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
2.11 purifying
the product

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. 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
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 is 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 it is 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, such as 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 because it is the cheapest). 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.

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