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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. (see Water resistivity measurement)

For example, for an observed conductance reading of 200 µS using a cell with K = 0. 1, the conductivity value is 200 x 0. 1 = 20 µS/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 µS/cm at 25oC. In summary, the calibration of a conductivity probe is to compensate for the fact that:

• K is not specifically known
• K changes as the electrode ages.
• Calibration simply adjusts the measured reading to the true value at a specified temperature.