Potassium Chlorate

UC · January 14, 2009

Preparation of Chlorates by Disproportionation of Calcium Hypochlorite in Aqueous Solution

Aside from their famous (perhaps infamous is more appropriate) use in pyrotechnics, chlorates have a fairly rich chemistry and make a valuable addition to the chemical inventory of the amateur experimenter. Iodic acid and iodates can be prepared by the reaction of iodine with solutions of chlorate acidified by nitric acid. Reduction of chlorates with hydrochloric acid generates ClO2 (generally mixed with varying amounts of Cl2 depending on stoichiometry), which makes for an interesting small scale demonstration (Do not scale up!). They can be used in place of nitrates as the oxidizer for the synthesis of chromates and permanganates by molten salt fusions and they can serve as heat boosting agents in exotic metal thermites. Perhaps most useful to the amateur chemist, however, is that chlorates can be catalytically decomposed with heating and small quantities of MnO2 to provide an ever ready and cheap source of oxygen gas without the need for a pressurized cylinder. The industrial preparation and essentially the only viable option for large scale preparation of chlorates is by electrolysis of sodium or potassium chloride solutions. For microscale and qualitative experiments, however, building a chlorate cell is time consuming and downright overkill for the quantities desired. An alternative procedure for the amateur chemist involves boiling down bleach, a rather odorous and time consuming affair, which tends to leave glassware etched for fairly small quantities of chlorate. Here, I present a simple and clean procedure for preparing potassium chlorate in small quantities from cheap and widely available calcium hypochlorite pool shock.

Safety

It should be more than apparent to anyone with even limited knowledge of chlorates that these compounds are dangerous. As a strong oxidizing agent, contamination of chlorates by organic material or other reducing agents can lead to fires or explosions. Spills of solutions soaked into wood or paper can later on accelerate fires. Acidification of chlorates, especially by hydrochloric acid (or by co-acidification with chlorides), can lead to the evolution of chlorine and chlorine dioxide gas. The former is a powerful respiratory irritant and inhalation may cause pulmonary edema. The latter is capable of spontaneous detonation when it reaches certain concentrations in air and is a severe respiratory irritant in lower concentrations. Contact of molten chlorates with organic material can lead to explosions. Chlorates are highly toxic by ingestion and are skin, eye, and respiratory tract irritants. They are used as nonspecific defoliants and weed killers, and are thus highly toxic to plant life. I urge you, at the bare minimum, to use gloves when working with any chlorates, and preferably goggles and a lab coat as well. As for the procedure in general, you are working with a close analog of bleach, and spills can and will damage or destroy clothing in the same way that bleach spills will. Contact with any part of your body should be avoided and the entire procedure should be carried out in a location with good ventilation as small amounts of chlorine gas are inevitably evolved during the procedure.

Legalese

Manufacture of any pyrotechnic device or explosive with the product of this procedure may or may not be legal, depending on your location and licensing. The manufacture of a pyrotechnics oxidizer (the completion of this procedure) also may or may not be legal in your area. All preparations are carried out at your own risk with respect to personal safety and legality. Neither Science Forums, nor the author will be held responsible for any and all damages that may be incurred as a result of following this procedure. By undertaking this experiment, you understand the risks associated with it (including any which may not be explicitly stated in the safety section) and agree to the above terms.

Required Chemicals:

Calcium hypochlorite (usually available with around 50% active material as pool shock)

Distilled water

Potassium chloride

An acid (anything even moderately acidic will work)

A reducing agent: a sulfite, bisulfite, metabisulfite (pyrosulfite), dithionite (hydrosulfite), or thiosulfate (hyposulfite) are all acceptable, with the first three being preferred.

Equipment:

Heat source sufficient to boil water

Stir rod

Beaker(s) and watch glass

Filtering Setup (vacuum filtration greatly preferred)

Ice/brine bath

Experimental

Into a 1 liter beaker containing 500mL of distilled water, 100.0g of 54.6% calcium hypochlorite pool shock was slowly added with stirring. The hypochlorite foams somewhat as it is added (hence the stirring) and small amounts of chlorine or chlorine oxides are evolved, giving rise to the familiar pool chemical smell. The liquid is a pale green-yellow, the same color as bleach, indicating dissolved hypochlorite ion. The beaker was covered with the watch glass to exclude as much air as possible. Heating was carried out in a microwave oven on maximum power. The mixture was watched at all times as it is highly prone to “boiling” over, especially at the beginning. The solution was brought to a hard boil and was kept boiling for 20 minutes. During this time the solution reduced to about half of its initial volume and the small pellets of hypochlorite broke down into a pasty suspension consisting (most likely) of calcium hydroxide. The smell of chlorine is detectable for the majority of the heating time.

The solution was vacuum filtered after cooling. Around 235mL of clear filtrate was collected. The solids were relatively dry and were discarded. The filtrate was transferred to a 600mL beaker and 19.0g of potassium chloride was added. This was covered with the watch glass and boiled until the solution volume was about 130mL. During the boiling, the solution became cloudy, likely due to the reaction of some traces of calcium hypochlorite with carbon dioxide in the air. The solution was gravity filtered while hot through a loose cotton plug to remove the cloudiness and placed in a salted ice bath. Potassium chlorate separated as flat, colorless, glittering plates which appeared more feathery in form when cooling was rapid.

The crystals were brought into suspension by mixing with a stir rod and were vacuum filtered. They were washed with two 25mL portions of ice cold distilled water and air was drawn over them for several minutes to dry them as much as possible. They were dried by leaving at room temperature for about 24 hours (the solid is not at all hygroscopic and readily dries without the need for a dessicator) to afford 15.2g of odorless, free-flowing, glittering plates. This was 48.7% of the theoretical yield based on calcium hypochlorite.

Waste Disposal

As stated in the safety section, chlorates are very toxic to all forms of life. They are also a very persistant drinking water contaminant. To prepare the liquid wastes for disposal, a spatula of the chosen reducing agent is added and acid is dripped in. The products of acidifying chlorates; chloric acid, chlorine dioxide, and chlorine, and readily reduced by free sulfur dioxide released in the acid environment. The waste is very carefully smelled. If sulfur dioxide is detectable, all chlorate in the waste has been reduced to chloride. If not, more reducing agent and acid are added. Afterwards, the waste can be flushed down the drain with no toxic effects.

Discussion

The reaction of core importance in this experiment is a disproportionation reaction. In this kind of reaction, one reactant is both oxidized and reduced. For this experiment, the reactant in question is the hypochlorite ion, which reacts as follows:

3ClO- => 2Cl- + ClO3-

The hypochlorite ion is stabilized in the presence of base and by low temperatures. Boiling the solution provides the necessary energy to drive the reaction to the right. The competing reaction is the decomposition of calcium hypochlorite:

Ca(OCl)2 => CaCl2 + O2

This reaction is notably observed when heating dry calcium hypochlorite. No chlorate is produced, only oxygen gas. The calcium hypochlorite in pool shock is probably prepared by exposing damp calcium hydroxide to chlorine gas, then mixing in some extra hydroxide to stabilize it. The product of this preparation is probably a mixture of calcium hypochlorite and basic calcium hypochlorite (Ca(OH)(OCl)), the latter which is probably only weakly soluble in water. On the scale of a swimming pool, it will dissolve, but not in 500mL of water. It is likely that this undissolved solid undergoes the above decomposition reaction when boiled instead of the disproportionation reaction. The evolution of oxygen is probably why the solution seems to “boil” when heating starts and why it seems very vigorous. The label of the pool shock, however need not distinguish between the two compounds as both will work in a pool. The meager 48.7% yield is thus likely a misrepresentation, as it should be based only on Ca(OCl)2 content.

Calcium chloride and chlorate are both extremely soluble. Potassium chloride is fairly soluble, but potassium chlorate is only weakly soluble at low temperatures. When the solution containing potassium ions, calcium ions, chloride ions, and chlorate ions is cooled, the least soluble combination is potassium chlorate, which crystallizes out. This is an example of a metathesis reaction (double displacement).

2KCl (aq) + Ca(ClO3)2 (aq) => CaCl2 (aq) + KClO3 (s)

Another reaction relevant to the procedure is the reaction of calcium hypochlorite with carbon dioxide and then the reaction with chlorides in solution:

Ca(OCl)2 + CO2 + H2O => CaCO3 + 2HOCl

HOCl + Cl- => Cl2 + OH-

This produces free hypochlorous acid, which is highly unstable. This may then react with chlorides, allowing chlorine to escape solution. This is especially relevant for calcium chloride, since the solution is already saturated with calcium hydroxide. The precipitation of Ca(OH)2 drives the release of chlorine, which may escape (especially when it forms near the surface, where CO2 is likely to be, and especially because the solution is boiling) before reacting with the excess of Ca(OH)2 in solution. The beakers used to heat the solutions in this experiment end up with a thin film of calcium carbonate on them from this reaction. It appears to be etching at first, but can be removed by vigorous scrubbing or by acids. Aside from stopping splatters, the watch glass is an attempt to minimize how much CO2 gets into the solution.

References and various links:

Initial discussion on sciencemadness.org:

http://www.sciencemadness.org/talk/viewthread.php?fid=2&tid=10727&action=printable

Preparation of chromates with chlorates:

http://webpages.charter.net/dawill/tmoranwms/Chem_Chromate.html

MSDS for potassium chlorate:

http://www.jtbaker.com/msds/englishhtml/p5620.htm

What not to do with chlorates:

http://www.destructve.com/bromicacid/mistakes.htm#4

Chlorine dioxide demonstration by Woelen:

http://81.207.88.128/science/chem/exps/clo2/index.html

Destruction of Chlorates: