Electrochemical cell

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A demonstration electrochemical cell setup resembling the Daniell cell. The two half-cells are linked by a salt bridge carrying ions between them. Electrons flow in the external circuit.

An electrochemical cell is a device used for creating an electromotive force (voltage) and current from chemical reactions. The current is caused by the reactions releasing and accepting electrons at the different ends of a conductor. A common example of an electrochemical cell is a standard 1.5-volt battery.

[edit] Overview

Each half-cell consists of an electrode with atoms, and an electrolyte with ions that undergo either oxidation or reduction. In a full electrochemical cell, ions from the electrolyte of one half-cell lose electrons (oxidation) to their electrode while ions from the electrolyte of other half-cell gain electrons (reduction) from their electrode. If the atoms/ions involved in the electrode reactions are metal, then the same metal can be used for each electrode. If the atoms/ions involved in the reaction at each half-cell are not metal, then no electrode can be constructed out of those ions in atomic form; nonreactive metals such as platinum then can be used as a substitute electrode (as in the standard hydrogen electrode). Finally, a salt bridge is often employed to provide electrical contact between two half-cells with very different electrolytes—to prevent the solutions from mixing. This can simply be a strip of filter paper soaked in saturated potassium nitrate (V) solution.

The Bunsen cell, invented by Robert Bunsen.

Each half-cell has a characteristic voltage. Different choices of substances for each half-cell give different potential differences. Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions—when equilibrium is reached the cell cannot provide further voltage. In the half-cell which is undergoing oxidation, the closer the equilibrium lies to the ion/atom with the more positive oxidation state the more potential this reaction will provide. Similarly, in the reduction reaction, the further the equilibrium lies to the ion/atom with the more negative oxidation state the higher the potential.

This potential can be predicted quantitatively through the use of electrode potentials (the voltage measured when the substance is connected to hydrogen). The difference in voltage between electrode potentials gives a prediction for the potential measured. Spontaneity of a chemical reaction is determined by the overall cell potential Eo. If Eo>0, the reaction is spontaneous and if Eo<0, the reaction will not be spontaneous.

The potential window is the electrochemical voltage range between which a substance does not get oxidized or reduced.

[edit] Cell types