Introduction Electrolytic Preparation

Chlorates and perchlorates are among the most often used oxidisers in pyrotechnics. Especially potassium perchlorate (KClO4), potassium chlorate (KClO3), ammonium perchlorate (NH4ClO4), and barium chlorate (Ba(ClO3)2) are used extensively. Chlorates and perchlorates are chemically related but from a pyrotechnical standpoint perchlorates are generally safer to use. Chlorates require special handling and mixtures made with chlorates tend to be more sensitive than similar mixtures based on perchlorates and some may ignite spontaneously. Potassium chlorate is extensively used commercially since it is cheaper than potassium perchlorate, and many accidents have happened that might have been prevented if the corresponding perchlorate was used instead. The use of chlorates is therefore not recommended. If you do decide to use them, make sure you really understand what you are getting involved with. The amateur pyrotechnist can prepare chlorates and perchlorates at home with some effort. Home made chemicals are usually more expensive than commercially produced chemicals and it is of course laborious, but taking shipping costs into account, the fact that perchlorates and chlorates are sometimes quite hard to obtain in most countries and their usefulness in pyrotechnics it may pay to try to make your own. Chlorates are the easiest to prepare, but again, the use of chlorates is not recommended for amateur use. Introduction On an industrial scale, chlorates are prepared by electrolysis. Electrolysing a solution of a chloride at elevated temperatures yields a chlorate. This method can be down scaled quite easily for amateur pyro purposes. Other methods of chlorate manufacture exist that may be of interest for small scale use. They are usually less efficient but the economy of the process is not as important for amateur pyro purposes as it is for industrial set-ups. A second method for example consists of heating a solution of hypochlorite. Sodium and calcium hypochlorite are both quite easily available as bleach and pool chlorinating agent respectively. Upon heating, the hypochlorite will decompose into both chloride and chlorate. The chlorate is separated and purified. Although slow and laborious, the method is simple and requires very little equipment. In the past chlorates were produced even on an industrial scale by bubbling chlorine gas through a hot hydroxide solution. This process is not very well suited for amateurs since chlorine gas is very dangerous to handle. The process is also extremely inefficient, for which reason it was abandoned quite soon after the electrochemical method became feasible at industrial scale. Electrolytic preparation The electrolysis is carried out in a diaphragm less cell, containing a solution of a chloride. Several chlorides may be used, but the use of sodium chloride has many advantages. Sodium chlorate is easily converted to a number of other chlorates by metathesis reactions. The most commonly used chlorates in pyrotechnics, potassium and barium chlorate, can both be made in this manner. Potassium chloride and barium chloride may also be used to obtain the respective chlorates directly, but this has many disadvantages as will be discussed below. Only sodium chlorate can be used in the manufacture of perchlorates, due to its high solubility. Ammonium chloride should never be used, and should in fact not even be present in the cells in trace amounts. It could result in the formation of two dangerously sensitive and explosive compounds, nitrogen trichloride (NCl3) and ammonium chlorate (NH4ClO3). The formation of both of these compounds should be avoided at all times. Not only can they explode by themselves when present in significant quantities, they can also lead to spontaneous ignition of pyrotechnic mixtures contaminated with even small amounts. Theory Mechanism of chlorate formation The reactions taking place in chlorate cells are not fully understood even today. The theory of Foerster and Mueller regarding the reactions in chlorate cells, developed about 80 years ago, is the most accepted. The following reactions are said to take place at the electrodes: At the anode: - 2Cl Cl2(aq) + 2 electrons At the cathode: - 2H2O + 2 electrons H2 + 2OH The dissolved chlorine gas can then react with water to give hypochlorous acid: + - Cl2(aq) + H2O HClO + H + Cl From this reaction it can be seen that if the chlorine does not dissolve but escapes to the atmosphere, no H+ will be generated to neutralise the OH- formed at the cathode and the pH of the electrolyte will increase. The hypochlorous acid thus formed will react in acid-base equilibrium reactions with water to give - hypochlorite ions and chlorine gas (dissolved). The exact concentrations of dissolved Cl2, ClO and HClO depend on the pH, temperature and pressure among other things. In the solution, chlorate will be formed (mainly) by the following reactions: - - + - 2HClO + ClO ClO3 + H + 2Cl and - - - - 2HClO + ClO +2OH ClO3 + 2Cl + H2O These reactions take place at a rather slow rate. Since this reaction pathway is the most efficient one as we will shortly come to see, the conditions in the cell are usually optimised to increase their reaction rate. The pH is kept within a range where HClO and ClO- are simultaneously at their maximum concentration (which is at around pH=6). The temperature is kept between 60 and 80 degrees centigrade, which is a good compromise between the temperatures required for a high reaction rate, low anode and cell body corrosion and high chlorine solubility (remember the chlorine evolved at the anode has to dissolve in the solution to start with). Many cells also have a large storage tank for electrolyte in which the electrolyte is kept for a while to give these reactions some time to take place. Alternatively, chlorate may also be formed by oxidation of hypochlorite at the anode as follows: 3 - - + 6HClO + 3 H2O /2 O2 + 4Cl + 2ClO3 + 12H + 6 electrons Oxygen is evolved in this reaction, which means a loss of current efficiency (the energy used for oxidising the oxygen in water to the free element is lost when the oxygen escapes to the atmosphere). When the reaction routes are worked out, it turns out that following this path 9 faradays of charge are required to produce 1 mole of chlorate, whereas only 6 faradays are required to do that following the route mentioned earlier. Therefore, optimising the conditions for that route improves current efficiency. To prevent the products from being reduced at the cathode again, a membrane around the cathode was employed in the past. Today, small amounts of chromates or dichromates are added. A layer of hydrated oxides of chromium will then form around the cathode effectively preventing hypochlorite and chlorate ions from reaching the cathode surface. Finally, it should be mentioned that the reactions forming perchlorates do not take place until the chloride concentration has dropped to below about 10%. Therefore, cells can be constructed and operated in such a way that chlorate is produced almost exclusively. The chlorate can then be purified and fed into a perchlorate cell. Depending on the type of anodes used in the chlorate cell, the purification step may also be skipped and the electrolysis continued until all chloride has been converted into perchlorate. Although slightly less efficient (and therefore not used a lot in industrial set-ups), this is much less laborious and therefore probably the preferred method for home set-ups. Cell voltage The current through a cell is related to the reaction rate. Therefore, to obtain a constant reaction rate that suits the cell design, a constant current is usually employed. The voltage over the cell will then fluctuate depending on conditions and cell design. The power consumed by the cell is the product of current and voltage, according to equation P = I * V. From that it can be seen that reducing the voltage over the cell results in a lower power consumption, an important fact for industrial operations. The factors influencing the cell voltage have been thoroughly investigated. Most important are the anode - cathode spacing, the concentration of the electrolyte, the surface area and materials of the electrodes, the temperature and the pH. Without going into details, the cell voltage usually lies in the range 3.5 - 4.5 volts. Of this, approximately 3 volts are required to get the oxidation of chloride to chlorate to take place (and the hydrogen reduction at the cathode), while the rest is used to overcome the resistance of the cell, according to Ohm's law V=I*R. From this law it can be seen that there are two ways to maintain a constant current through a cell: either the voltage over it may be varied or its resistance may be changed. Adjusting the voltage over a cell to maintain a constant current can be done manually or with an electronic circuit. If the power supply does not allow voltage adjustment (such as old PC power supplies or battery chargers for example) or the required electronics are not available, adjusting the resistance of the cell is another option. This could in principle be done by adjusting each of the factors mentioned earlier, the most practical of which is probably the anode-cathode distance. By increasing the distance between the electrodes the resistance of the cell is increased, which reduces the current through the cell. One thing to keep in mind when doing this is that it with decreasing resistance, the heat generated in the cell is increased. Depending on the anode material used it may then be necessary to cool the cell to prevent excessive erosion, more on that later. Cell construction Cells can range in complexity from a glass jar with a nail and a old battery electrode to well designed, corrosion resistant cells with thermostats, pH control, circulating electrolyte and coulometers. Even the simplest of cells will work, but it will require more maintenance.

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