Water electrolysis

Electrolysis is an electrochemical process in which electrical energy is the driving force of chemical reactions. Substances are decomposed, by passing a current through them. The first observation of this phenomenon was recorded in 1789. Nicholson and Carlisle were the first who developed this technique back in 1800 and by the beginning of the 20 century there were already 400 industrial water electrolysis units in use. As mentioned before, water is decomposed to hydrogen and oxygen, by passing a current through it in the presence of suitable substances, called electrolytes. Electric current causes positively charged hydrogen ions to migrate to the negatively charged cathode, where a reduction takes place in order to form hydrogen atoms. The atoms formed then combine to form gaseous hydrogen molecules (H2). On the other hand, oxygen is formed at the other electrode (the positively charged anode). The stoichiometry of the reaction is two volumes of hydrogen to one volume of oxygen. The most important part of the construction of electrolysis units is to use adequate electrodes to avoid unwanted reactions, which produce impurities in the hydrogen gas. Another necessary component of such a unit is a separating membrane that allows the passage of ions, or electrons and not oxygen, or hydrogen atoms. This membrane allows the gases to be kept separate in order to avoid the risk of an explosive mixture being formed in the electrolysis unit. In the initial discovery of electrolysis, an acidic water solution was used, but nowadays there is a trend towards alkaline electrolytes such as potassium hydroxide (KOH). This technology offers the advantages of materials which are cheaper and less susceptible to corrosion compared to those required to handle acids. Electrolysis plants with normal or slightly elevated pressure usually operate at electrolyte temperature of 70-90C, cell voltage of 1.85-2.05 V and consume 4-5 KWh / m of hydrogen, which is obtained at a purity of 99.8% and more. Pressure electrolysis units run at 6-200 bar and there is no significant influence on the power consumption. Because of its high energy consumption and also of the quite substantial investment, water electrolysis is currently used for only 4% of world hydrogen production. Nowadays research and development into high efficiency electrolysers is flourishing in many areas. A way of improving electrolysis units efficiency is by increasing the process temperature which lowers the voltage required to electrolyse the water, but also requires more expensive materials. Despite the fact that the total energy needed for the electrochemical decomposition of water decreases only slightly with increasing temperature, the reversible part of the energy requirement (∆F), which is supplied as electrical energy, decreases considerably. Therefore an increasing amount of the total energy could be supplied as heat. At elevated temperatures (800-900C) the electric power consumption is approximately only 3 kWh / m of hydrogen. It must be noted that this technology is still in the development stage. Electrolysis is considered as the “cleanest” way to produce hydrogen, when the required electricity is derived from renewable energy sources. In countries with a lot of waterfalls, hydroelectricity can be used as the energy source for water electrolysis. Other renewable sources that could be used for supplying electrolysis units are solar, aeolic and geothermal energy. Photoelectrolysis, in which the photovoltaic cells are also electrodes that decompose water to hydrogen and oxygen gas could be used for the production of hydrogen. These technologies could be used in order to store energy as hydrogen, which can be transformed to electricity in fuel cells, when the natural source of energy is not available. The production of hydrogen through electrolysis using renewable energy sources has the smallest impact on the environment. 1. Historical background The history of water electrolysis started as early as the first industrial revolution, in the year 1800, when Nicholson and Carlisle were the first to discover the ability of electrolytical water decomposing. By 1902 more than 400 industrial water electrolysis units were in operation and in 1939 the first large water electrolysis plant with a capacity of 10,000 Nm H2/h went into operation. In 1948, the first pressurized industrial electrolyser was manufactured by Zdansky/Lonza. In 1966, the first solid polymer electrolyte system (SPE) was built by General Electric, and in 1972 the first solid oxide water electrolysis unit was developed. The first advanced alkaline systems started in 1978. The history ends up in our days with the development of proton exchange membranes, usable for water electrolysis units and fuel cells, by DuPont and other manufacturers, due to the developments in the field of high temperature solid oxide technology and by the optimization and reconstruction of alkaline water electrolysers [1]. 2. Theory of water electrolysis The electrolysis of water is considered a well-known principle to produce oxygen and hydrogen gas. In Fig.1 a schematic of an electrochemical cell is presented. The core of an electrolysis unit is an electrochemical cell, which is filled with pure water and has two electrodes connected with an external power supply. At a certain voltage, which is called critical voltage, between both electrodes, the electrodes start to produce hydrogen gas at the negatively biased electrode and oxygen gas at the positively biased electrode. The amount of gases produced per unit time is directly related to the current that passes through the electrochemical cell. In water, there is always a certain percentage found as ionic species; H and OH represented by the equilibrium equation: Fig.1: Sketch of an electrochemical cell [2]. H2O (l)↔ H (aq) + OH (aq) (1) Oxygen and hydrogen gas can be generated at noble metal electrodes by the electrolysis of water: + electrode (anode): 4OH ↔ 2H2O + O2 + 4e (2a) electrode (cathode): 2H (aq) + 2e ↔ H2 (g) (2b) In case of acidic or basic water, the reactions which occur at the electrode interface are slightly different. In water electrolysis there are no side reactions that could yield undesired byproducts, therefore the net balance is: 2H2O → (4e) → O2 + 2H2 (3) The minimum necessary cell voltage for the start-up of electrolysis, E ocell , is given under standard conditions (P, T constant) by the following equation: E ocell = nF G o ∆ −