A B C D E F G H I J K L M N O P Q R S T U V W X Y Z

ELECTROLYSIS

DOI: 10.1615/AtoZ.e.electrolysis

Electrolysis is the process which occurs in an electrochemical cell when electrons pass from the anode to the cathode via an external circuit connecting the electrodes. For this to occur, there must be a mechanism for charge transport in the electrolyte between the electrodes. This charge transport is due to the movement of ions in the electrolyte. These ions, on approaching the electrodes, undergo electrolysis and either accept electrons from the cathode or release electrons to the anode. In this way chemical change occurs at both electrodes.

The extent of chemical change depends on the charge passed, according to Faraday’s Laws of electrolysis. The rate of chemical change depends upon current density, j = I/S (Am−2), at the electrode and is given by;

(1)

where n is the number of electrons in the electrode reaction and F is the Faraday constant. The rate of chemical change and thus current density is determined by the kinetics of the electrode reaction(s).

Under conditions of electrolysis, the cell is operating away from its equilibrium (reversible) potentials determined from Thermodynamics. Certain electrode reactions are very fast and depart very little from the equilibrium potential. Such reactions are frequently referred to as reversible (see Figure 1). Other electrode reactions are inherently slow and require a potential, E, significantly greater in magnitude than the equilibrium potential to achieve a reasonable current density. This potential is called the overpotential, η (=E–Ee), and the electrode is the said to be polarized. Such reactions are referred to as irreversible (Figure 2). As overpotential is increased in magnitude (more negative for cathodic processes, more positive for anodic processes) current density increases, typically exponentially at high overpotentials. The relationship between current density and electrode potential is the subject of electrode kinetics.

Current potential curve at a polarized electrode.

Figure 1. Current potential curve at a polarized electrode.

Electrolysis of NaCl solution to form chlorine and sodium hydroxide. Other examples of electrolysis, or more specifically “half-cell reactions,” include:

Figure 2. Electrolysis of NaCl solution to form chlorine and sodium hydroxide. Other examples of electrolysis, or more specifically “half-cell reactions,” include:

  • Simple electron transfer, e.g., anodic oxidation of Ce (TO) ions

  • Metal deposition, e.g., nickel plating

  • Surface film transformation, e.g., in lead acid batteries

  • Anodic dissolution, e.g., of iron

  • Gas reduction, e.g., oxygen in porous gas diffusion electrodes used in fuel cells

There are a wide variety of electrochemical cells in practice: batteries where electrical energy is produced from the electrode reactions, electrolytic cells where chemical change is derived from an applied potential and fuel cells where electrical energy is produced continuously by the supply of a fuel. Electrolysis can involve different types of processes including reactions in the liquid phase, reactions in the solid phase, located on the surface of the electrodes, and reactions involving the gas phase. An important example of gas phase processes is the electrolytic generation of gas, e.g.,

(2)
(3)
(4)

The anodic production of chlorine and the cathodic formation of H2, and thus hydroxide ion, is the basis for the chlor-alkali industry utilizing a sodium chloride electrolyte. For this process, the anode and cathode reactions in the cells are separated by a diaphragm or Membrane to prevent chlorine gas reacting with the hydroxide (of sodium). This technique of separating the anode and cathode reactions during electrolysis is common. The appropriate separator, by necessity, must allow transport of suitable ions through its structure. In the electrolysis of sodium chloride for chlorine production, the separator should ideally enable transport of sodium ion alone, as shown in Figure 2, to form sodium hydroxide with hydroxide ion generated at the cathode (reaction 3).

In all cells, electrolysis involves reactions at both electrodes and the overall movement of ions in the solutions maintains a neutrality of charge in the electrolyte. Generally therefore, electrolysis processes are written as overall reactions of two-electrode processes, e.g., in chlor-alkali electrolysis the reaction

(5)

represents the electrolysis of sodium chloride in water to give sodium hydroxide, chlorine gas and hydrogen.

For electrolysis to occur, a source of energy is required to move ions in electrolytes and to overcome overpotentials at the electrodes. This energy is either supplied by an external power source, as in electrolytic processes (e.g., recharging of lead/acid batteries) or is obtained from available “chemical” or free energy of the reaction (e.g., discharging of batteries). It is essentially transformed into heat in the electrolyte and is directly related to the internal resistance of the cell R and the applied current, i.e., I2R (Joules s−1).

In electrolysis, this energy requirement can be reduced by factors which lower internal resistance (small distance between the electrodes, high-electrolyte conductivity) and which reduce overpotential, e.g., by the use of electrocatalysts or alternative electrode materials. Electrocatalysts provide alternative reaction pathways for the inherently kinetic, slow step of the electrode process and enable reaction at higher current densities closer to equilibrium potential. Many electrolyses have benefited from the use of electrocatalysts. For example, in chlorine cells anode design has been revolutionized in the 1960s by the use of coated titanium electrodes (so-called dimensional stable anodes or DSAs) as replacements to carbon. The coating, based on ruthenium oxide, valve metals, precious metals and transition metals, gave significant reductions in overpotential (< 50 mV for chlorine generation). Such materials are widely used for many commercial electrolyses.

Number of views: 10521 Article added: 2 February 2011 Article last modified: 10 February 2011 © Copyright 2010-2017 Back to top