter. Source: Geass

Basic parts of pH meter (cont.)

The pH-meter is an instrument used for rapidly measuring the hydrogen ion concentration of solutions, and then reporting the value in as accurate a way as possible. It is thus useful for checking the acidity of a variety of liquid and semi-solid foodstuffs. Finally, it is a useful instrument in acid-base titration, both when a visual indicator cannot be used, and in cases where the pH leap is smaller than three units.




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