ELECTROCHEMICAL CELLS

Galvanic and electrolytic cells:

The difference between potential of the two metals results in a potential difference (also called a electromotive force or voltage, emf) between the two half-cells. That can be measured by means of a high impedance voltmeter which measures the driving force or voltage for reaction without allowing current to flow from which can be calculated thermodynamic data. Alternatively the reaction can be allowed to proceed by connecting the two half-cells by circuit (a wire or a resistor) and allowing the current to flow. These are both examples of galvanic cells, where the chemical reaction occurs. Electrons flow from the electrode with the most negative potential (the anode, where oxidation occurs) to that with the most positive potential (the cathode, where reduction occurs). The salt bridge (or porous glass frit) allows ions to transfer into each half-cell. That flow counteracts the imbalance of charge that would develop in each half cell as electrons (e-) pass from one electrode to the other, which would inhibit the reaction. The need for a frit or salt bridge is avoided if both half cells can share a common electrolyte. It is a special case, where all redox active ions in the solution react at one half cell electrode only and therefore do not have to be separated from the other electrode.

Fig. 1. Examples of (a) an electrolytic cell incorporating a salt bridge; (b) a galvanic cell incorporating a porous frit.

Other half-cell reactions:

The metal-insoluble salt electrode consists of a metal M coated with a porous insoluble salt MX in a solution of X-. A good example is the silver/silver chloride electrode (Fig. 2a) for which the half-cell reaction is, where the reduction of solid silver chloride produces solid silver and releases chloride ion into solution.

Fig. 2. (a) The silver/silver chloride half-cell; (b) the ferric (Fe3+)/ferrous (Fe2+) half-cell; (c) the

fluorine/fluoride ion half-cell.