Electrochemical cell

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An electrochemical cell is a device capable of either deriving electrical energy from chemical reactions, or facilitating chemical reactions through the introduction of electrical energy. A common example of an electrochemical cell is a standard 1.5-volt "battery". (Actually a single "Galvanic cell"; a battery properly consists of multiple cells.)

Contents

Half-cells

An electrochemical cell consists of two half-cells. Each half-cell consists of an electrode, and an electrolyte. The two half-cells may use the same electrolyte, or they may use different electrolytes. The chemical reactions in the cell may involve the electrolyte, the electrodes or an external substance (as in fuel cells which may use hydrogen gas as a reactant). In a full electrochemical cell, species from one half-cell lose electrons (oxidation) to their electrode while species from the other half-cell gain electrons (reduction) from their electrode. A salt bridge (i.e. filter paper soaked in KNO3) is often employed to provide ionic contact between two half-cells with different electrolytes—to prevent the solutions from mixing and causing unwanted side reactions. As electrons flow from one half-cell to the other, a difference in charge is established. If no salt bridge were used, this charge difference would prevent further flow of electrons. A salt bridge allows the flow of ions to maintain a balance in charge between the oxidation and reduction vessels while keeping the contents of each separate. Other devices for achieving separation of solutions are porous pots and gelled solutions. A porous pot is used in the Bunsen cell (right).

Equilibrium reaction

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.

Electrode potential

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