Chlorate, Perchlorate production cells -archive file

[megalomania]

PYRO500
Moderator
Posts: 1466
From: somewhere in florida
Registered: SEP 2000
posted February 18, 2001 12:00 AM
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I know production of a chlorate, or perchlorate has been covered and is on a faq that I downloaded. What I want to know is details on the construction and your experiences you have had with your cells, basicly I want to know the details of your chlorate/perchlorate cells like the annode,cathode materials effiency success rates production steps etc. the reason I am asking is that I am thinking of setting a cell up and I want to get all the info available before attempting. I beleve I may be able to produce this stuff by the 5 gallon paint jug full and I just want some input first. Thanks for any insight any of you may provide


ALENGOSVIG1
Moderator
Posts: 766
From: Vancouver, Canada
Registered: NOV 2000
posted February 18, 2001 12:53 AM
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I have been wondering about this for awhile. You get small amounts of sodium chlorate which must be converted to potassium chlorate but it would be incrediablly cheap to make and run. A small cell runs off of approximately 4.5 volts.
I will be making one of these this week hopefully and will tell you how it goes. Is the anode in a car battery lead dioxide? I need to find lead dioxide for a good anode. I cant exactly affort a ingot of platinum or titanium.

It is some work getting it all set up but once you've got it running it will be a constant sourve of almost free chlorate or perchloarate depending what you want.


kv21
A new voice
Posts: 6
From: Berlin
Registered: FEB 2001
posted February 18, 2001 02:04 PM
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quote:
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Originally posted by ALENGOSVIG1:
Is the anode in a car battery lead dioxide? I need to find lead dioxide for a good anode. I cant exactly affort a ingot of platinum or titanium.
It is some work getting it all set up but once you've got it running it will be a constant sourve of almost free chlorate or perchloarate depending what you want.

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The anodes in car battery's can't be used. It's plated on a lead plate and it's is very porous. The PbO2 will just release after some time. <a href="http://huizen.dds.nl/~wfvisser" target="_blank">http://huizen.dds.nl/~wfvisser</a> describes the construction of a PbO2 anode.

You'll be terribly disappointed if it comes to the mainenance of the cell. The connections with the anodes are very easily broken, and you will notice all sorts of problems after some time. You also need a very powerfull transformer. At least 50 Amps. About 10 V for perchlorates and 5 V for chlorates.

The costs are about 11 USD / kg. But it is essential that your amperage is high, else it will take weeks to make a pound of KClO3.


ALENGOSVIG1
Moderator
Posts: 766
From: Vancouver, Canada
Registered: NOV 2000
posted February 18, 2001 02:44 PM
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I can easily get all sorts of transformers for very cheap. pennies to be exact. Im definataly going to make one. Its just the anode that is the problem. I dont want to use graphite because of its high corrosion rate. And the page you listed has been down for a couple of days now. I have been told that manganese dioxide works well and it can be found in dry cell batteries. I have opened dry cells before and sure enough, there was manganese dioxide. but it was a powder. And i am assuming it needs to be pressed into a rod. I would think you would need a very well build press because the rod would have to stand up to corrosion. Any ideas on anodes anyone?
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technology is a wonderful servant, but a bitch of a master.

Explosives Archive


vehemt
Frequent Poster
Posts: 580
From: Canada
Registered: SEP 2000
posted February 18, 2001 05:41 PM
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<a href="http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/chlorate.html" target="_blank">http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/chlorate.html</a>


firebreether
Frequent Poster
Posts: 108
From:
Registered: NOV 2000
posted February 19, 2001 09:29 PM
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Where can you get transformers for that cheap?


PYRO500
Moderator
Posts: 1466
From: somewhere in florida
Registered: SEP 2000
posted February 19, 2001 10:06 PM
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oh come on! you can get transformers anywhere. I am thinking of rewinding the secondary of a microwave oven transformer to get 12 volts (plus a current limiting circuit) I will probably get that chlorate to be made in a day! (mw transformers are rated at 2000 watts)


angelo
Frequent Poster
Posts: 279
From:
Registered: SEP 2000
posted February 19, 2001 11:02 PM
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then you can make KNO3 from the KClO3
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angelo's place
have a good link? add it here


ALENGOSVIG1
Moderator
Posts: 766
From: Vancouver, Canada
Registered: NOV 2000
posted February 19, 2001 11:52 PM
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Why would you bother making kno3 when you can have powerful oxidisers such as chlorates and perchlorates?
------------------
technology is a wonderful servant, but a bitch of a master.

Explosives Archive


Mr Cool
Frequent Poster
Posts: 991
From: None of your bloody business!
Registered: DEC 2000
posted February 20, 2001 10:19 AM
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Preparing chlorates
Chapter 1: Introduction
Chapter 2: Electrolytic preparation

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

Chapter 3: Thermal decomposition of hypochlorites

3.1 starting materials
3.2 method
3.3 purifying the product

Chapter 4: Literature

Chapter 1: 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 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

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 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- 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 effient 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:

6HClO + 3 H2O 3/2 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.

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 decresing 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 methatesis 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

This is the main factor affecting a cells capacity, provided the power supply can provide the necessary current. As a rule of
thumb no more than 2 amperes per 100 ml of electrolyte must be passed through a chlorate cell. Under more optimal
conditions a higher amperage may be tolarable, still maintaining reasonable efficiency whereas in less optimal conditions 2
amperes may be too high and a lot of chlorine will be lost, leading to lower efficiency and rising pH. A current of 2 amperes will
convert approximately 0.73 gram of sodium chloride to 1.32g of sodium chlorate per hour (assuming 100% efficiency). After
extracting, metathesis reactions and recrystallising that will yield 1.53 g potassium chlorate. So, for example, to produce 100
grams of potassium chlorate a day at least 100 grams / 1.53 grams / 24 hours * 100 ml = 272 ml of electrolyte are required.
To maintain that rate of conversion the cell will then require 272 ml / 100 ml * 2 amperes = 5.44 amperes. If a cell is less
efficient than 100%, which every cell is, increase these figures proportionally (so at 50% efficiency: 100% / 50 % * 272 ml =
544 ml of electrolyte, consuming 10.88 amperes of current to maintain the same rate of production). The example cell
described above contains 200 ml of electrolyte. Thus, it should be operated at a current of 4 amperes, and the maximum daily
yield is 100/272 * 200 = 74 g of potassium chlorate after processing the electrolyte. These figures were also mentioned in the
cell description without explanation.

2.5 Cell body materials

One of the main problems in chlorate cells is the corrosiveness of the electrolyte. Only very few materials do not corrode when
in contact with the electrolyte or its fumes. Most metals corrode, many plastics will and even glass does under some
circumstances.

Some metals, such as steel, can be used if they are protected from corrosion in some way. For that purpose it can be coated
with a resistant material such as teflon or some types of rubber, or it can be 'cathodically protected'. This means means the cell
walls are used as a cathode. The negative potential prevents the steel from being oxidised if the current density (current per unit
of surface area) on the steel is high enough.

Some metals, particularly titanium, zirconium, tantalum and niobium, form a protective film when they are in contact with the
electrolyte. This prevents them from further corrosion, and they therefore find extensive use in industrial setups (particularly
titanium). For amateurs the difficulties in working with these metals and their high price restricts their use somewhat. In small
scale setups glass and plastics such as PVC are more easily available, easier to work with and much cheaper.

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

The range of suitable electrode materials is very limited. Especially the anode material is critical. The positive charge on the
anode promotes oxidition and the evolving oxygen attacks many anode materials. Several anode materials have been
considered over the years. Todays main options are listed below along with a short description.

Anode materials

Graphite: graphite is cheap and easy to obtain. It does however corrode at a comparatively fast rate. This makes it necessary
to replace the anodes every so often and to filter the electrolyte before further processing which can be difficult and laborous
due to the small size of the carbon particles. Graphite is not suitable for making perchlorates. When the chloride concentration
of the electrolyte drops to the point where perchlorate formation begins (about 10% w/v), the graphite begins to oxidise at a
great rate, yielding no or only traces of perchlorate. Cells operating with graphite anodes must also be maintained at a relative
low temperature to limit anode erosion, which translates to a lower cell capacity. Graphite rods can be found in old manganese
dioxide-zinc batteries or in welding shops where they are sold as 'gouging rods'. They can be treated with linseed oil to reduce
corrosion. A practical method for the home experimenter has been devised by Rich Weaver, and is well described on Mike
Brown's page. Old battery electrodes do not need to be treated with linseed oil.

Platinum: The obvious disadvantage of platinum is its high price. However, platinum anodes corrode only at a very slow rate
and are suitable for perchlorate production. They therefore provide an almost ideal anode material. High efficiency can be
reached with platinum and processing of the electrolyte is greatly simplified.

Lead dioxide: Lead dioxide provides an economical alternative to platinum. Lead dioxide anodes can be made at home. This
takes some work and effort, but the anodes are cheap, fairly resistant to corrosion even at higher temperatures and are suitable
for perchlorate production. More information on lead dioxide electrodes is given elsewhere on this homepage.

Manganese dioxide: Another oxide that is conductive and resistant to oxidation. It is made by thermal decomposition of
manganese nitrate pasted onto a substrate. This type of anode seems quite promising for amateur pyro use. For more
information, the user is referenced to patents in the literature list below. If anyone has experiences with these anodes and their
preparation I'd be most interested to hear about them.

DSA: DSA stands for Dimensionally Stable Anode. This is the common term used to refer to anodes consisting of a layer of
noble metal oxides (usually RuO2 and TiO2) coated onto a substrate, usually titanium. This type of anode is finding increased
use in industrial cells because of its comparatively low cost when compared to platinum and its resistance to corrosion. Some of
the chemicals required to manufacture these anodes (particularly RuCl3 and tetra-butyl titanate, Ti(OBu)4) are expensive and
perhaps difficult to handle safely. However, if the chemicals can be obtained and suitable equipment is available, the procedure
to make the anode seems fairly straightforward and may be an option. For the preparation of these, the reader is referenced to
the literature. Again, I'd be most interested in anyones experiences with this type of anode.

Magnetite: This has found use in industry in the past, but is rarely used nowadays. It corrodes, but not very quickly and it can
be used for perchlorate manufacture. The anodes are made by melting and casting FeO.Fe2O3 into the required anode shapes.
I have little literature available on this material, so it is not further discussed here.

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

Although not essential for chlorate manufacture, controlling the temperature and pH will increase cell efficiency and therefore
the capacity of a cell. Temperature control can be anything from a sophisticated thermostat and heating element (or a cooling
element) to simple insulation around the cell or a cold water bath. As mentioned earlier, part of the electric energy is lost as heat
in the cell. Small cells operating at high currents can sometimes reach temperatures of 80 to 90 deg C. Though high
temperatures will improve efficiency, temperatures as high as that will also increase anode corrosion and it is therefore usually
considered better to maintain a temperature in the range 60 to 80 deg C to get the best of both worlds. Graphite anodes tend
to erode faster than other types though, especially at higher temperatures, and cells employing these are therefore usually
operated at 40 deg C to limit anode erosion.

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 just starting a first batch of chlorate a fresh electrolyte has to be prepared. When the cell has been operated before, the
electrolyte from the previous cell is available to prepare the electrolyte for a new batch. Also, the material left behind from the
extraction and purification steps can be added to the next cell as it may contain some residual chlorate. That way no product
left in the solution after procesing is lost.

As mentioned, it is common to use sodium chloride to prepare sodium chlorate first, which is then converted to potasisum or
barium chlorate later. Even though both compounds may be prepared directly from potassium or barium chloride, using sodium
chloride as a starting material has advantages. Mainly, it makes processing of the electrolyte much easier since sodium chlorate
is very soluble. It is therefore easily separated from insoluble impurities which are almost always present. It will be assumed that
sodium chloride is used. If for some reason the use of other compounds is desired, the procedure and amounts may need to be
adjusted.

Preparing fresh electrolyte

1. Prepare a saturated solution of sodium chloride. Take about 40 grams for every 100 ml of solution and bring the solution to
a boil. Then allow to cool to room temperature again. Some sodium chloride will crystalise as the solution cools. The solution is
then filtered to obtain a clear saturated solution.

2. Optionally, 2 to 4 g/l of potassium dichromate, potassium chromate, sodium chromate or sodium dichromate may be added
to improve efficiency. These compounds are suspected carcinogens, so if you choose to add any, know the hazards involved
and act accordingly. If lead dioxide anodes are used, do not add potassium dichromate as it will only reduce efficiency. Instead,
2 to 4 g/l of sodium or potassium fluoride may be used. Although not carcinogenic, the fluorides are nasty compounds as well
and should be handled properly.

3. Finally, the pH of the solution can be adjusted. A pH of around 6 is optimal, but anything between 5.5 and 6.5 is reasonable.
The pH can be increased by addition of sodium hydroxide solution and it can be decreased by adding hydrochloric acid. Do
not use too concentrated solutions for adjusting the pH. A concentration of 2% (w/v) for both solutions is convenient to work
with.

Recycling old 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

Voltage and current

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.


Mr Cool
Frequent Poster
Posts: 991
From: None of your bloody business!
Registered: DEC 2000
posted February 20, 2001 10:21 AM
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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

The electrolyte usually contains suspended solid particles, even though they are not always visible. Suspended particles can be
detected with the use of the Tyndall effect. Shine a bright flashlight through the side of a glass container containing the solution.
If no suspended particles are present the light beam cannot be seen going through the solution. If suspended particles are
present they will scatter the light and make the beam visible.

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

Depending on the starting material, sodium or calcium hypochlorite, a different procedure must be followed. Each is described
separately below.

Procedure when using sodium hypochlorite

It is assumed bleach will be used, which is usually a 4% solution of sodium hypochlorite in water. If a less or more concentrated
solution is used, adjust the amounts accordingly.

1. Take 1 liter of bleach, and place this in heat resistant glass or stainless steel container. Bring it to a boil.

2. Boil the solution untill only about 140 ml of solution is left. The exact volume is not critical, a deviation of 10 to 20 ml is
acceptable.

3. Allow the solution to cool. If crystals form upon cooling, filter the solution after it has completely cooled. The crystals are
sodium chloride and can be discarded.

4. In a separate container, prepare a solution of potassium chloride. Dissolve 28 grams of potassium chloride in the smallest
volume of water possible (about 80 ml). This can be done by dissolving the potassium chloride in about 90 ml of boiling 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.

5. Mix the boiled bleach solution with the potassium chloride solution. A white precipitate should form. This is potassium
chlorate.

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

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

8. 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.

Procedure when using calcium hypochlorite

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.
Back to main page


Mr Cool
Frequent Poster
Posts: 991
From: None of your bloody business!
Registered: DEC 2000
posted February 20, 2001 10:22 AM
--------------------------------------------------------------------------------
Preparing perchlorates
Chapter 1: Introduction
Chapter 2: Electrolytic preparation

2.1 2.1 cell construction
2.2 2.2 electrode materials
2.3 2.3 preparing the electrolyte
2.4 2.4 operating the cell
2.5 2.5 Processing the electrolyte
2.6 2.6 Purification

Chapter 3: Preparation by thermal decomposition of chlorate
Chapter 4: Preparation by chemical oxidation
Chapter 5: Literature

Chapter 1: Introduction

Industrially, perchlorates are exclusively prepared by the electrochemical method. In the past, thermal decomposition of
chlorate has been used but since this process is very inefficient it has been abandoned long ago. Chemical oxidation of chlorates
is currently not very economical either, but it may become an option in the future. For amateur pyros, each of these methods
may be used as we need not be concerned with the economy of the process as much as commercial operations do. The
electrochemical method is convenient to use if you have a chlorate cell with the right anodes already, since then making
perchlorates is simply a matter of operating the cell for a bit longer. If you can get chlorates in quantity for cheap, for example
as a herbicide, the thermal decomposition method is an option. The method is quick, and requires no chemicals other than the
chlorate starting material. Perchlorates can also be prepared by chemical oxidation of chlorates. The required chemicals are
relatively expensive, but the method is quick and simple.

Chapter 2: Electrolytic preparation

Chlorates can be oxidised in an electrochemical cell to yield perchlorates. The prefered starting material for this method is
sodium chlorate, since it is very soluble. Potassium chlorate is seldom used due to its low solubility, and ammonium chlorate
should never be used since it leads to the formation of sensitive and explosive NCl3 in the cell. Sodium perchlorate is
conveniently converted in high yield to a number of other perchlorates (such as potassium and ammonium perchlorate) by
double decomposition (metathesis) reactions. It is assumed from here on that sodium chlorate is used as starting material.

Cell construction

Electrochemical cells for perchlorate synthesis do not differ much from chlorate cells. The most important difference lies in the
anode material. Not all anode materials suitable for chlorate synthesis can also be used for perchlorate synthesis. Most cell
body materials used in chlorate cells may also be used in perchlorate cells. The effect of temperature and pH deviating from the
optimal values is of much less importance in perchlorate cells. Finally, the voltage at which perchlorate cells operate is
somewhat higher because the potential at which the conversion reactions take place is higher. The general structure of both cell
types is the same: two working electrodes, and no diaphragm.

Electrode materials

Like in chlorate cells, stainless steel is a suitable cathode material. Mild steel may also be used. While copper and brass will
also work, they may cause problems with copper contamination when they erode.

Anode materials for perchlorate cells should have a high oxygen overpotential. What exactly that means is not further discussed
here; it suffices to say that if the oxygen overpotential at a certain anode material is not high enough oxygen will be evolved
instead of chlorate oxidised to perchlorate. No perchlorate will be formed, and the anode material is usually attacked
comparatively quickly. This holds also for chlorate cells, but the problem is less severe there since lower potentials are involved.
Anode materials suitable for perchlorate synthesis are listed below. These are also described in the chlorate synthesis section
but are repeated here for convenience.

Platinum: The obvious disadvantage of platinum is its high price. However, it corrodes only at a very slow rate and therefore
provides an almost ideal anode material. High efficiency can be reached with platinum and processing of the electrolyte is
greatly simplified due to the absence of insoluble anode erosion products.

Lead dioxide: Lead dioxide provides an economical alternative to platinum. Efficiency of lead dioxide anode based cells is
usually slightly lower than that of platinum based cells, but the difference is small. Lead dioxide anodes are not easily bought and
must be made. This takes some work and effort, but the anodes are cheap, farily resistant to corrosion even at higher
temperatures. More information on lead dioxide anodes of several types is given elsewhere on this homepage.

DSA: DSA stands for Dimensionally Stable Anode. This is the common term used to refer to anodes consisting of a layer of
noble metal oxides (usually RuO2 and TiO2) coated onto a substrate, usually titanium. This type of anode is finding increased
use in industrial cells because of its comparatively low cost when compared to platinum and its resistance to corrosion. The
chemicals required to manufacture these anodes are expensive and difficult to handle. However, if the chemicals can be
obtained and suitable equipment is available, the procedure to make the anode seems fairly straightforward and may be an
option. For the preparation of these, the reader is referenced to the literature. Again, I'd be most interested in anyones
experiences with this type of anode.

Magnetite: This material has found use in industry in the past, but is rarely used nowadays due to its relatively high corrossion
rate and low efficiency for perchlorate manufacture. The anodes are made by melting and casting FeO.Fe2O3 into the required
anode shapes. I have little literature available on this material, so it is not further discussed here.

Preparing the electrolyte

Sodium perchlorate can be made directly from sodium chloride by electrolysis in which case no special electrolyte for the
chlorate to perchlorate step has to be prepared. The preparation of a chloride electrolyte is described in the text on chlorates.

A cell can also be operated purely for the chlorate to perchlorate conversion. An electrolyte has to be prepared before each
batch in this case. If the cell has been operated before, it is best to 'recycle' the old electrolyte and all the impure fractions
obtained during extraction and purification of the product. That way no product is wasted. When the cell is operated for the
first time, a fresh electrolyte has to be prepared which can be done as follows.

Preparing fresh electrolyte

1. Prepare a saturated solution of sodium chlorate. Take about 60 grams of sodium chlorate for every 100 ml of solution and
bring the solution to a boil. Then allow to cool to room temperature again. Sodium chlorate will crystalise as the solution cools.
The solution is then filtered to obtain a clear saturated solution.

2. Optionally, 2 to 4 g/l of potassium dichromate, potassium chromate, sodium chromate or sodium dichromate may be added
to improve efficiency. These compounds are suspected carcinogens, so if you choose to add any, know the hazards involved
and act accordingly. If lead dioxide anodes are used, do not add potassium dichromate as it will only reduce efficiency. Instead,
2 to 4 g/l of sodium or potassium fluoride may be used. Although not carcinogenic, the fluorides are nasty compounds as well
and should be handled properly.

Recycling old electrolyte

1. If the electrolyte is not clear but contains suspended particles, remove these by filtration.

2. Dissolve any impure material left over from purification steps.

3. Re-saturate the solution with sodium chlorate, following the same procedure as described in step 1 of 'preparing a fresh
electrolyte'.

4. Like described in step 2 of 'preparing a fresh electrolyte' you may now add dichromates, chromates or fluorides if you
choose to do so.

Operating the cell

Perchlorate cells are operated at a higher voltage than chlorate cells and temperature and pH do not need to be controlled
within strict limits for optimal efficiency. Other than that, operation is much like that of chlorate cells.

Like explained in the theory section on chlorate cells, the voltage over a cell may fluctuate if the current is kept constant. In
typical chlorate cells this results in a cell voltage of 3 to 4 volts, whereas in perchlorate cells the voltage is higher, 5 to 7 volts
usually. The current is kept constant at an acceptable level with respect to anode erosion. A maximum current per volume as
exists in chlorate cells does not exist in perchlorate cells. The current could in theory be increased indefinately to increase the
reaction rate, were it not that anode erosion increases with increasing current density (the current per unit of anode surface
area). The current is therefore usually set by the surface area of the anode. As a rule of thumb maintain a current density of 200
mA/cm2.

The influence of cell temperature is two-fold: anode erosion increases with increasing cell temperature and the cell voltage is
reduced at higher temperatures. The former is obviously unwanted, while a lower cell voltage means energy is saved. In
industry the temperature is ofcourse chosen to get the best of both worlds, depending on what is more expensive: energy or
anodes. We need not be concerned with economy as industry does so the temperature does not matter a lot. Try to keep it
between 40 and 80 deg C. When using lead dioxide anodes, it is probably best to prevent the temperature from going very
high. It can make the lead dioxide crack.

Finally, some wat er should be added from time to time to make up for what has evaporated. Try to maintain a constant
electrolyte volume.

Running times

The required time to operate a cell depends on the current. The higher the current, the less time needed. In fact, the current is a
measure of the reaction rate. Therefore, the amount of electricity that went through a cell is calculated by multiplying the current
going through the cell (in amperes) by the time it has been flowing (in hours). The resulting number is measured in
amperage-hours (abbreviated as Ah). To convert 100 grams of sodium to sodium perchlorate 50 Ah are required if the cell
operates at 100% efficiency. In real life a cell will never reach 100% efficiency, and more electricity is needed.

Example: A 200 ml cell contains initially about 100 grams of sodium chlorate (the solubility of NaClO3 in water is about 50
g/100ml at room temperature). Per 100 grams, 50 Ah are needed. So, if a cell operates at a current of 2 amperes, it would
take 50/2 = 25 hours to convert all chlorate to perchlorate in a cell operating at 100% efficiency. If the cell actually operates at
80% efficiency, 100/80 * 25 = 31.25 hours (or 31 hours, 15 minutes) are needed.

If perchlorate is prepared by continuing to run a chlorate cell after all chloride has been consumed, the total run time is ofcourse
the sum of the time required for the chloride to chlorate conversion and the chlorate to perchlorate conversion.

Processing the electrolyte

When done, raw potassium perchlorate or ammonium perchlorate may be prepared from the electrolyte as follows:

Filtering

The first step is filtration. The electrolyte usually contains suspended solid particles. These 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. However, even that 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.

Chromate and dichromate removal

If chromates or dichromates were used to increase the cells efficiency they should now be removed. Adding a solution of
barium chloride to the electrolyte will precipitate any chromate or dichromate as the corresponding barium compounds. Add
small amounts of a 10% barium chloride solution to the electrolyte. A yellow precipitate will form. If no more yellow precipitate
is formed, filter to remove the barium compounds. A white (sometimes clearly crystalline) precipitate may form instead of a
yellow precipitate. This is barium chlorate or perchlorate. If this happens, do not add any more barium chloride solution and
filter to remove the precipitate.

Destruction of chlorate

The next step is the desctruction of residual chlorate. Even when a cell is operated for many times the required running time,
some chlorate is still present in the electrolyte. Since chlorates and perchlorates behave very differently in pyrotechnic
compositions (see the safety page among others) chlorate contamination can be very dangerous. If the cell is not operated long
enough to convert most of the chlorate to perchlorate the chlorate contamination may be very severe and it is unpractical and
very inefficient to attemp to destroy it all. If this is the case I suggest tthis step is skipped, and the raw product is extracted.
Ammonium perchlorate cannot be made this way, as it would result in the formation of the dangerously usntable explosive
compound ammonium chlorate (also see the safety page). Conversion to potassium perchlorate may however be tried. This will
ofcourse afford a heavily chlorate contaminted batch of perchlorate, probably even something that can better be considered to
be a chlorate/perchlorate mixture or perchlorate contaminated chlorate. This material could be used to an extremely limited
extend in pyrotechnic compositions when it is treated as a chlorate, or, a better option, it can be used as a starting material for
the thermal decomposition or chemical oxidation preparations of potassium perchlorate described later.

Residual chlorate is usually destroyed by the action of reducing agents. Sodium sulfite is used for this purpose in industrial
setups. Iron(II)sulfate is another option, and the chemical can be bought in some gardening supply shops as it is used to supply
plants with iron. A third method destroys chlorates by the action of strong acids. The cell electrolyte is acidified with
hydrochloric acid to a pH of 1 to 2, and the solution is boiled. Chlorates will decompose and yield a yellow gas, chlorine
dioxide. The gas will partially dissolve in the solution, imparting a bright yellow color to it. In high concentrations chlorine
dioxide is dangerously explosive and sensitive but if done using proper ventilation the small amounts evolved in this reaction are
very unlikely to cause dangerous levels. The gas is however quite toxic and inhalation should be avoided at all times. Never do
this step inside unless a well functioning fumehood can be used. After boiling for 15 minutes raise the pH to around 8 or 9 again
using sodium hydroxide. This should afford a colorless solution. A qualitative test for chlorate should now be performed to
make sure all chlorate has properly been destroyed. Such tests are described on the product analysis page.

If any other method is used to destroy residual chlorate the pH must always be adjusted afterwards to slightly above 7.
Otherwise, traces of acid may be incorporated into the product in later steps which can make it very dangerous to use in
pyrotechnic compositions.

Double decomposition

A decission will have to be made at this point wheter the intended product is ammonium or potassium perchlorate. Other
perchlorates can be made as well but are not discussed here as they find very little use in pyrotechnics.

If residual chlorate was not destroyed the choice is simple since ammonium perchlorate is not an option. Ammonium chlorate
could be formed in the process which is, as mentioned several times already, a dangerously unstable explosive compound the
formation of which should at all times be avoided (even in trace amounts). If chlorate was properly destroyed, and a qualitative
test indicates so, ammonium perchlorate may be prepared. 1. If potassium perchlorate is the intended product, take 70 grams
of potassium chloride for every 100 grams of sodium chlorate that was started with and dissolve this in the smallest volume of
water possible. If instead ammonium perchlorate is required, take 50 grams of ammonium chloride and dissolve in the smallest
volume of water possible.

2. Mix this solution of either potassium or ammonium chloride with the electrolyte. A white precipitate of the corresponding
perchlorate should form.

3. Boil the solution and add small amounts of water untill all the perchlorate has dissolved. Due to the low solubility of
potassium perchlorate a large volume of water may be needed then. If a sufficiently large container is not available the solution
may be split up in several portions that are later recombined.

4. When all has dissolved, check the pH of the solution. It should be neutral or slightly alcaline (above 7). If it is not, add some
dilute sodium hydroxide solution to increase the pH to between 7 and 8. When this value is overshot, hydrochloric acid may be
used to lower the pH again.

5. Allow the solution to cool slowly to room temperature. The perchlorate will crystallise during this. Cool the solution further to
0 deg C, and filter. Rinse the crystals in the filter with some ice-cold water. This raw product may be further purified as
described below.

Purification

The product can be purified by recrystallisation. This method is especially suitable for potassium perchlorate due to its low
solubility. Little prodcut will be lost, and the purity is greatly increased. Ammonium perchlorate suffers slightly worse losses
when recrystallised, but still acceptable. The impure ammonium perchlorate should be discarded since recycling could result in
NCl3 formation in the cells. The losses occuring when potassium perchlorate is recrystallised are so slight recycling is hardly
worth the effort (a liter of recrystallisation solution contains only a few grams of perchlorate). Recrystallising is done as follows:

1. Place the crude product in a pan and add 100 ml of water for every 20 g of raw potassium perchlorate, or 100 ml of water
for every 50 gram of raw ammonium perchlorate. Bring this to a boil.

2. After it has boiled for a few minutes, add 10 ml amounts of water to the boiling solution in 5 minite intervals untill all the
product 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 ammonium perchlorate is the intended product, use ammonia instead. If this is not
done, traces of acid may be included in the final product making it very dangerous to use in pyrotechnic compositions.

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

5. Filter and rinse the crystals well with ice cold water. The filtrate should be discarded

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

Chapter 3: Preparation by thermal decomposition of chlorate

Chlorates are thermodynamically unstable. Over time they will decompose into a mixture of chlorides and perchlorates. In
essence, the chlorate will undergo a redox reaction with itself (a so-called auto-oxidation reaction). Due to the kinetic stability
of chlorates however, the process is slow at room temperature. A well known example of the same phenomenon is diamond:
Diamonds are unstable at common pressures and temperatures. They turn into graphite extremely slowly. If a diamond is
heated, the process is sped up. Similarly, if chlorates are heated the reaction rate is increased enough for it to be used as
reaction pathway in the preparation of perchlorates. When potassium chlorate is used, the resulting perchlorate is easily
separated from the chloride by recrystallisation. There may be several other chlorates that this method can be used with, but
potassium chlorate seems to be the most well investigated option. In any case, ammonium perchlorate cannot be prepared
directly with this method due to the explosive and instable nature of ammonium chlorate, as mentioned earlier. In the first large
scale perchlorate plants this method was used to prepare potassium perchlorate. Ideally, this preparation is performed in an
oven since it involves heating the chlorate for several hours. The chlorate used should be free of impurities that catalyse chlorate
decomposition (such as most d-block metals). The following description assumes potassium chlorate will be used as a starting
material.

1. Heat pure potassium chlorate to slightly over its melting point. A colorless clear liquid is obtained. Before heating make sure
no organic material or other fuels are present in the chlorate or able to fall into the molten chlorate. This would result in a violent
reaction. Do not overheat since this will decompose the chlorate, yielding chloride only.

2. Maintain this temperature for several hours. During this time the potassium chlorate will undergo the auto-oxidation reaction.
Due to the higher melting point of potassium chloride and potassium perchlorate the melt will solidify slowly during this time. It
becomes quite hard to judge the correct temperature as the melt solidifies, and if an oven is not used a thermometer is essential
to judge the correct temperature. Too high a temperature will cause the perchlorate to decompose, a low temperature will
result in incomplete conversion.

3. After the mixture has completely solidified, allow it to cool to room temperature. Test a sample of the cooled residue for
chlorate, as described in the analysis page. If chlorate is present it needs to be destroyed before extraction of the perchlorate.
Destruction of chlorate is described earlier in this text, in the section dealing with processing the electrolyte from the electrolytic
preparation of perchlorates.

4. Recrystallise the residue, as described earlier. Two recrystallisation steps are sometimes needed to separate the potassium
perchlorate completely from the chloride as there is quite a lot present.

Chapter 4: Preparation by chemical oxidation

A third method to convert chlorates to perchlorates is by chemical oxidation. A sufficiently strong oxidiser added to a chlorate
can oxidise a chlorate to a perchlorate. Suitable oxidisers are persulfates and lead dioxide in concentrated sulfuric acid.
Hydrogen peroxide does not seem to work. I must admit my experience with this method is limited and I have not been able to
obtain much literature about it. It is mentioned in literature however, and it is definately a possible method that can be used with
good results. If anyone has some experience with it, I would be most interested to hear about it. The experiments I have
conducted involved sodium persulfate as the oxidiser. I will give an account of the general method I used, the ideas behind it
and the results obtained with it. Anyone with comments, results, ideas, anything is very welcome to comment on it. I can
currently not do any further experiments untill july/august this year. Soon after that, this text should be updated. If you wish to
try this method in the meantime this may be used as a starting point:

Theory

Persulfates are strong oxidisers. They are reduced according to the following half-reaction:

[ REACTION ]

The H+ generated in this reaction will prevent it from taking place below a certain pH. When a base is added to neutralise the
acid generated the reaction may go to completion. All persulfate may be consumed in the reaction. Persulfate being a stronger
oxidiser than perchlorate may be used to convert chlorates to perchlorates according to the following half reaction:

[ REACTION ]

Neither persulfates, chlorates or perchlorates are destroyed at a significant rate at the temperature of boiling water, so the
reaction rate may be increased by boiling a solution containing the reactants.

Oxidation of KClO3 by persulfate

The general procedure I tried:

A solution of 30 g/l sodium persulfate was prepared and the pH raised to 14 by the addition of a concentrated sodium
hydroxide solution. The sodium persulfate was obtained from an electronics supply store, where it was sold for etching printed
circuit boards. A foggy solution was obtained, which was filtered. 5 ml of the clear solution obtained after filtering was added to
a test tube. Approximately 1 gram of potassium chlorate was added to the same tube, and solution heated. The solution was
boiled vigorously for 15 minutes during which time water was added occasionally to make up for what had evaporated. The
solution was then allowed to cool. Upon cooling white crystals formed. These were filtered, washed with ice cold water and
recrystallised. The crystal shape during recrystallisation was observed and found to resemble that of potassium perchlorate
best. As described in the analysis page, crystal shape is not a reliable way to determine the identity of a product however. The
crystals obtained after recrystallisation were tested qualitatively for chlorate with phenylanthranilic acid (as described in the
analysis page). Chlorate was shown to be present. It seems most likely that the chlorate was only partially converted to
perchlorate. Maybe the addition of more persulfate or allowing a longer reaction time will convert more chlorate to perchlorate.
Destruction of the chlorate followed by recrystallisation should afford a chlorate free product.

Chapter 5: Literature

1. Schumacher, J.C., "Perchlorates", New York, Reinhold Publishing Corp., 1960

2. Remy, H. "Treatise on Inorganic Chemistry", New York, Elsevier Publishing Co., 1956

3. Mike Brown's homepage

list under construction

Copyright Wouter Visser, May 1999.
Back to main page


Mr Cool
Frequent Poster
Posts: 991
From: None of your bloody business!
Registered: DEC 2000
posted February 20, 2001 10:23 AM
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I hope some of that helped.


PYRO500
Moderator
Posts: 1466
From: somewhere in florida
Registered: SEP 2000
posted February 20, 2001 09:22 PM
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yep it did, when I get well (I think I have pneumonia, and I know I have strep throat) I will go to the junk yard and get a mw oven transformer, I'm still thinking of a way to current limit it though


Jhonbus
Frequent Poster
Posts: 345
From:
Registered: SEP 2000
posted February 21, 2001 11:23 AM
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I current limited a MOT with another MOT that had the secondary shorted. It got quite hot in continuous shorting of the other MOT though (Jacob's Ladder) so I put it under oil....which boiled


firebreether
Frequent Poster
Posts: 108
From:
Registered: NOV 2000
posted February 22, 2001 04:05 PM
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Could Al be used as an anode or cathode since it forms a protective layer of oxide when electrolized? I dont know if the oxide(Al2O3) conducts. The reason I ask this is that it says titanium is the perfect cathode ( the positive one right?) because of its low wear rate because of the protective oxide layer it forms around itself. Aluminum forms this layer so could it be used? Also many people use a stainless steel cathod, could this be as simple as a kitchen butter fork that is stainless? If it isn't optimal no big deal, not like I would use that anyway
Also, maybe this is why PbO2 anodes have such little wear.