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Overview An ionic compound is dissolved with an appropriate solvent, or otherwise melted by heat, so that its ions are available in the liquid. An electrical current is applied between a pair of inert electrodes immersed in the liquid. The negatively charged electrode is called the cathode, and the positively charged one the anode. Each electrode attracts ions which are of the opposite charge. Therefore, positively charged ions (called cations) move towards the cathode, while negatively charged ions (termed anions) move toward the anode. The energy required to separate the ions, and cause them to gather at the respective electrodes, is provided by an electrical power supply. At the probes, electrons are absorbed or released by the ions, forming a collection of the desired element or compound. The amount of electrical energy that must be added equals the change in Gibbs free energy of the reaction plus the losses in the system. The losses can (theoretically) be arbitrarily close to zero, so the maximum thermodynamic efficiency equals the enthalpy change divided by the free energy change of the reaction. In most cases the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance in the electrolysis of steam into hydrogen and oxygen at high temperature, the opposite is true. Heat is absorbed from the surroundings, and the heating value of the produced hydrogen is higher than the electric input. In this case the efficiency can be said to be greater than 100%. (It is worth noting that the maximum theoretic efficiency of a fuel cell is the inverse of that of electrolysis. It is thus impossible to create a perpetual motion machine by combining the two processes. See water fuel cell for an example of such an attempt.) In electrolysis, the anode is the positive electrode, meaning it has a deficit of electrons; species in contact with the anode can be stripped of electrons (i.e., they are oxidized). The cathode is the negative electrode, meaning it has a surplus of electrons. Species in contact with the cathode tend to gain electrons (i.e., they are reduced). A higher current flow (amperage) through the cell means it will be passing more electrons through it at any given time. This means a faster rate of reduction at the cathode and a faster rate of oxidation at the anode. This corresponds to a greater number of moles of product. The amount of current that passes depends on the conductance of the electrodes and electrolyte, though it also depends on how much current the power source itself can generate. Current also makes a difference in that it can shift chemical equilibria by sheer mass action. The processes in an electrolytic cell with just two or three reactants can become very, very complex. Most of the time it's best to search the literature to see what current density works best for a desired process. For instance, metals plated at a certain current density might form a durable and shiny coating on the substrate, while some other current density might form an excessively grainy, dull coating. A higher potential difference (voltage) applied to the cell means the cathode will have more energy to bring about reduction, and the anode will have more energy to bring about oxidation. Higher potential difference enables the electrolytic cell to oxidize and reduce energetically more "difficult" compounds. This can drastically change what products will form in a given experiment. On a practical level, both current and voltage determine what will form in a cell. The following technologies are related to electrolysis: Electrolysis of water One important use of electrolysis is to produce hydrogen. The reaction that occurs is 2H2O(aq) → 2H2(g) + O2(g) This has been suggested as a way of shifting society toward using hydrogen as an energy carrier for powering electric motors and internal combustion engines. (See hydrogen economy.) Electrolysis of water can be achieved in a simple hands-on project, where electricity from a battery or low-voltage DC power supply (e.g. computer power supply 5 volt rail) is passed through a cup of water (in practice a saltwater solution or other electrolyte will need to be used otherwise no result will be observed). Using platinum electrodes, hydrogen gas will be seen to bubble up at the cathode, and oxygen will bubble at the anode. If, however, any other metal is utilised for the anode the oxygen will react with the anode instead of being released as a gas. For example using iron electrodes in a sodium chloride solution electrolyte, iron oxide will be produced at the anode, which will react to form iron hydroxide. When producing large quantites of hydrogen, this can significantly contaminate the electrolytic cell - which is why iron is not used for commercial electrolysis. The energy efficiency of water electrolysis varies widely. The efficiency is a measure of what fraction of electrical energy used is actually contained within the hydrogen. Some of the electrical energy is converted to heat, a useless by-product. Some reports quote efficiencies between 50–70%* This efficiency is based on the Lower Heating Value of Hydrogen. The Lower Heating Value of Hydrogen is thermal energy released when Hydrogen is combusted. This does not represent the total amount of energy within the Hydrogen, hence the efficiency is lower than a more strict definition. Other reports quote the theoretical maximum efficiency of electrolysis. The theoretical maximum efficiency is between 80–94%.*. The theoretical maximum considers the total amount of energy absorbed by both the hydrogen and oxygen. These values only refer to the efficiency of converting electrical energy into hydrogen's chemical energy. The energy lost in generating the electricity is not included. For instance, when considering a power plant that converts the heat of nuclear reactions into hydrogen via electrolysis, the total efficiency is more like 25–40%.* Experimenters Scientific pioneers of electrolysis included: More recently, electrolysis of heavy water was performed by Fleischmann and Pons in their famous experiment, resulting in anomalous heat generation and the controversial claim of cold fusion. First law of electrolysis In 1832, Michael Faraday reported that the quantity of elements separated by passing an electrical current through a molten or dissolved salt was proportional to the quantity of electric charge passed through the circuit. This became the basis of the first law of electrolysis. Second law of electrolysis Faraday also discovered that the mass of the resulting separated elements was directly proportional to the atomic masses of the elements when an appropriate integral divisor was applied. This provided strong evidence that discrete particles of electricity existed as parts of the atoms of elements. Industrial uses Military uses As well as producing hydrogen, electrolysis also produces oxygen. Nuclear submarines are able to generate breathable oxygen from the water around them, so can remain underwater for as long as their fuel lasts. Space stations can also use electrolysis to produce amounts of extra oxygen from waste water or surplus water produced from the Space Shuttle fuel cells. Both these applications depend on having an abundant electrical supply, from either the reactor or solar panels. Examples Electrolysis of an aqueous solution of table salt (NaCl, or sodium chloride) produces aqueous sodium hydroxide and chlorine, although usually only in minute amounts. NaCl(aq) can be reliably electrolysed to produce hydrogen. In order to produce chlorine commercially, molten sodium chloride is electrolysed to produce sodium metal and chlorine gas. These will react violently, so a mercury cell is used to ensure they do not come into contact with each other. Electrolysis can be conducted relatively safely in the house, with use of any low voltage battery and a solution of salt water. A two volt battery is capable of producing a visible stream of gas from both electrodes. However caution should be exercised to avoid splashing any of the solutions in the eyes, or mixing the gases, which could result in the formation an explosive mixture See also | ||||||||||
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