SODIUM

SYMBOLATOMIC WEIGHTDENSITYMELTING POINTFIRST PREPAREDLithiumLi7.030.59186.°Davy 1820SodiumNa23.050.9797.6°" 1807PotassiumK39.150.8762.5°" 1807RubidiumRb85.51.5238.5°Bunsen 1861CæsiumCs132.91.8826.5°" 1860

The family.The metals listed in the above table constitute the even family in Group I in the periodic arrangement of the elements, and therefore form a natural family. The name alkali metals is commonly applied to the family for the reason that the hydroxides of the most familiar members of the family, namely sodium and potassium, have long been called alkalis.

1.Occurrence.While none of these metals occur free in nature, their compounds are very widely distributed, being especially abundant in sea and mineral waters, in salt beds, and in many rocks. Only sodium and potassium occur in abundance, the others being rarely found in any considerable quantity.

2.Preparation.The metals are most conveniently prepared by the electrolysis of their fused hydroxides or chlorides, though it is possible to prepare them by reducing their oxides or carbonates with carbon.

3.Properties.They are soft, light metals, having low melting points and small densities, as is indicated in the table. Their melting points vary inversely with their atomic weights, while their densities (sodium excepted) vary directly with these. The pure metals have a silvery luster but tarnish at once when exposed to the air, owing to the formation of a film of oxide upon the surface of the metal. They are therefore preserved in some liquid, such as coal oil, which contains no oxygen. Because of their strong affinity for oxygen they decompose water with great ease, forming hydroxides and liberating hydrogen in accordance with the equation

M + H2O = MOH + H,

where M stands for any one of these metals. These hydroxides are white solids; they are readily soluble in water and possess very strong basic properties. These bases are nearly equal in strength, that is, they all dissociate in water to about the same extent.

4.Compounds.The alkali metals almost always act as univalent elements in the formation of compounds, the composition of which can be represented by such formulas as MH, MCl, MNO3, M2SO4, M3PO4. These compounds, when dissolved in water, dissociate in such a way as to form simple, univalent metallic ions which are colorless. With the exception of lithium these metals form very few insoluble compounds, so that it is not often that precipitates containing them are obtained. Only sodium and potassium will be studied in detail, since the other metals of the family are of relatively small importance.

The compounds of sodium and potassium are so similar in properties that they can be used interchangeably formost purposes. Other things being equal, the sodium compounds are prepared in preference to those of potassium, since they are cheaper. When a given sodium compound is deliquescent, or is so soluble that it is difficult to purify, the corresponding potassium compound is prepared in its stead, provided its properties are more desirable in these respects.

Occurrence in nature.Large deposits of sodium chloride have been found in various parts of the world, and the water of the ocean and of many lakes and springs contains notable quantities of it. The element also occurs as a constituent of many rocks and is therefore present in the soil formed by their disintegration. The mineral cryolite (Na3AlF6) is an important substance, and the nitrate, carbonate, and borate also occur in nature.

Preparation.In 1807 Sir Humphry Davy succeeded in preparing very small quantities of metallic sodium by the electrolysis of the fused hydroxide. On account of the cost of electrical energy it was for many years found more economical to prepare it by reducing the carbonate with carbon in accordance with the following equation:

Na2CO3+ 2C = 2Na + 3CO.

The cost of generating the electric current has been diminished to such an extent, however, that it is now more economical to prepare sodium by Davy's original method, namely, by the electrolysis of the fused hydroxide or chloride. When the chloride is used the process is difficult to manage, owing to the higher temperature required to keep the electrolyte fused, and because of the corroding action of the fused chloride upon the containing vessel.

SIR HUMPHRY DAVY (English) (1778-1829) Isolated sodium, lithium, potassium, barium, strontium, and calcium by means of electrolysis; demonstrated the elementary nature of chlorine; invented the safety lamp; discovered the stupefying effects of nitrous oxideSIR HUMPHRY DAVY (English) (1778-1829)Isolated sodium, lithium, potassium, barium, strontium, and calcium by means of electrolysis; demonstrated the elementary nature of chlorine; invented the safety lamp; discovered the stupefying effects of nitrous oxide

Technical preparation.The sodium hydroxide is melted in a cylindrical iron vessel (Fig. 76) through the bottom of which rises the cathodeK. The anodesA, several in number, are suspended around the cathode from above. A cylindrical vesselCfloats in the fused alkali directly over the cathode, and under this cap the sodium and hydrogen liberated at the cathode collect. The hydrogen escapes by lifting the cover, and the sodium, protected from the air by the hydrogen, is skimmed or drained off from time to time. Oxygen is set free upon the anode and escapes into the air through the openingsOwithout coming into contact with the sodium or hydrogen. This process is carried on extensively at Niagara Falls.

Fig. 76Fig. 76

Properties.Sodium is a silver-white metal about as heavy as water, and so soft that it can be molded easily by the fingers or pressed into wire. It is very active chemically, combining with most of the non-metallic elements, such as oxygen and chlorine, with great energy. It will often withdraw these elements from combination with other elements, and is thus able to decompose water and the oxides and chlorides of many metals.

Sodium peroxide(NaO). Since sodium is a univalent element we should expect it to form an oxide of the formula Na2O. While such an oxide can be prepared, the peroxide (NaO) is much better known. It is a yellowish-white powder made by burning sodium in air. Its chief use is as an oxidizing agent. When heated with oxidizable substances it gives up a part of its oxygen, as shown in the equation

2NaO = Na2O + O.

Water decomposes it in accordance with the equation

2NaO + 2H2O = 2NaOH + H2O2.

Acids act readily upon it, forming a sodium salt and hydrogen peroxide:

2NaO + 2HCl = 2NaCl + H2O2.

In these last two reactions the hydrogen dioxide formed may decompose into water and oxygen if the temperature is allowed to rise:

H2O2= H2O + O.

Peroxides.It will be remembered that barium dioxide (BaO_{2}) yields hydrogen dioxide when treated with acids, and that manganese dioxide gives up oxygen when heated with sulphuric acid. Oxides which yield either hydrogen dioxide or oxygen when treated with water or an acid are called peroxides.

Sodium hydroxide(caustic soda) (NaOH). 1.Preparation.Sodium hydroxide is prepared commercially by several processes.

(a) In the older process, still in extensive use, sodium carbonate is treated with calcium hydroxide suspended in water. Calcium carbonate is precipitated according to the equation

Na2CO3+ Ca(OH)2= CaCO3+ 2NaOH.

The dilute solution of sodium hydroxide, filtered from the calcium carbonate, is evaporated to a paste and is then poured into molds to solidify. It is sold in the form of slender sticks.

(b) The newer methods depend upon the electrolysis of sodium chloride. In the Castner process a solution of salt is electrolyzed, the reaction being expressed as follows:

NaCl + H2O = NaOH + H + Cl.

The chlorine escapes as a gas, and by an ingenious mechanical device the sodium hydroxide is prevented from mixing with the salt in the solution.

In the Acker process the electrolyte isfusedsodium chloride. The chlorine is evolved as a gas at the anode, while the sodium alloys with the melted lead which forms the cathode. When this alloy is treated with water the following reaction takes place:

Na + H2O = NaOH + H.

Fig. 77Fig. 77

Technical process.A sketch of an Acker furnace is represented in Fig. 77. The furnace is an irregularly shaped cast-iron box, divided into three compartments,A,B, andC. CompartmentAis lined with magnesia brick. CompartmentsBandCare filled with melted lead, which also covers the bottom ofAto a depth of about an inch. Above this layer inAis fused salt, into which dip carbon anodesD. The metallic box and melted lead is the cathode.

When the furnace is in operation chlorine is evolved at the anodes, and is drawn away through a pipe (not represented) to the bleaching-powder chambers. Sodium is set free at the surface of the melted lead inA, and at once alloys with it. Through the pipeEa powerful jet of steam is driven through the lead inBupwardsinto the narrow tubeF. This forces the lead alloy up through the tube and over into the chamberG.In this process the steam is decomposed by the sodium in the alloy, forming melted sodium hydroxide and hydrogen. The melted lead and sodium hydroxide separate into two layers inG, and the sodium hydroxide, being on top, overflows into tanks from which it is drawn off and packed in metallic drums. The lead is returned to the other compartments of the furnace by a pipe leading fromHtoI. CompartmentCserves merely as a reservoir for excess of melted lead.

In this process the steam is decomposed by the sodium in the alloy, forming melted sodium hydroxide and hydrogen. The melted lead and sodium hydroxide separate into two layers inG, and the sodium hydroxide, being on top, overflows into tanks from which it is drawn off and packed in metallic drums. The lead is returned to the other compartments of the furnace by a pipe leading fromHtoI. CompartmentCserves merely as a reservoir for excess of melted lead.

2.Properties.Sodium hydroxide is a white, crystalline, brittle substance which rapidly absorbs water and carbon dioxide from the air. As the name (caustic soda) indicates, it is a very corrosive substance, having a disintegrating action on most animal and vegetable tissues. It is a strong base. It is used in a great many chemical industries, and under the name of lye is employed to a small extent as a cleansing agent for household purposes.

Sodium chloride(common salt) (NaCl). 1.Preparation.Sodium chloride, or common salt, is very widely distributed in nature. Thick strata, evidently deposited at one time by the evaporation of salt water, are found in many places. In the United States the most important localities for salt are New York, Michigan, Ohio, and Kansas. Sometimes the salt is mined, especially if it is in the pure form called rock salt. More frequently a strong brine is pumped from deep wells sunk into the salt deposit, and is then evaporated in large pans until the salt crystallizes out. The crystals are in the form of small cubes and contain no water of crystallization; some water is, however, held in cavities in the crystals and causes the salt to decrepitate when heated.

2.Uses.Since salt is so abundant in nature it forms the starting point in the preparation of all compoundscontaining either sodium or chlorine. This includes many substances of the highest importance to civilization, such as soap, glass, hydrochloric acid, soda, and bleaching powder. Enormous quantities of salt are therefore produced each year. Small quantities are essential to the life of man and animals. Pure salt does not absorb moisture; the fact that ordinary salt becomes moist in air is not due to a property of the salt, but to impurities commonly occurring in it, especially calcium and magnesium chlorides.

Sodium sulphate(Glauber's salt) (Na2SO4·10H2O). This salt is prepared by the action of sulphuric acid upon sodium chloride, hydrochloric acid being formed at the same time:

2NaCl + H2SO4= Na2SO4+ 2HCl.

Some sodium sulphate is prepared by the reaction represented in the equation

MgSO4+ 2NaCl = Na2SO4+ MgCl2.

The magnesium sulphate required for this reaction is obtained in large quantities in the manufacture of potassium chloride, and being of little value for any other purpose is used in this way. The reaction depends upon the fact that sodium sulphate is the least soluble of any of the four factors in the equation, and therefore crystallizes out when hot, saturated solutions of magnesium sulphate and sodium chloride are mixed together and the resulting mixture cooled.

Sodium sulphate forms large efflorescent crystals. The salt is extensively used in the manufacture of sodium carbonate and glass. Small quantities are used in medicine.

Sodium sulphite(Na2SO3·7H2O). Sodium sulphite is prepared by the action of sulphur dioxide upon solutionsof sodium hydroxide, the reaction being analogous to the action of carbon dioxide upon sodium hydroxide. Like the carbonate, the sulphite is readily decomposed by acids:

Na2SO3+ 2HCl = 2NaCl + H2O + SO2.

Because of this reaction sodium sulphite is used as a convenient source of sulphur dioxide. It is also used as a disinfectant and a preservative.

Sodium thiosulphate(hyposulphite of soda or "hypo") (Na2S2O3·5H2O). This salt, commonly called sodium hyposulphite, or merely hypo, is made by boiling a solution of sodium sulphite with sulphur:

Na2SO3+ S = Na2S2O3.

It is used in photography and in the bleaching industry, to absorb the excess of chlorine which is left upon the bleached fabrics.

Thio compounds.The prefix "thio" means sulphur. It is used to designate substances which may be regarded as derived from oxygen compounds by replacing the whole or a part of their oxygen with sulphur. The thiosulphates may be regarded as sulphates in which one atom of oxygen has been replaced by an atom of sulphur. This may be seen by comparing the formula Na2SO4(sodium sulphate) with the formula Na2S2O3(sodium thiosulphate).

Sodium carbonate(sal soda)(Na2CO3·10H2O). There are two different methods now employed in the manufacture of this important substance.

1.Le Blanc process.This older process involves several distinct reactions, as shown in the following equations.

(a) Sodium chloride is first converted into sodium sulphate:

2NaCl + H2SO4= Na2SO4+ 2HCl.

(b) The sodium sulphate is next reduced to sulphide by heating it with carbon:

Na2SO4+ 2C = Na2S + 2CO2.

Na2S + CaCO3= CaS + Na2CO3.

Technical preparation of sodium carbonate.In a manufacturing plant the last two reactions take place in one process. Sodium sulphate, coal, and powdered limestone are heated together to a rather high temperature. The coal reduces the sulphate to sulphide, which in turn reacts upon the calcium carbonate. Some limestone is decomposed by the heat, forming calcium oxide. When treated with water the calcium oxide is changed into hydroxide, and this prevents the water from decomposing the insoluble calcium sulphide.The crude product of the process is a hard black cake called black ash. On digesting this mass with water the sodium carbonate passes into solution. The pure carbonate is obtained by evaporation of this solution, crystallizing from it in crystals of the formula Na2CO3·10H2O. Since over 60% of this salt is water, the crystals are sometimes heated until it is driven off. The product is called calcined soda, and is, of course, more valuable than the crystallized salt.

The crude product of the process is a hard black cake called black ash. On digesting this mass with water the sodium carbonate passes into solution. The pure carbonate is obtained by evaporation of this solution, crystallizing from it in crystals of the formula Na2CO3·10H2O. Since over 60% of this salt is water, the crystals are sometimes heated until it is driven off. The product is called calcined soda, and is, of course, more valuable than the crystallized salt.

2.Solvay process.This more modern process depends upon the reactions represented in the equations

NaCl + NH4HCO3= NaHCO3+ NH4Cl,2NaHCO3= Na2CO3+ H2O + CO2.

NaCl + NH4HCO3= NaHCO3+ NH4Cl,

2NaHCO3= Na2CO3+ H2O + CO2.

The reason the first reaction takes place is that sodium hydrogen carbonate is sparingly soluble in water, while the other compounds are freely soluble. When strong solutions of sodium chloride and of ammonium hydrogen carbonate are brought together the sparingly soluble sodium hydrogen carbonate is precipitated. This is converted into the normal carbonate by heating, the reaction being represented in the second equation.

Technical preparation.In the Solvay process a very concentrated solution of salt is first saturated with ammonia gas, and a current of carbon dioxide is then conducted into the solution. In this way ammonium hydrogen carbonate is formed:NH3+ H2O + CO2= NH4HCO3.This enters into double decomposition with the salt, as shown in the first equation under the Solvay process. After the sodium hydrogen carbonate has been precipitated the mother liquors containing ammonium chloride are treated with lime:2NH4Cl + CaO = CaCl2+ 2 NH3+ H2O.The lime is obtained by burning limestone:CaCO3= CaO + CO2.The ammonia and carbon dioxide evolved in the latter two reactions are used in the preparation of an additional quantity of ammonium hydrogen carbonate. It will thus be seen that there is no loss of ammonia. The only materials permanently used up are calcium carbonate and salt, while the only waste product is calcium chloride.Historical.In former times sodium carbonate was made by burning seaweeds and extracting the carbonate from their ash. On this account the salt was calledsoda ash, and the name is still in common use. During the French Revolution this supply was cut off, and in behalf of the French government Le Blanc made a study of methods of preparing the carbonate directly from salt. As a result he devised the method which bears his name, and which was used exclusively for many years. It has been replaced to a large extent by the Solvay process, which has the advantage that the materials used are inexpensive, and that the ammonium hydrogen carbonate used can be regenerated from the products formed in the process. Much expense is also saved in fuel, and the sodium hydrogen carbonate, which is the first product of the process, has itself many commercial uses. The Le Blanc process is still used, however, since the hydrochloric acid generated is of value.By-products.The substances obtained in a given process, aside from the main product, are called the by-products. The success of many processes depends upon the value of the by-products formed.Thus hydrochloric acid, a by-product in the Le Blanc process, is valuable enough to make the process pay, even though sodium carbonate can be made cheaper in other ways.

NH3+ H2O + CO2= NH4HCO3.

This enters into double decomposition with the salt, as shown in the first equation under the Solvay process. After the sodium hydrogen carbonate has been precipitated the mother liquors containing ammonium chloride are treated with lime:

2NH4Cl + CaO = CaCl2+ 2 NH3+ H2O.

The lime is obtained by burning limestone:

CaCO3= CaO + CO2.

The ammonia and carbon dioxide evolved in the latter two reactions are used in the preparation of an additional quantity of ammonium hydrogen carbonate. It will thus be seen that there is no loss of ammonia. The only materials permanently used up are calcium carbonate and salt, while the only waste product is calcium chloride.

Historical.In former times sodium carbonate was made by burning seaweeds and extracting the carbonate from their ash. On this account the salt was calledsoda ash, and the name is still in common use. During the French Revolution this supply was cut off, and in behalf of the French government Le Blanc made a study of methods of preparing the carbonate directly from salt. As a result he devised the method which bears his name, and which was used exclusively for many years. It has been replaced to a large extent by the Solvay process, which has the advantage that the materials used are inexpensive, and that the ammonium hydrogen carbonate used can be regenerated from the products formed in the process. Much expense is also saved in fuel, and the sodium hydrogen carbonate, which is the first product of the process, has itself many commercial uses. The Le Blanc process is still used, however, since the hydrochloric acid generated is of value.

By-products.The substances obtained in a given process, aside from the main product, are called the by-products. The success of many processes depends upon the value of the by-products formed.

Thus hydrochloric acid, a by-product in the Le Blanc process, is valuable enough to make the process pay, even though sodium carbonate can be made cheaper in other ways.

Properties of sodium carbonate.Sodium carbonate forms large crystals of the formula Na2CO3· 10 H2O. It has a mild alkaline reaction and is used for laundry purposes under the name of washing soda. Mere mention of the fact that it is used in the manufacture of glass, soap, and many chemical reagents will indicate its importance in the industries. It is one of the few soluble carbonates.

Sodium hydrogen carbonate(bicarbonate of soda) (NaHCO3). This salt, commonly called bicarbonate of soda, or baking soda, is made by the Solvay process, as explained above, or by passing carbon dioxide into strong solutions of sodium carbonate:

Na2CO3+ H2O + CO2= 2NaHCO3.

The bicarbonate, being sparingly soluble, crystallizes out. A mixture of the bicarbonate with some substance (the compound known as cream of tartar is generally used) which slowly reacts with it, liberating carbon dioxide, is used largely in baking. The carbon dioxide generated forces its way through the dough, thus making it porous and light.

Sodium nitrate(Chili saltpeter) (NaNO3). This substance is found in nature in arid regions in a number of places, where it has been formed apparently by the decay of organic substances in the presence of air and sodium salts. The largest deposits are in Chili, and most of the nitrate of commerce comes from that country. Smaller deposits occur in California and Nevada. The commercial salt is prepared by dissolving the crude nitrate in water,allowing the insoluble earthy materials to settle, and evaporating the clear solution so obtained to crystallization. The soluble impurities remain for the most part in the mother liquors.

Since this salt is the only nitrate found extensively in nature, it is the material from which other nitrates as well as nitric acid are prepared. It is used in enormous quantities in the manufacture of sulphuric acid and potassium nitrate, and as a fertilizer.

Sodium phosphate(Na2HPO4·12H2O). Since phosphoric acid has three replaceable hydrogen atoms, three sodium phosphates are possible,—two acid salts and one normal. All three can be made without difficulty, but disodium phosphate is the only one which is largely used, and is the salt which is commonly called sodium phosphate. It is made by the action of phosphoric acid on sodium carbonate:

Na2CO3+ H3PO4= Na2HPO4+ CO2+ H2O.

It is interesting as being one of the few phosphates which are soluble in water, and is the salt commonly used when a soluble phosphate is needed.

Normal sodium phosphate(Na3PO4). Although this is a normal salt its solution has a strongly alkaline reaction. This is due to the fact that the salt hydrolyzes in solution into sodium hydroxide and disodium phosphate, as represented in the equation

Na3PO4+ H2O = Na2HPO4+ NaOH.

Sodium hydroxide is strongly alkaline, while disodium phosphate is nearly neutral in reaction. The solution as a whole is therefore alkaline. The salt is prepared by adding a large excess of sodium hydroxide to a solution of disodiumphosphate and evaporating to crystallization. The excess of the sodium hydroxide reverses the reaction of hydrolysis and the normal salt crystallizes out.

Sodium tetraborate(borax) (Na2B4O7·10H2O). The properties of this important compound have been discussed under the head of boron.

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 ofthese 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:

SylvineKCl.AnhydriteCaSO4.CarnalliteKCl·MgCl2·6H2O.KainiteK2SO4·MgSO4·MgCl2·6H2O.PolyhaliteK2SO4·MgSO4·2CaSO4·2H2O.KieseriteMgSO4·H2O.SchöniteK2SO4·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. 78Fig. 78

Potassium hydroxide(caustic potash) (KOH). Potassium hydroxide is prepared by methods exactly similar to thoseused 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 withpotassium 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 aretreated 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.

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 andslow 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 equation2KNO3+ 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.

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.

Of the three remaining elements of the family—lithium, rubidium, and cæsium—lithium is by far the most common, the other two being very rare. Lithium chloride and carbonate are not infrequently found in natural mineral waters, and as these substances are supposed to increase the medicinal value of the water, they are very often added to artificial mineral waters in small quantities.

General.As explained in a previous chapter, when ammonia is passed into water the two compounds combine to form the base NH4OH, known as ammonium hydroxide. When this base is neutralized with acids there are formed the corresponding salts, known as the ammonium salts. Since the ammonium group is univalent, ammonium salts resemble those of the alkali metals in formulas; they also resemble the latter salts very much in their chemical properties, and may be conveniently described in connection with them. Among the ammonium salts the chloride, sulphate, carbonate, and sulphide are the most familiar.

Ammonium chloride(sal ammoniac) (NH4Cl). This substance is obtained by neutralizing ammonium hydroxide with hydrochloric acid. It is a colorless substance crystallizing in fine needles, and, like most ammonium salts, is very soluble in water. When placed in a tube and heated strongly it decomposes into hydrochloric acid and ammonia. When these gases reach a cooler portion of the tube theyat once recombine, and the resulting ammonium chloride is deposited on the sides of the tube. In this way the salt can be separated from nonvolatile impurities. Ammonium chloride is sometimes used in preparation of ammonia; it is also used in making dry batteries and in the laboratory as a chemical reagent.

Ammonium sulphate((NH4)2SO4). This salt resembles the chloride very closely, and, being cheaper, is used in place of it when possible. It is used in large quantity as a fertilizer, the nitrogen which it contains being a very valuable food for plants.

Ammonium carbonate((NH4)2CO3). This salt, as well as the acid carbonate (NH4HCO3), is used as a chemical reagent. They are colorless solids, freely soluble in water. The normal carbonate is made by heating ammonium chloride with powdered limestone (calcium carbonate), the ammonium carbonate being obtained as a sublimate in compact hard masses:

2NH4Cl + CaCO3= (NH4)2CO3+ CaCl2.

The salt always smells of ammonia, since it slowly decomposes, as shown in the equation

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