2.1 theory
2.2 cell construction
2.3 an example
2.4 cell volume
2.5 cell body materials
2.6 electrode materials
2.7 pH and temperature control
2.8 preparing the electrolyte
2.9 operating the cell
2.10 processing the electrolyte
2.11 purifying the product

3.1 starting materials
3.2 method
3.3 purifying the product

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 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 laborous, 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 inneficient, for which reason it was abondoned quite soon after the electrochemical method became feasible at industrial scale.

Chapter 2: Electrolytic preparation

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

The reactions taking place in chlorate cells are not fully understood even today. A summarised 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^- Cl_2(aq) + 2 electrons

At the cathode:

2H₂O + 2 electrons H₂ + 2OH^-

Cl_2(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^-

2HClO + ClO^- +2OH^- ClO₃^- + 2Cl^- + H₂O

Alternatively, chlorate may also be formed by oxidation of hypochlorite at the anode as follows:

6HClO + 3 H₂O ³/₂ O₂ + 4Cl^- + 2ClO₃^- + 12H⁺ + 6 electrons

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 untill 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 untill all chloride has been converted into perchlorate. Although slightly less efficient (and therefore not used a lot in industrial setups), this is much less laborous and therefore probably the prefered 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 maintanance. 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 is 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 hotmelt 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 low, styrofoam isolation is provided. The cell is operated outside, causing the temperature to fluctuate between day and night. The pH is checked about twice a day and adjusted if necessary with hydrochloric acid. The power source used is an old computer power supply. The output voltage can be regulated within certain limits and this is done to maintain a current of about 4 amperes. An other model computer power supply was used previously that did not allow control over the output voltage. Current adjustment was done by widening or narrowing the cathode spiral, effectively reducing or increasing the anode-cathode distance.

Theoratically, if 100% efficiency could be reached, the cell would have the capacity to convert approximately 35 grams of sodium chloride to 64 grams of sodium chlorate per day. Using a metathesis reaction with potassium chloride this would yield 74g of potassium chlorate. In practice the average yield is about 40 grams of potassium chlorate a day from which an efficiency of 55% can be calculated.

2.4 Cell volume

The table below gives some idea of how well a number of materials stand up to corrosion. The column 'protected' lists how well metals resist corrosion when cathodically protected. The column 'unprotected' lists materials used as is.

material	corrosion resistance when unprotected	corrosion resistance when cathodically protected
Iron	--	+
Stainless steel	-	+
Titanium	++	++
Copper	--	++
Brass	--	+
Tantalum	++	++
Platinum	++	++
Aluminum	--	+-
PET	++	X
Poly ethylene	+	X
Poly propylene	+	X
PVC	++	X
Rubber	+-	X
Hot melt glue	+-	X
Styrene	-	X
Graphite	+	+
Silicone rubber	--	X
Concrete	+	X
Glass	++	X
Ceramics	++	X
Wood	--	X
Polyester	-	X

2.6 Electrode materials

Cathode materials

Both stainless and mild steel find widespread use as cathode materials. Brass and copper may also be used. Each of these metals is protected to a certain extent by the negative charge present on the cathode as long as they are submerged and the current per surface area is high enough. Unsubmerged parts of the cathode corrode at a high rate however due to the action of evolving gasses and droplets of cell electrolyte.

It seems that under some conditions the chromium in stainless steel can dissolve, even though the cathode does not seem to corrode. A yellow electrolyte is the result from which barium chromate can be precipitated even if no chromate was added, which will be described later (see processing the electrolyte). The presence of chromates could lower the efficiency of cells employing lead dioxide anodes.

Finally, contamination of the final product with copper (from brass or copper erosion products) can be dangerous when the product is to be used in pyrotechnic purposes. Although this is unlikely to be a great problem since the impurities are usually removed easily and completely by filtration (as will be described later) it is good to be aware of the possibility.

2.7 pH and temperature control

Like temperature control, pH control is not essential for chlorate manufacture. Efficiency is improved greatly however if the pH is kept within the range 5.5 - 6.5 (slightly acidic) as explained in the theory section. Graphite anodes also tend to erode faster at high pH, so maintaining the pH will extend graphite anode life. In commercial setups pH control is done manually by periodic additions of hydrochloric acid. Automated pH control seems to be difficult and expensive to realise. If anyone devices a practical method of doing this, I'd be interested to hear about it.

2.8 Preparing the electrolyte

When electrolyte from a previous batch of chlorate is available the following steps can be used to recycle the electrolyte.

1. If the electrolyte is not clear but has solid particles in it, filter to remove these. See the section on filtering below.

2. dissolve any impure chlorate from the purification and extraction steps.

3. Now, re-saturate the solution with sodium chloride. The procedure mentioned above in step 1 of 'preparing a fresh electrolyte' may be used.

4. The chromate, dichromate or fluoride if added is still present so does not need to be replenished. The pH should be readjusted, like in step 3 for preparing a fresh solution above.

2.9 Operating the cell

As explained in the theory section the voltage over the cell may vary. The current should be kept more or less constant at a value determined by the cell design. As a rule of thumb, supply 2 amperes of current per 100 ml of electrolyte. If graphite anodes are used it is better to supply less current since that will increase anode life (30 mA per square centimere of anode surface area is typical). A constant current supply is ofcourse the most convenient to use for regulating current, but manually adjusting the voltage from time to time also works well. The current usually only changes very gradually, and the precise value is not very critical. In any case, measure the current at regular intervals and record them. That information is required to determine when a batch is complete, as described below in the paragraph 'running times'.

Maintaining optimal conditions

As explained, the pH of the electrolyte will tend to rise. Also, some of the water will evaporate and some will be consumed in the reaction. The temperature may also vary with ambient temperatures. For good efficiency these variables must be kept within certain limits.

Every once in a while, make up for evaporated and consumed water. This can be done with water, but it is better to use a saturated sodium or potassium chloride solution. That way, the chloride concentration will be kept at a higher level which improves efficiency. It will also prevent excessive formation of perchlorate, and in graphite anode based cells it reduces anode wear.

The pH will rise during operation of the cell, and it is best to lower the pH regularly to a value of about 6. A high pH value is best corrected by adding hydrochloric acid occasionally. If too much HCl is added, it may be corrected with sodium hydroxide solution. The pH is self correcting to a certain extend as well, since at very low pH the cell will produce more chlorine gas. This will then escape, raising the pH again. Measuring the pH of the solution can be done with common pH paper. However, if the paper is simply dipped directly into the electrolyte the hypochlorite present will usually bleach the paper making a measurement impossible. This problem can be overcome by boiling a sample of the solution for 5 minutes and measuring the pH of that. The boiling destroys the hypochlorite.

The temperature will usually reach a more or less constant value quickly. If the cell is placed outside, the temperature may vary between night and day and between seasons ofcourse, so then some sort of control may be necessary to maintain the optimum. Usually, cooling is necessary but it will depend on the specific cell design as explained in the theory section.

Running times

The time required to convert a certain amount of chloride to chlorate depends on the current and the efficiency. The latter can vary dramatically from cell to cell and it is therefore not possible to state precise running times. It is however possible to calculate the required running times for a hypothetical cell operating with 100% current efficiency. The calculation will show that the conversion of one mole of chloride to chlorate requires 160.8 amperage hours. So, for example, a cell containing 100 grams of sodium chloride will require 100/58.6 * 160.8 = 274.4 Ah if it operates at 100% efficiency. For a current effiency other than 100%, increase the running times in proportion (to convert 100g of NaCl with 80% current efficiency one needs 274.4/80*100 = 343 Ah). So, if a current of 3 amperes flows through the cell, it requires 274.4 / 3 = 91.47 hours (91 hours, 28 minutes) to finish.

2.10 Processing the electrolyte

When done, the product must be extracted from the electrolyte and the electrolyte can be recycled for the next batch (see preparing the electrolyte).

Filtering

Usually, the impurities consist of erosion products of the anodes, the cell walls, and the unsubmerged parts of the cathodes. These particles may be very small and are not always easily removed with common filtering paper. Filtering through a layer of diatomeous earth (sold in shops for aquarium supplies) in a filter or on a piece of cloth sometimes solves the problem. Another great idea for a filter comes from E.S. However, just filtering will not always remove all solid impurities. A common impurity that is hard to remove is suspended iron hydroxide, originating from corrosion of (stainless) steel cathodes. The fluffy, voluminous form of the material often gives it a white or yellowish foggy appearance. This is next to impossible to remove unless some sodium hydroxide or pool coagulant is added first. This causes the iron hydroxide particles to coagulate, making them easy to remove by filtration. Another possibility is to add hydrochloric acid to lower the pH to between 2 and 3. This will dissolve the iron hydroxide. If sodium hydroxide is then added to raise the pH to above 7 again, the iron hydroxide is precipitated in a more dense form which is easily removed by filtration, even with common filter paper.

In this step, the advantages of using sodium chloride will become evident. When potassium chloride is used instead potassium chlorate crystallises during operation of the cell due to its relatively low solubility. To separate the potassium chlorate from insoluble impurities the electrolyte has to be filtered hot. The solution usually takes quite a long time to pass through the filter and if it cools during this time, potassium chlorate will crystallise and block the filter. Alternative methods have been developed to separate potassium chlorate from insoluble impurities. For example, the solution may be boiled and sufficient water added to dissolve all potassium chlorate. If the solution is then allowed to cool slowly, crystals of potassium chlorate will form on the suspended insoluble impurites. These will sink to the bottom, usually leaving a clear solution. The clear solution is then carefully decanted and allowed to cool further . This method will not remove the insoluble purities as well as filtering will but it is much less laborous.

Destruction of hypochlorite

Next, the electrolyte is boiled to decompose remaining hypochlorite. 15 minutes of vigorous boiling is sufficient. After that, the pH of the solution is checked and it is made slightly alcaline by adding sodium hydroxide solution. Bring the pH to between 8 and 9.

Metathesis reaction

At this point, a clear solution of sodium chlorate (with residual chloride) has been obtained. This can be used either to prepare potassium or barium chlorate (or other chlorates which are not further elaborated upon here), or it can be used to prepare perchlorates, described elsewhere.

Potassium chlorate is by far the most commonly used chlorate in pyrotechnics. For practical purposes, the preparation of this compound is discussed here. For the preparation of barium chlorate the amounts will have to be adjusted.

1. Weigh out either 127g of potassium chloride or 355g barium chloride for every 100 g of sodium chloride that was started with, depending on wheter you want to prepare potassium or barium chlorate. Dissolve this in as little water as possible (dissolve in minimum amount of boiling hot water, add a bit more water and allow to cool. Nothing should precipitate. If it does, add some more water and heat again)

2. Add this solution to the electrolyte. A white precipitate of potassium or barium chlorate should form.

3. Bring the solution to a boil. Add 20 ml amounts of water to the solution in 5 minute intervals untill all chlorate has dissolved. If all chlorate dissolves upon heating without the addition of extra water, allow the water to evaporate untill a thin crust of chlorate forms on the surface (indicating that the saturation point has been reached). Then add 20 ml of water and boil for a minute to redissolve the crust.

4. Allow the solution to cool to room temperature. Potassium or barium chlorate will crystallise. If it has cooled to room temperature, cool further to 0 deg C.

5. Filter to obtain the crude chlorate crystals. Rinse them thoroughly with ice-cold water. The filtrate can be saved to prepare the next electrolyte, as is described in the section on recycling electrolyte.

2.11 Purifying the product

The crude product can be purified by recrystallisation. The low solubility of potassium chlorate makes this method very convenient to use and will greatly improve the purity with a relatively small loss of product. Barium chlorate is somewhat more soluble and to prevent losses it is a good idea to use the impure barium chlorate 'waste' from this procedure in the electrolyte of a new cell. Some treatment is necessary, which was described earlier. If a single recrystallisation step does not yield a sufficiently pure product, the method can be repeated to further increase the purity. Usually one or two recrystallisations will yield a product that does not impart the characteristic yellow color of sodium impurities to a flame.

1. Place the crude product in a pan and add 100 ml of water for every 35g of crude potassium chlorate or 50g of barium chlorate. Bring this to a boil.

2. Add 20 ml amounts of water to the boiling solution untill all the chlorate has dissolved.

3. Check the pH of the boiling solution. It should be neutral or slightly alcaline. If it is acidic, add potassium hydroxide solution untill it is slightly alcaline (pH 7..8) again. If this is not done, traces of acid may be included in the product making it very dangerous to use in pyrotechnic compositions.

4. Allow the solution to cool to room temperature. The chlorate will crystalise.

5. Filter and rinse the crystals in the filter well with ice cold water. The filtrate may be used to prepare the electrolyte for a new cell, as was described in the section on recycling old electrolyte.

6. The crystals may be dried in an oven at 100 deg C.

Thermal decomposition of hypochlorites

This is an alternative method of chlorate manufacture. It is more laborous than the electrolytic method, and can only be used for small batches at a time. The starting materials are quite easily available however as bleach and pool chlorinating agents and it only requires the use of simple tools.

3.1 Starting materials

Possible starting materials are sodium hypochlorite and calcium hypochlorite. The former is available in solution as bleach and antifungal spray for bathrooms. Calcium hypochlorite finds use as a chlorinating agent for pools. However, different varieties exist. Carefully read the package to make sure you have the right material. It usually states a '85% available chlorine' content for calcium hypochlorite. A higher available chlorine content may mean it is something else, most likely trichlorohydrocyanuric acid.

3.2 Method

warning: On one occasion an small explosion occured when I was doing this preparation. I am not sure exactly what caused the explosion. It seems to have been a steam explosion. I was also not sure wheter I was using calcium hypochlorite or trichlorohydrocyanuric acid, another common pool chlorinating agent. It seems to be very uncommon that explosions happen and they can probably be prevented by vigorous stirring, but I thought everyone attempting this method should know so proper precautions can be taken. The procedure below has been optimised to reduce the chances of an explosion happening.

1. Place 250 ml of water in a heat resistant glass or stainless steel container, large enough to hold twice that volume.

2. Bring the water to a boil.

3. To the boiling water, add 125 gram of calcium hypochlorite in 10 gram portions. The calcium hypochlorite usually comes in tablets, which need to be crushed first. Stir vigorously during this step, occasionally scraping over the bottom to prevent the formation of a cake of calcium chloride. The solution will foam a lot. If too much foam is developed, do not add any more calcium hypochlorite and boil untill the foam subsides. Then continue adding calcium hypochlorite.

4. When all calcium hypochlorite has been added, continue boiling untill no more foaming is observed. Stir continuously.

5. Allow the solution to cool down, and filter to remove the precipitated calcium chloride.

6. In a separate container, dissolve 68 grams of potassium chloride in the smallest volume of water possible (approximately 195 ml). This can be done by dissolving the potassium chloride in about 200 ml of water, and allowing it to cool. If crystals form, add some more water, boil again to dissolve the potassium chloride, and allow to cool again. If crystals form, repeat. If not, the solution is ready to use.

7. Mix this solution with the boiled calcium hypochlorite solution. A white precipitate of potassium chlorate should form.

8. Bring the solution to a boil and add water untill all potassium chlorate has dissolved.

9. Allow the solution to cool slowly. Crystals of potassium chlorate will form. Cool to 0 deg C.

10. Filter to obtain the raw potassium chlorate. Rinse the crystals in the filter with ice-cold water. The product can be further purified as described below.

3.3 Purifying the product

The product can be purified by recrystallisation, just like the product of the electrolytic procedure. For convenience, the same procedure is given again here:

1. Place the crude product in a pan and add 100 ml of water for every 20 g of crude product. Bring this to a boil.

2. Add 20 ml amounts of water to the boiling solution untill all the potassium chlorate has dissolved.

3. Check the pH of the boiling solution. It should be neutral or slightly alcaline. If it is acidic, add potassium hydroxide solution untill it is slightly alcaline (pH 7..8) again. If this is not done, traces of acid may be included in the product making it very dangerous to use in pyrotechnic compositions.

4. Allow the solution to cool to room temperature. Potassium chlorate will crystalise.

5. Filter and rinse the crystals well with ice cold water. The filtrate may be discarded or concentrated by evaporation and the residue added to the electrolyte for a next batch.

6. The crystals may be dried in an oven at 100 deg C.

Chapter 4: Literature

The amount of literature available is overwhelming. A short list of interesting reading material follows.

1. F Hine, "Electrode processes and electrochemical engineering", Plenum Press, New York (1985)

2. F. Foerster and E. Muller, Z. Elektrochem, 8, 8, 515, 633, 923 (1902); 9 171 (1903); 10, 781 (1904).

3. Webpage of Mike Brown

This list is under construction

Copyright Wouter Visser, May 1999.
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