An Elementary Study of Chemistry, William McPherson [icecream ebook reader txt] 📗
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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:
(b) The sodium sulphate is next reduced to sulphide by heating it with carbon:
(c) The sodium sulphide is then heated with calcium carbonate, when double decomposition takes place:
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.
2. Solvay process. This more modern process depends upon the reactions represented in the equations
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 called soda 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:
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:
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
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 disodium phosphate 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.
POTASSIUMOccurrence 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:
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.
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:
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:
Equation (1), after being multiplied by 3, may be combined with equation (2), giving the following:
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
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