Electrochemical cell

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 generating an electromotive force (voltage) and current from chemical reactions, or the reverse, inducing a chemical reaction by a flow of current. 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. Batteries are composed of usually multiple Galvanic cells.

Half-cells

The Bunsen cell, invented by Robert Bunsen.

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, ions, atoms, or molecules from one half-cell lose electrons (oxidation) to their electrode while ions, atoms, or molecules from the other half-cell gain electrons (reduction) from their electrode. 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 solution. 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

The cell potential can be predicted through the use of electrode potentials (the voltages of each half-cell). (See table of standard electrode potentials). The difference in voltage between electrode potentials gives a prediction for the potential measured.

Cell potentials have a possible range of about zero to 6 volts. Cells using water-based electrolytes are usually limited to cell potentials less than about 2.5 volts, because the very powerful oxidising and reducing agents which would be required to produce a higher cell potential tend to react with the water.

Electrical cells

An electrical cell is a device that is used to generate electricity, or one that is used to make chemical reactions possible by applying electricity.

Cells producing electricity

An example is a Primary cell.

Cells using electricity

Some chemical reactions need high energy to happen. An example is the breakdown of water into hydrogen and oxygen in a process known as electrolysis.^[1] An electrochemical cell (or an electrolytic cell) is used for these reactions. Another example is the reduction of bauxite ore to make aluminum, which uses large cells and currents on the order of thousands of amperes.

Cell types

• Accumulator
• Concentration cell
• Electrolytic cell
• Galvanic cell
• Lasagne cell
• Lemon battery

See also

energy portal

• Alkaline battery
• Battery
• Cell notation
• Electrochemical potential
• Nickel Cadmium battery
• Activity (chemistry)

References

1. ^ "Electrolytic cell". Answers.com. http://www.answers.com/topic/electrolytic-cell-1. Retrieved on 2007-07-10.