There are several variants of this process: the Le Sueur cell (1893), the Hargreaves-Bird cell (1901), the Gibbs cell (1908), and the Townsend cell (1904). This technology was also developed at the end of the nineteenth century. In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode. In Europe, mercury cells accounted for 43% of capacity in 2006 and Western European producers have committed to closing or converting all remaining chloralkali mercury plants by 2020. In the United States, there will be only five mercury plants remaining in operation by the end of 2008.
In Japan, mercury-based chloralkali production was virtually phased out by 1987 (except for the last two potassium chloride units shut down in 2003). It is estimated that there are still around 100 mercury-cell plants operating worldwide. The mercury process is the least energy-efficient of the three main technologies (mercury, diaphragm and membrane) and there are also concerns about mercury emissions. The mercury is then recycled to the primary cell by a pump situated at the bottom. This flows continuously into a separate reactor (" denuder" or "secondary cell"), where it is usually converted back to mercury by reaction with water, producing hydrogen and sodium (or potassium) hydroxide at a commercially useful concentration (50% by weight). When a potential difference is applied and current flows, chlorine is released at the titanium anode and sodium (or potassium) dissolves in the mercury cathode forming an amalgam. Today, in the "primary cell", titanium anodes clad with platinum or conductive metal oxides (formerly graphite anodes) are placed in a sodium (or potassium) chloride solution flowing over a liquid mercury cathode. The "rocking" cells used have been improved over the years. Mercury cell electrolysis, also known as the Castner–Kellner process, was the first method used at the end of the nineteenth century to produce chlorine on an industrial scale. The sodium–mercury amalgam flows to the center cell, where it reacts with water to produce sodium hydroxide and regenerate the mercury. Overall process: 2 NaCl (or KCl) + 2 H 2O → Cl 2 + H 2 + 2 NaOH (or KOH)Ĭastner–Kellner cell: Sodium chloride is electrolyzed between the "A" anode and "M" mercury cathode in the side cells, with chlorine bubbling up into the space above the NaCl and the sodium dissolving in the mercury. There are three industrial methods for the extraction of chlorine by electrolysis of chloride solutions, all proceeding according to the following equations:Ĭathode: 2 H + (aq) + 2 e − → H 2 (g) Anode: 2 Cl − (aq) → Cl 2 (g) + 2 e − Chlorine can also be produced by the electrolysis of a solution of potassium chloride, in which case the co-products are hydrogen and caustic potash ( potassium hydroxide). These two products, as well as chlorine itself, are highly reactive. The production of chlorine results in the co-products caustic soda ( sodium hydroxide, NaOH) and hydrogen gas (H 2). 3.6 Caustic handling, evaporation, storage and loadingĬhlorine can be manufactured by the electrolysis of a sodium chloride solution ( brine), which is known as the Chloralkali process.1.2 Diaphragm cell electrolysis (bipolar).