Silver Faraday constant measurement

In the 1930s, electrochemistry was a major part of physical chemistry and laboratory measurements were related to easily reproducible experiments. Thus plating out 1 mole of silver metal from a solution of AgNOa was an easy way to measure coulombs with an ammeter to measure current and a clock measuring seconds. The Faraday constant then requires fiufher definitions of an ampere etc., but those constants can be obtained through measurements and calculations from electroplating silver. Today, the modem values are all subjected to a least squares fit of all the known constants with the best experimental data except, as mentioned above, the value of c is now fixed and not subject to further measurement. The value of c is the kingpin of most of all the other constants. 

The quantity of electricity passed through a cell may be determined by measuring the current as a function of time, and determining the area under the current-time curve. A calibrated galvanometer with a short response time is used. Alternatively a chemical coulometer is commonly employed. This consists of an electrolytic cell in series with the experimental cell, the same amount of electricity therefore passing through both the chemical reaction at the cathode or anode (or both) of the coulometer must occur with 100% current efficiency, and should be easily and accurately estimated. The deposition of silver at the cathode of a silver coulometer (q.v.), the dissolution of silver from a silver anode (see Faraday constant) and the reaction 2e + l2 2r of the iodine coulometer (q.v.) all satisfy these conditions, and another common and convenient device is the copper coulometer (q.v.). 

The accepted reference method for determining chloride in blood serum, plasma, urine, sweat, and other body fluids is the coulometric titration procedure. In this technique, silver ions are generated coulometrically. The silver ions then react with chloride ions to form insoluble silver chloride. The end point is usually detected by amperometry (see Section 23B-4) when a sudden increase in current occurs on the generation of a slight excess of Ag. In principle, the absolute amount of Ag" needed to react quantitatively with Cl can be obtained from application of Faraday s law. In practice, calibration is used. First, the time required to titrate a chloride standard solution with a known number of moles of chloride (nci )s using a constant current I is measured. The same constant current is next used in the titration of the unknown solution, and the time r is measured. The number of moles of chloride in the unknown (ncr)u is then obtained as follows ... 

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