POTASSIUM

Occurrence in nature. Potassium is a constituent of many common rocks and minerals, and is therefore a rather abundant element, though not so abundant as sodium. Feldspar, which occurs both by itself and as a constituent of granite, contains considerable potassium. The element is a constituent of all clay and of mica and also occurs in very large deposits at Stassfurt, Germany, in the form of the chloride and sulphate, associated with compounds of sodium and magnesium. In small quantities it is found as nitrate and in many other forms.

The natural decomposition of rocks containing potassium gives rise to various compounds of the element in all fertile soils. Its soluble compounds are absorbed by growing plants and built up into complex vegetable substances; when these are burned the potassium remains in the ash in the form of the carbonate. Crude carbonate obtained from wood ashes was formerly the chief source of potassium compounds; they are now mostly prepared from the salts of the Stassfurt deposits.

Stassfurt salts. These salts form very extensive deposits in middle and north Germany, the most noted locality for working them being at Stassfurt. The deposits are very thick and rest upon an enormous layer of common salt. They are in the form of a series of strata, each consisting largely of a single mineral salt. A cross section of these deposits is shown in Fig. 78. While these strata are salts from a chemical standpoint, they are as solid and hard as many kinds of stone, and are mined as stone or coal would be. Since the strata differ in general appearance, each can be mined separately, and the various minerals can be worked up by methods adapted to each particular case. The chief minerals of commercial importance in these deposits are the following:

Sylvine	KCl.
Anhydrite	CaSO₄.
Carnallite	KCl·MgCl₂·6H₂O.
Kainite	K₂SO₄·MgSO₄·MgCl₂·6H₂O.
Polyhalite	K₂SO₄·MgSO₄·2CaSO₄·2H₂O.
Kieserite	MgSO₄·H₂O.
Schönite	K₂SO₄·MgSO₄·6H₂O.

Preparation and properties. The metal is prepared by the same method used in the preparation of sodium. In most respects it is very similar to sodium, the chief difference being that it is even more energetic in its action upon other substances. The freshly cut, bright surface instantly becomes dim through oxidation by the air. It decomposes water very vigorously, the heat of reaction being sufficient to ignite the hydrogen evolved. It is somewhat lighter than sodium and is preserved under gasoline.

Fig. 78

Potassium hydroxide (caustic potash) (KOH). Potassium hydroxide is prepared by methods exactly similar to those used in the preparation of sodium hydroxide, which compound it closely resembles in both physical and chemical properties. It is not used to any very great extent, being replaced by the cheaper sodium hydroxide.

Action of the halogen elements on potassium hydroxide. When any one of the three halogen elements—chlorine, bromine, and iodine—is added to a solution of potassium hydroxide a reaction takes place, the nature of which depends upon the conditions of the experiment. Thus, when chlorine is passed into a cold dilute solution of potassium hydroxide the reaction expressed by the following equation takes place:

(1) 2KOH + 2Cl = KCl + KClO + H₂O.

If the solution of hydroxide is concentrated and hot, on the other hand, the potassium hypochlorite formed according to equation (1) breaks down as fast as formed:

(2) 3KClO = KClO₃ + 2KCl.

Equation (1), after being multiplied by 3, may be combined with equation (2), giving the following:

(3) 6KOH + 6Cl = 5KCl + KClO₃ + 3H₂O.

This represents in a single equation the action of chlorine on hot, concentrated solutions of potassium hydroxide. By means of these reactions one can prepare potassium chloride, potassium hypochlorite, and potassium chlorate. By substituting bromine or iodine for chlorine the corresponding compounds of these elements are obtained. Some of these compounds can be obtained in cheaper ways.

If the halogen element is added to a solution of sodium hydroxide or calcium hydroxide, the reaction which takes place is exactly similar to that which takes place with potassium hydroxide. It is possible, therefore, to prepare in this way the sodium and calcium compounds corresponding to the potassium compounds given above.

Potassium chloride (KCl). This salt occurs in nature in sea water, in the mineral sylvine, and, combined with magnesium chloride, as carnallite (KCl·MgCl₂·6H₂O). It is prepared from carnallite by saturating boiling water with the mineral and allowing the solution to cool. The mineral decomposes while in solution, and the potassium chloride crystallizes out on cooling, while the very soluble magnesium chloride remains in solution. The salt is very similar to sodium chloride both in physical and chemical properties. It is used in the preparation of nearly all other potassium salts, and, together with potassium sulphate, is used as a fertilizer.

Potassium bromide (KBr). When bromine is added to a hot concentrated solution of potassium hydroxide there is formed a mixture of potassium bromide and potassium bromate in accordance with the reactions already discussed. There is no special use for the bromate, so the solution is evaporated to dryness, and the residue, consisting of a mixture of the bromate and bromide, is strongly heated. This changes the bromate to bromide, as follows:

KBrO₃ = KBr +3O.

The bromide is then crystallized from water, forming large colorless crystals. It is used in medicine and in photography.

Potassium iodide (KI). Potassium iodide may be made by exactly the same method as has just been described for the bromide, substituting iodine for bromine. It is more frequently made as follows. Iron filings are treated with iodine, forming the compound Fe₃I₈; on boiling this substance with potassium carbonate the reaction represented in the following equation occurs:

Fe₃I₈ + 4K₂CO₃ = Fe₃O₄ + 8KI + 4CO₂.

Potassium iodide finds its chief use in medicine.

Potassium chlorate (KClO₃). This salt, as has just been explained, can be made by the action of chlorine on strong potassium hydroxide solutions. The chief use of potassium chlorate is as an oxidizing agent in the manufacture of matches, fireworks, and explosives; it is also used in the preparation of oxygen and in medicine.

Commercial preparation. By referring to the reaction between chlorine and hot concentrated solutions of potassium hydroxide, it will be seen that only one molecule of potassium chlorate is formed from six molecules of potassium hydroxide. Partly because of this poor yield and partly because the potassium hydroxide is rather expensive, this process is not an economical one for the preparation of potassium chlorate. The commercial method is the following. Chlorine is passed into hot solutions of calcium hydroxide, a compound which is very cheap. The resulting calcium chloride and chlorate are both very soluble. To the solution of these salts potassium chloride is added, and as the solution cools the sparingly soluble potassium chlorate crystallizes out:

Ca(ClO₃)₂ + 2KCl = 2KClO₃ + CaCl₂.

Electro-chemical processes are also used.

Potassium nitrate (saltpeter) (KNO₃). This salt was formerly made by allowing animal refuse to decompose in the open air in the presence of wood ashes or earthy materials containing potassium. Under these conditions the nitrogen in the organic matter is in part converted into potassium nitrate, which was obtained by extracting the mass with water and evaporating to crystallization. This crude and slow process is now almost entirely replaced by a manufacturing process in which the potassium salt is made from Chili saltpeter:

NaNO₃ + KCl = NaCl + KNO₃.

This process has been made possible by the discovery of the Chili niter beds and the potassium chloride of the Stassfurt deposits.

The reaction depends for its success upon the apparently insignificant fact that sodium chloride is almost equally soluble in cold and hot water. All four factors in the equation are rather soluble in cold water, but in hot water sodium chloride is far less soluble than the other three. When hot saturated solutions of sodium nitrate and potassium chloride are brought together, sodium chloride precipitates and can be filtered off, leaving potassium nitrate in solution, together with some sodium chloride. On cooling, potassium nitrate crystallizes out, leaving small amounts of the other salts in solution.

Potassium nitrate is a colorless salt which forms very large crystals. It is stable in the air, and when heated is a good oxidizing agent, giving up oxygen quite readily. Its chief use is in the manufacture of gunpowder.

Gunpowder. The object sought for in the preparation of gunpowder is to secure a solid substance which will remain unchanged under ordinary conditions, but which will explode readily when ignited, evolving a large volume of gas. When a mixture of carbon and potassium nitrate is ignited a great deal of gas is formed, as will be seen from the equation

2KNO₃ + 3C = CO₂ + CO + N₂ + K₂CO₃.

By adding sulphur to the mixture the volume of gas formed in the explosion is considerably increased:

2KNO₃ + 3C + S = 3CO₂ + N₂ + K₂S.

Gunpowder is simply a mechanical mixture of these three substances in the proportion required for the above reaction. While the equation represents the principal reaction, other reactions also take place. The gases formed in the explosion, when measured under standard conditions, occupy about two hundred and eighty times the volume of the original powder. Potassium sulphide (K₂S) is a solid substance, and it is largely due to it that gunpowder gives off smoke and soot when it explodes. Smokeless powder consists of organic substances which, on explosion, give only colorless gases, and hence produce no smoke. Sodium nitrate is cheaper than potassium nitrate, but it is not adapted to the manufacture of the best grades of powder, since it is somewhat deliquescent and does not give up its oxygen so readily as does potassium nitrate. It is used, however, in the cheaper grades of powder, such as are employed for blasting.

Potassium cyanide (KCN). When animal matter containing nitrogen is heated with iron and potassium carbonate, complicated changes occur which result in the formation of a substance commonly called yellow prussiate of potash, which has the formula K₄FeC₆N₆. When this substance is heated with potassium, potassium cyanide is formed:

K₄FeC₆N₆ + 2 K = 6KCN + Fe.

Since sodium is much cheaper than potassium it is often used in place of it:

K₄FeC₆N₆ + 2Na = 4KCN + 2NaCN + Fe.

The mixture of cyanides so resulting serves most of the purposes of the pure salt. It is used very extensively in several metallurgical processes, particularly in the extraction of gold. Potassium cyanide is a white solid characterized by its poisonous properties, and must be used with extreme caution.

Potassium carbonate (potash) (K₂CO₃). This compound occurs in wood ashes in small quantities. It cannot be prepared by the Solvay process, since the acid carbonate is quite soluble in water, but is made by the Le Blanc process. Its chief use is in the manufacture of other potassium salts.

Other salts of potassium. Among the other salts of potassium frequently met with are the sulphate (K₂SO₄), the acid carbonate (KHCO₃), the acid sulphate (KHSO₄), and the acid sulphite (KHSO₃). These are all white solids.