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Conductivity Cell

In theory, a conductivity measuring cell is formed by two 1-cm square surfaces spaced 1-cm apart. Cells of different physical configuration are characterized by their cell constant, K. This cell constant (K) is a function of the electrode areas, the distance between the electrodes and the electrical field pattern between the electrodes. The theoretical cell just described has a cell constant of K = 1.0. Often, for considerations having to do with sample volume or space, a cell's physical configuration is designed differently. Cells with constants of 1.0 cm-1 or greater normally have small, widely spaced electrodes. Cells with constants of K = 0. 1 or less normally have large closely spaced electrodes. Since K (cell constant) is a "factor" that reflects a particular cell's physical configuration, it must be multiplied by the observed conductance to obtain the actual conductivity reading.

For example, for an observed conductance reading of 200 ÁS using a cell with K = 0. 1, the conductivity value is 200 x 0. 1 = 20 ÁS/cm.

In a simplified approach, the cell constant is defined as the ratio of the distance between the electrodes, d, to the electrode area, A. This however neglects the existence of a fringe-field effect, which affects the electrode area by the amount AR. Therefore K = d/(A + AR). Because it is normally impossible to measure the fringe-field effect and the amount of AR to calculate the cell constant, K, the actual K of a specific cell is determined by a comparison measurement of a standard solution of known electrolytic conductivity.

The most commonly used standard solution for calibration is 0.01 M KCl. This solution has a conductivity of 1412 ÁS/cm at 25oC. In summary, the calibration of a conductivity probe is to compensate for the fact that:


See also: Conductivity cell, Corrosion cell definition, Daniell cell, Impressed current cathodic protection system, Natural corrosion cells, Rotating cylinder test cell, Sacrificial anode cathodic protection