Chapter 1: Introduction
On an industrial scale, chlorates are prepared by electrolysis. Electrolyzing a solution of a chloride at elevated temperatures yields a chlorate. This method can be downscaled 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 setups. 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.
Chapter 2: Electrolytic preparation
The electrolysis is carried out in a diaphragmless 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.
2.1 theory
Mechanism of chlorate formation
The reactions taking place in chlorate cells are not fully understood even today. A summarized description of the process will be given here, and the interested reader is referred to the literature

listed below for a more extensive description. 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 neutralize 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 effect one as we will shortly come to see, the conditions in the cell are usually optimized 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:

electrons Oxygen is evolved in this reaction, which means a loss of current efficiency (the energy used for oxidizing 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, optimizing 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 dichromate 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 setups), this is much less laborious and therefore probably the preferred method for home setups.
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.
2.2 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. If the chlorates are going to be prepared on a more or less regular basis, it probably pays to spend some more time designing a cell. It will also improve efficiency somewhat, but unlike in industrial setups where high efficiency is mandatory to be able to compete, the home experimenter can do with less efficient cells. The two main disadvantages of a low efficiency are that it takes more time for the conversion to complete, and that more electricity is required. To give some indication of the power consumption of the process: typical figures for industrial cells lie in the range 4.5 to 5.5 kWh per kg of sodium chlorate. In this section some of the things to consider when building and designing chlorate cells will be discussed. The reader can design his own cell based on the information given. An example of a cell, the small test cell I currently use to experiment with, has been given but it is by no means perfect, and it is probably better to design your own. The example has merely been given to illustrate some principles.
2.3 An example
The example given here consists of a small cell, of 200 ml electrolyte volume. The cell is normally operated with graphite or graphite substrate lead dioxide anodes. Platinum sheet has also been tried with, unsurprisingly, good success. The electrolyte consists of sodium chloride with either some potassium dichromate or potassium fluoride added, depending on wheter graphite or lead dioxide anodes are used. The cathode consists of a stainless steel wire spiraling down. The wire is corroded where it is not submerged, so it has to be replaced occasionally. The connections to the anode and cathode are made outside the cell but do corrode from the gasses and electrolyte mist. This is partially prevented by leading the gasses away from the connections with a vent tube, as shown in the picture. Covering the connections with hot melt glue also helps, but the heat generated in a faulty connection may cause the hot melt to melt.. The temperature is controlled by placing the cell in a water bath, which acts as a heat sink. If the temperature is too