ZINC

Magnesium cement.A paste of magnesium hydroxide and water slowly absorbs carbon dioxide from the air and becomes very hard. The hardness of the product is increased by the presence of a considerable amount of magnesium chloride in the paste. The hydroxide, with or without the chloride, is used in the preparation of cements for some purposes.

Magnesium carbonate(MgCO3). Magnesium carbonate is a very abundant mineral. It occurs in a number of localities as magnesite, which is usually amorphous, but sometimes forms pure crystals resembling calcite. More commonly it is found associated with calcium carbonate.The mineral dolomite has the composition CaCO3·MgCO3. Limestone containing smaller amounts of magnesium carbonate is known as dolomitic limestone. Dolomite is one of the most common rocks, forming whole mountain masses. It is harder and less readily attacked by acids than limestone. It is valuable as a building stone and as ballast for roadbeds and foundations. Like calcium carbonate, magnesium carbonate is insoluble in water, though easily dissolved by acids.

Basic carbonate of magnesium.We should expect to find magnesium carbonate precipitated when a soluble magnesium salt and a soluble carbonate are brought together:

Na2CO3+ MgCl2= MgCO3+ 2NaCl.

Instead of this, some carbon dioxide escapes and the product is found to be a basic carbonate. The most common basic carbonate of magnesium has the formula 4MgCO3·Mg(OH)2, and is sometimes called magnesia alba. This compound is formed by the partial hydrolysis of the normal carbonate at first precipitated:

5MgCO3+ 2H2O = 4MgCO3·Mg(OH)2+ H2CO3.

Magnesium chloride(MgCl2·6H2O). Magnesium chloride is found in many natural waters and in many salt deposits (see Stassfurt salts). It is obtained as a by-product in the manufacture of potassium chloride from carnallite. As there is no very important use for it, large quantities annually go to waste. When heated to drive off the water of crystallization the chloride is decomposed as shown in the equation

MgCl2·6H2O = MgO + 2HCl + 5H2O.

Owing to the abundance of magnesium chloride, this reaction is being used to some extent in the preparation of both magnesium oxide and hydrochloric acid.

Boiler scale.When water which contains certain salts in solution is evaporated in steam boilers, a hard insoluble material calledscaledeposits in the boiler. The formation of this scale may be due to several distinct causes.1.To the deposit of calcium sulphate.This salt, while sparingly soluble in cold water, is almost completely insoluble in superheated water. Consequently it is precipitated when water containing it is heated in a boiler.2.To decomposition of acid carbonates.As we have seen, calcium and magnesium acid carbonates are decomposed on heating, forming insoluble normal carbonates:Ca(HCO3)2= CaCO3+ H2O + CO2.3.To hydrolysis of magnesium salts.Magnesium chloride, and to some extent magnesium sulphate, undergo hydrolysis when superheated in solution, and the magnesium hydroxide, being sparingly soluble, precipitates:MgCl2+ 2H2O <--> Mg(OH)2+ 2HCl.This scale adheres tightly to the boiler in compact layers and, being a non-conductor of heat, causes much waste of fuel. It is very difficult to remove, owing to its hardness and resistance to reagents. Thick scale sometimes cracks, and the water coming in contact with the overheated iron occasions an explosion. Moreover, the acids set free in the hydrolysis of the magnesium salts attack the iron tubes and rapidly corrode them. These causes combine to make the formation of scale a matter which occasions much trouble in cases where hard water is used in steam boilers. Water containing such salts should be softened, therefore, before being used in boilers.

1.To the deposit of calcium sulphate.This salt, while sparingly soluble in cold water, is almost completely insoluble in superheated water. Consequently it is precipitated when water containing it is heated in a boiler.

2.To decomposition of acid carbonates.As we have seen, calcium and magnesium acid carbonates are decomposed on heating, forming insoluble normal carbonates:

Ca(HCO3)2= CaCO3+ H2O + CO2.

3.To hydrolysis of magnesium salts.Magnesium chloride, and to some extent magnesium sulphate, undergo hydrolysis when superheated in solution, and the magnesium hydroxide, being sparingly soluble, precipitates:

MgCl2+ 2H2O <--> Mg(OH)2+ 2HCl.

This scale adheres tightly to the boiler in compact layers and, being a non-conductor of heat, causes much waste of fuel. It is very difficult to remove, owing to its hardness and resistance to reagents. Thick scale sometimes cracks, and the water coming in contact with the overheated iron occasions an explosion. Moreover, the acids set free in the hydrolysis of the magnesium salts attack the iron tubes and rapidly corrode them. These causes combine to make the formation of scale a matter which occasions much trouble in cases where hard water is used in steam boilers. Water containing such salts should be softened, therefore, before being used in boilers.

Magnesium sulphate(Epsom salt) (MgSO4·7H2O). Like the chloride, magnesium sulphate is found rather commonly in springs and in salt deposits. A very large deposit of the almost pure salt has been found in Wyoming. Its namewas given to it because of its abundant occurrence in the waters of the Epsom springs in England.

Magnesium sulphate has many uses in the industries. It is used to a small extent in the preparation of sodium and potassium sulphates, as a coating for cotton cloth, in the dye industry, in tanning, and in the manufacture of paints and laundry soaps. To some extent it is used in medicine.

Magnesium silicates.Many silicates containing magnesium are known and some of them are important substances. Serpentine, asbestos, talc, and meerschaum are examples of such substances.

Occurrence.Zinc never occurs free in nature. Its compounds have been found in many different countries, but it is not a constituent of common rocks and minerals, and its occurrence is rather local and confined to definite deposits or pockets. It occurs chiefly in the following ores:

Sphalerite (zinc blende)ZnS.ZinciteZnO.SmithsoniteZnCO3.WillemiteZn2SiO4.FrankliniteZnO·Fe2O3.

One fourth of the world's output of zinc comes from the United States, Missouri being the largest producer.

Metallurgy.The ores employed in the preparation of zinc are chiefly the sulphide, oxide, and carbonate. They are first roasted in the air, by which process they are changed into oxide:

ZnCO3= ZnO + CO2,ZnS + 3O = ZnO + SO2.

The oxide is then mixed with coal dust, and the mixture is heated in earthenware muffles or retorts, natural gas being used as fuel in many cases. The oxide is reduced by this means to the metallic state, and the zinc, being volatile at the high temperature reached, distills and is collected in suitable receivers. At first the zinc collects in the form of fine powder, called zinc dust or flowers of zinc, recalling the formation under similar conditions of flowers of sulphur. Later, when the whole apparatus has become warm, the zinc condenses to a liquid in the receiver, from which it is drawn off into molds. Commercial zinc often contains a number of impurities, especially carbon, arsenic, and iron.

Physical properties.Pure zinc is a rather heavy bluish-white metal with a high luster. It melts at about 420°, and if heated much above this temperature in the air takes fire and burns with a very bright bluish flame. It boils at about 950° and can therefore be purified by distillation.

Many of the physical properties of zinc are much influenced by the temperature and previous treatment of the metal. When cast into ingots from the liquid state it becomes at ordinary temperatures quite hard, brittle, and highly crystalline. At 150° it is malleable and can be rolled into thin sheets; at higher temperatures it again becomes very brittle. When once rolled into sheets it retains its softness and malleability at ordinary temperatures. When melted and poured into water it forms thin brittle flakes, and in this condition is called granulated or mossy zinc.

Chemical properties.Zinc is tarnished superficially by moist air, but beyond this is not affected by it. It does not decompose even boiling water. When the metal is quite pure, sulphuric and hydrochloric acids have scarcely any action upon it; when, however, it contains smallamounts of other metals such as magnesium or arsenic, or when it is merely in contact with metallic platinum, brisk action takes place and hydrogen is evolved. For this reason, when pure zinc is used in the preparation of hydrogen a few drops of platinum chloride are often added to the solution to assist the chemical action. Nitric acid dissolves the metal readily, with the formation of zinc nitrate and various reduction products of nitric acid. The strong alkalis act upon zinc and liberate hydrogen:

Zn + 2KOH = Zn(OK)2+ 2H.

The product of this reaction, potassium zincate, is a salt of zinc hydroxide, which is thus seen to have acid properties, though it usually acts as a base.

Uses of zinc.The metal has many familiar uses. Rolled into sheets, it is used as a lining for vessels which are to contain water. As a thin film upon the surface of iron (galvanized iron) it protects the iron from rust. Iron is usually galvanized by dipping it into a bath of melted zinc, but electrical methods are also employed. Zinc plates are used in many forms of electrical batteries. In the laboratory zinc is used in the preparation of hydrogen, and in the form of zinc dust as a reducing agent.

One of the largest uses of zinc is in the manufacture of alloys. Brass, an alloy of zinc and copper, is the most important of these; German silver, consisting of copper, zinc, and nickel, has many uses; various bronzes, coin metals, and bearing metals also contain zinc. Its ability to alloy with silver finds application in the separation of silver from lead (see silver).

Compounds of zinc.In general, the compounds of zinc are similar in formula and appearance to those of magnesium,but in other properties they often differ markedly. A number of them have value in commercial ways.

Zinc oxide(zinc white) (ZnO). Zinc oxide occurs in impure form in nature, being colored red by manganese and iron compounds. It can be prepared just like magnesium oxide, but is more often made by burning the metal.

Zinc oxide is a pure white powder which becomes yellow on heating and regains its white color when cold. It is much used as a white pigment in paints, under the name of zinc white, and has the advantage over white lead in that it is not changed in color by sulphur compounds, while lead turns black. It is also used in the manufacture of rubber goods.

Commercial preparation of zinc oxide.Commercially it is often made from franklinite in the following way. The franklinite is mixed with coal and heated to a high temperature in a furnace, by which process the zinc is set free and converted into vapor. As the vapor leaves the furnace through a conduit it meets a current of air and takes fire in it, forming zinc oxide. The oxide passes on and is filtered from the air through canvas bags, which allow the air to pass but retain the oxide. It is thus made by burning the metal, though the metal is not actually isolated in the process.

Soluble salts.The soluble salts of zinc can be made by dissolving the metal or the oxide in the appropriate acid. They are all somewhat poisonous. The sulphate and chloride are the most familiar.

Zinc sulphate(white vitriol) (ZnSO4·7H2O). This salt is readily crystallized from strong solutions in transparent colorless crystals. It is prepared commercially by careful roasting of the sulphide:

ZnS + 4O = ZnSO4.

Zinc chloride(ZnCl2·H2O). When a solution of zinc chloride is slowly evaporated a salt of the composition ZnCl2·H2O crystallizes out. If the water is completely expelled by heat and the residue distilled, the anhydrous chloride is obtained and may be cast into sticks or broken into lumps. In this distillation, just as in heating magnesium chloride, some of the chloride is decomposed:

ZnCl2·H2O = ZnO + 2HCl.

The anhydrous chloride has a great affinity for water, and is used as a dehydrating agent. It is also a germicide, and wood which is to be exposed to conditions which favor decay, as, for example, railroad ties, is often soaked in solutions of this salt.

Insoluble compounds.The insoluble compounds of zinc can be prepared by precipitation. The most important are the sulphide, carbonate, and hydroxide.

Zinc sulphide(ZnS). This substance occurs as the mineral sphalerite, and is one of the most valued ores of zinc. Very large deposits occur in southwestern Missouri. The natural mineral is found in large crystals or masses, resembling resin in color and luster. When prepared by precipitation the sulphide is white.

The element.This element occurs in small quantities in some zinc ores. In the course of the metallurgy of zinc the cadmium compounds undergo chemical changes quite similar to those of the zinc compounds, and the cadmium distills along with the zinc. Being more volatile, it comes over with the first of the zinc and is prepared from the first portions of the distillate by special methods of purification.The element very closely resembles zinc in most respects. Some of its alloys are characterized by having low melting points.

Compounds of cadmium.Among the compounds of cadmium may be mentioned the chloride (CdCl2·2H2O), the sulphate (3CdSO4·8H2O), and the nitrate (Cd(NO3)2·4H2O). These are white solids soluble in water. The sulphide (CdS) is a bright yellow substance which is insoluble in water and in dilute acids. It is valuable as a pigment in fine paints.

1.What properties have the metals of the magnesium family in common with the alkali metals; with the alkaline-earth metals?

2.Compare the action of the metals of the magnesium group on water with that of the other metals studied.

3.What metals already studied are prepared by electrolysis?

4.Write the equations representing the reactions between magnesium and hydrochloric acid; between magnesium and dilute sulphuric acid.

5.What property of magnesium was taken advantage of in the isolation of argon?

6.With phosphoric acid magnesium forms salts similar to those of calcium. Write the names and formulas of the corresponding magnesium salts.

7.How could you distinguish between magnesium chloride and magnesium sulphate? between Glauber's salts and Epsom salts?

8.What weight of carnallite is necessary in the preparation of 500 g. of magnesium?

9.Account for the fact that paints made of zinc oxide are not colored by hydrosulphuric acid.

10.What hydroxide studied, other than zinc hydroxide, has both acid and basic properties?

11.Write equations showing how the following compounds of zinc may be obtained from metallic zinc: the oxide, chloride, nitrate, carbonate, sulphate, sulphide, hydroxide.

The family.The element aluminium is the most abundant member of the group of elements known as the aluminium family; indeed, the other members of the family—gallium, indium, and thallium—are of such rare occurrence that they need not be separately described. The elements of the family are ordinarily trivalent, so that the formulas for their compounds differ from those of the elements so far studied. Their hydroxides are practically insoluble in water and are very weak bases; indeed, the bases are so weak that their salts are often hydrolyzed into free base and free acid in solution. The salts formed from these bases usually contain water of crystallization, which cannot be driven off without decomposing them more or less.

The trivalent metals, which in addition to aluminium include also iron and chromium, are sometimes called theearth metals. The name refers to the earthy appearance of the oxides of these metals, and to the fact that many earths, soils, and rocks are composed in part of these substances.

Occurrence.Aluminium never occurs in the free state in nature, owing to its great affinity for oxygen. In combined form, as oxides, silicates, and a few other salts, it is both abundant and widely distributed, being an essentialconstituent of all soils and of most rocks excepting limestone and sandstone. Cryolite (Na3AlF6), found in Greenland, and bauxite, which is an aluminium hydroxide usually mixed with some iron hydroxide, are important minerals. It is estimated that aluminium composes about 8% of the earth's crust. In the industries the metal is called aluminum, but its chemical name is aluminium.

Fig. 82Fig. 82

Preparation.Aluminium was first prepared by Wöhler, in 1827, by heating anhydrous aluminium chloride with potassium:

AlCl3+ 3K = 3KCl + Al.

This method was tried after it was found impossible to reduce the oxide of aluminium with carbon. The metal possessed such interesting properties and promised to be so useful that many efforts were made to devise a cheap way of preparing it. The method which has proved most successful consists in the electrolysis of the oxide dissolved in melted cryolite.

Metallurgy.An iron boxA(Fig. 82) about eight feet long and six feet wide is connected with a powerful generator in such a way as to serve as the cathode upon which the aluminium is deposited. Three or four rows of carbon rodsBdip into the box and serve as the anodes. The box is partially filled with cryolite and the current is turned on, generating enough heat to melt the cryolite. Aluminium oxide is then added, and under the influence of the electric current it decomposes into aluminium and oxygen. The temperature is maintained above the melting point of aluminium, and the liquid metal, being heavier than cryolite, sinks to the bottom of the vessel, from which it is tapped off from time to time through the tap holeC. The oxygen in part escapes as gas, and in part combines with the carbon of the anode, the combustion being very brilliant. The process is carried on at Niagara Falls.The largest expense in the process, apart from the cost of electrical energy, is the preparation of aluminium oxide free from other oxides, for most of the oxide found in nature is too impure to serve without refining. Bauxite is the principal ore used as a source of the aluminium because it is converted into pure oxide without great difficulty. Since common clay is a silicate of aluminium and is everywhere abundant, it might be expected that this would be utilized in the preparation of aluminium. It is, however, very difficult to extract the aluminium from a silicate, and no practical method has been found which will accomplish this.

The largest expense in the process, apart from the cost of electrical energy, is the preparation of aluminium oxide free from other oxides, for most of the oxide found in nature is too impure to serve without refining. Bauxite is the principal ore used as a source of the aluminium because it is converted into pure oxide without great difficulty. Since common clay is a silicate of aluminium and is everywhere abundant, it might be expected that this would be utilized in the preparation of aluminium. It is, however, very difficult to extract the aluminium from a silicate, and no practical method has been found which will accomplish this.

Physical properties.Aluminium is a tin-white metal which melts at 640° and is very light, having a density of 2.68. It is stiff and strong, and with frequent annealing can be rolled into thin foil. It is a good conductor of heat and electricity, though not so good as copper for a given cross section of wire.

Chemical properties.Aluminium is not perceptibly acted on by boiling water, and moist air merely dims its luster. Further action is prevented in each case by the formation of an extremely thin film of oxide upon the surface of the metal. It combines directly with chlorine, and when heated in oxygen burns with great energy and the liberation of much heat. It is therefore a good reducing agent. Hydrochloric acid acts upon it, forming aluminium chloride:nitric acid and dilute sulphuric acid have almost no action on it, but hot, concentrated sulphuric acid acts upon it in the same way as upon copper:

2Al + 6H2SO4= Al2(SO4)3+ 6H2O + 3SO2.

Alkalis readily attack the metal, liberating hydrogen, as in the case of zinc:

Al + 3KOH = Al(OK)3+ 3H.

Salt solutions, such as sea water, corrode the metal rapidly. It alloys readily with other metals.

Uses of aluminium.These properties suggest many uses for the metal. Its lightness, strength, and permanence make it well adapted for many construction purposes. These same properties have led to its extensive use in the manufacture of cooking utensils. The fact that it is easily corroded by salt solutions is, however, a disadvantage. Owing to its small resistance to electrical currents, it is replacing copper to some extent in electrical construction, especially for trolley and power wires. Some of its alloys have very valuable properties, and a considerable part of the aluminium manufactured is used for this purpose. Aluminium bronze, consisting of about 90% copper and 10% aluminium, has a pure golden color, is strong and malleable, is easily cast, and is permanent in the air. Considerable amounts of aluminium steel are also made.

Goldschmidt reduction process.Aluminium is frequently employed as a powerful reducing agent, many metallic oxides which resist reduction by carbon being readily reduced by it. The aluminium in the form of a fine powder is mixed with the metallic oxide, together with some substance such as fluorspar to act as a flux. The mixture is ignited, and the aluminium unites with theoxygen of the metallic oxide, liberating the metal. This collects in a fused condition under the flux.

An enormous quantity of heat is liberated in this reaction, and a temperature as high as 3500° can be reached. The heat of the reaction is turned to practical account in welding car rails, steel castings, and in similar operations where an intense local heat is required. A mixture of aluminium with various metallic oxides, ready prepared for such purposes, is sold under the name ofthermite.

Fig. 83Fig. 83

Preparation of chromium by the Goldschmidt method.A mixture of chromium oxide and aluminium powder is placed in a Hessian crucible (A, Fig. 83), and on top of it is placed a small heapBof a mixture of sodium peroxide and aluminium, into which is stuck a piece of magnesium ribbonC. Powdered fluorsparDis placed around the sodium peroxide, after which the crucible is set on a pan of sand and the magnesium ribbon ignited. When the flame reaches the sodium peroxide mixture combustion of the aluminium begins with almost explosive violence, so that great care must be taken in the experiment. The heat of this combustion starts the reaction in the chromium oxide mixture, and the oxide is reduced to metallic chromium. When the crucible has cooled a button of chromium will be found in the bottom.

Aluminium oxide(Al2O3). This substance occurs in several forms in nature. The relatively pure crystals are called corundum, while emery is a variety colored dark gray or black, usually with iron compounds. In transparent crystals, tinted different colors by traces of impurities, it forms such precious stones as the sapphire, oriental ruby, topaz, and amethyst. All these varieties are veryhard, falling little short of the diamond in this respect. Chemically pure aluminium oxide can be made by igniting the hydroxide, when it forms an amorphous white powder:

2Al(OH)3= Al2O3+ 3H2O.

The natural varieties, corundum and emery, are used for cutting and grinding purposes; the purest forms, together with the artificially prepared oxide, are largely used in the preparation of aluminium.

Aluminium hydroxide(Al(OH)3). The hydroxide occurs in nature as the mineral hydrargyllite, and in a partially dehydrated form called bauxite. It can be prepared by adding ammonium hydroxide to any soluble aluminium salt, forming a semi-transparent precipitate which is insoluble in water but very hard to filter. It dissolves in most acids to form soluble salts, and in the strong bases to form aluminates, as indicated in the equations

Al(OH)3+ 3HCl = AlCl3+ 3H2O,Al(OH)3+ 3NaOH = Al(ONa)3+ 3H2O.

It may act, therefore, either as a weak base or as a weak acid, its action depending upon the character of the substances with which it is in contact. When heated gently the hydroxide loses part of its hydrogen and oxygen according to the equation

Al(OH)3= AlO·OH + H2O.

This substance, the formula of which is frequently written HAlO2, is a more pronounced acid than is the hydroxide, and its salts are frequently formed when aluminium compounds are fused with alkalis. The magnesium salt Mg(AlO2)2is called spinel, and many other of its salts, called aluminates, are found in nature.

When heated strongly the hydroxide is changed into oxide, which will not again take up water on being moistened.

Mordants and dyeing.Aluminium hydroxide has the peculiar property of combining with many soluble coloring materials and forming insoluble products with them. On this account it is often used as a filter to remove objectionable colors from water. This property also leads to its wide use in the dye industry. Many dyes will not adhere to natural fibers such as cotton and wool, that is, will not "dye fast." If, however, the cloth to be dyed is soaked in a solution of aluminium compounds and then treated with ammonia, the aluminium salts which have soaked into the fiber will be converted into the hydroxide, which, being insoluble, remains in the body of it. If the fiber is now dipped into a solution of the dye, the aluminium hydroxide combines with the color material and fastens, or "fixes," it upon the fiber. A substance which serves this purpose is called amordant, and aluminium salts, particularly the acetate, are used in this way.

Aluminium chloride(AlCl3·6 H2O). This substance is prepared by dissolving the hydroxide in hydrochloric acid and evaporating to crystallization. When heated it is converted into the oxide, resembling magnesium in this respect:

2(AlCl3·6 H2O) = Al2O3+ 6HCl + 9H2O.

The anhydrous chloride, which has some important uses, is made by heating aluminium turnings in a current of chlorine.

Alums.Aluminium sulphate can be prepared by the action of sulphuric acid upon aluminium hydroxide. It has the property of combining with the sulphates of the alkali metals to form compounds calledalums. Thus, with potassium sulphate the reaction is expressed by the equation

K2SO4+ Al2(SO4)3+ 24H2O = 2(KAl(SO4)2·12H2O).

Under similar conditions ammonium sulphate yields ammonium alum:

(NH4)2SO4+ Al2(SO4)3+ 24H2O = 2(NH4Al(SO4)2·12H2O).

Other trivalent sulphates besides aluminium sulphate can form similar compounds with the alkali sulphates, and these compounds are also called alums, though they contain no aluminium. They all crystallize in octahedra and contain twelve molecules of water of crystallization. The alums most frequently prepared are the following:

Potassium alumKAl(SO4)2·12H2O.Ammonium alumNH4Al(SO4)2·12H2O.Ammonium iron alumNH4Fe(SO4)2·12H2O.Potassium chrome alumKCr(SO4)2·12H2O.

An alum may therefore be regarded as a compound derived from two molecules of sulphuric acid, in which one hydrogen atom has been displaced by the univalent alkali atom, and the other three hydrogen atoms by an atom of one of the trivalent metals, such as aluminium, iron, or chromium.

Very large, well-formed crystals of an alum can be prepared by suspending a small crystal by a thread in a saturated solution of the alum, as shown in Fig. 84. The small crystal slowly grows and assumes a very perfect form.

Fig. 84Fig. 84

Other salts of aluminium.While aluminium hydroxide forms fairly stable salts with strong acids, it is such a weak base that its salts with weak acids are readily hydrolyzed. Thus, when an aluminium salt and a soluble carbonate arebrought together in solution we should expect to have aluminium carbonate precipitated according to the equation

3Na2CO3+ 2AlCl3= Al2(CO3)3+ 6NaCl.

But if it is formed at all, it instantly begins to hydrolyze, the products of the hydrolysis being aluminium hydroxide and carbonic acid,

Al2(CO3)3+ 6H2O = 2Al(OH)3+ 3H2CO3.

Similarly a soluble sulphide, instead of precipitating aluminium sulphide (Al2S3), precipitates aluminium hydroxide; for hydrogen sulphide is such a weak acid that the aluminium sulphide at first formed hydrolyzes at once, forming aluminium hydroxide and hydrogen sulphide:

3Na2S + 2AlCl3+ 6H2O = 2Al(OH)3+ 6NaCl + 3H2S.

Alum baking powders.It is because of the hydrolysis of aluminium carbonate that alum is used as a constituent of some baking powders. The alum baking powders consist of a mixture of alum and sodium hydrogen carbonate. When water is added the two compounds react together, forming aluminium carbonate, which hydrolyzes into aluminium hydroxide and carbonic acid. The carbon dioxide from the latter escapes through the dough and in so doing raises it into a porous condition, which is the end sought in the use of a baking powder.

Aluminium silicates.One of the most common constituents of rocks is feldspar (KAlSi3O8), a mixed salt of potassium and aluminium with the polysilicic acid (H4Si3O8). Under the influence of moisture, carbon dioxide, and changes of temperature this substance is constantly being broken down into soluble potassium compounds and hydrated aluminium silicate. This compound has the formula Al2Si2O7·2H2O. In relatively pure condition it is called kaolin; in the impure state, mixed with sand and othersubstances, it forms common clay. Mica is another very abundant mineral, having varying composition, but being essentially of the formula KAlSiO4. Serpentine, talc, asbestos, and meerschaum are important complex silicates of aluminium and magnesium, and granite is a mechanical mixture of quartz, feldspar, and mica.

Ceramic industries.Many articles of greatest practical importance, ranging from the roughest brick and tile to the finest porcelain and chinaware, are made from some form of kaolin, or clay. No very precise classification of such ware can be made, as the products vary greatly in properties, depending upon the materials used and the treatment during manufacture.Porcelain is made from the purest kaolin, to which must be added some less pure, plastic kaolin, since the pure substance is not sufficiently plastic. There is also added some more fusible substance, such as feldspar, gypsum, or lime, together with some pure quartz. The constituents must be ground very fine, and when thoroughly mixed and moistened must make a plastic mass which can be molded into any desired form. The article molded from such materials is then burned. In this process the article is slowly heated to a point at which it begins to soften and almost fuse, and then it is allowed to cool slowly. At this stage, a very thin vessel will be translucent and have an almost glassy fracture; if, however, it is somewhat thicker, or has not been heated quite so high, it will still be porous, and partly on this account and partly to improve its appearance it is usually glazed.Glazing is accomplished by spreading upon the object a thin layer of a more fusible mixture of the same materials as compose the body of the object itself, and again heating until the glaze melts to a transparent glassy coating upon the surface of the vessel. In some cases fusible mixtures of quite different composition from that used in fashioning the vessel may be used as a glaze. Oxides of lead, zinc, and barium are often used in this way.When less carefully selected materials are used, or quite thick vessels are made, various grades of stoneware are produced. The inferior grades are glazed by throwing a quantity of common salt into the kiln towards the end of the first firing. In the form of vaporthe salt attacks the surface of the baked ware and forms an easily fusible sodium silicate upon it, which constitutes a glaze.Vitrified bricks, made from clay or ground shale, are burned until the materials begin to fuse superficially, forming their own glaze. Other forms of brick and tile are not glazed at all, but are left porous. The red color of ordinary brick and earthenware is due to an oxide of iron formed in the burning process.The decorations upon china are sometimes painted upon the baked ware and then glazed over, and sometimes painted upon the glaze and burned in by a third firing. Care must be taken to use such pigments as are not affected by a high heat and do not react chemically with the constituents of the baked ware or the glaze.

Porcelain is made from the purest kaolin, to which must be added some less pure, plastic kaolin, since the pure substance is not sufficiently plastic. There is also added some more fusible substance, such as feldspar, gypsum, or lime, together with some pure quartz. The constituents must be ground very fine, and when thoroughly mixed and moistened must make a plastic mass which can be molded into any desired form. The article molded from such materials is then burned. In this process the article is slowly heated to a point at which it begins to soften and almost fuse, and then it is allowed to cool slowly. At this stage, a very thin vessel will be translucent and have an almost glassy fracture; if, however, it is somewhat thicker, or has not been heated quite so high, it will still be porous, and partly on this account and partly to improve its appearance it is usually glazed.

Glazing is accomplished by spreading upon the object a thin layer of a more fusible mixture of the same materials as compose the body of the object itself, and again heating until the glaze melts to a transparent glassy coating upon the surface of the vessel. In some cases fusible mixtures of quite different composition from that used in fashioning the vessel may be used as a glaze. Oxides of lead, zinc, and barium are often used in this way.

When less carefully selected materials are used, or quite thick vessels are made, various grades of stoneware are produced. The inferior grades are glazed by throwing a quantity of common salt into the kiln towards the end of the first firing. In the form of vaporthe salt attacks the surface of the baked ware and forms an easily fusible sodium silicate upon it, which constitutes a glaze.

Vitrified bricks, made from clay or ground shale, are burned until the materials begin to fuse superficially, forming their own glaze. Other forms of brick and tile are not glazed at all, but are left porous. The red color of ordinary brick and earthenware is due to an oxide of iron formed in the burning process.

The decorations upon china are sometimes painted upon the baked ware and then glazed over, and sometimes painted upon the glaze and burned in by a third firing. Care must be taken to use such pigments as are not affected by a high heat and do not react chemically with the constituents of the baked ware or the glaze.

1.What metals and compounds studied are prepared by electrolysis?

2.Write the equation for the reaction between aluminium and hydrochloric acid; between aluminium and sulphuric acid (in two steps).

3.What hydroxides other than aluminium hydroxide have both acid and basic properties?

4.Write equations showing the methods used for preparing aluminium hydroxide and sulphate.

5.Write the general formula of an alum, representing an atom of an alkali metal by X and an atom of a trivalent metal by Y.

6.What is meant by the term polysilicic acid, as used in the discussion of aluminium silicates?

7.Compare the properties of the hydroxides of the different groups of metals so far studied.

8.In what respects does aluminium oxide differ from calcium oxide in properties?

9.Supposing bauxite to be 90% aluminium hydroxide, what weight of it is necessary for the preparation of 100 kg. of aluminium?

SYMBOLATOMIC WEIGHTDENSITYAPPROXIMATE MELTING POINTOXIDESIronFe55.97.931800°FeO, Fe2O3CobaltCo59.08.551800°CoO, Co2O3NickelNi58.78.91600°NiO, Ni2O3

The family.The elements iron, cobalt, and nickel form a group in the eighth column of the periodic table. The atomic weights of the three are very close together, and there is not the same gradual gradation in the properties of the three elements that is noticed in the families in which the atomic weights differ considerably in magnitude. The elements are very similar in properties, the similarity being so great in the case of nickel and cobalt that it is difficult to separate them by chemical analysis.

The elements occur in nature chiefly as oxides and sulphides, though they have been found in very small quantities in the native state, usually in meteorites. Their sulphides, carbonates, and phosphates are insoluble in water, the other common salts being soluble. Their salts are usually highly colored, those of iron being yellow or light green as a rule, those of nickel darker green, while cobalt salts are usually rose colored. The metals are obtained by reducing the oxides with carbon.

Occurrence.The element iron has long been known, since its ores are very abundant and it is not difficult to prepare the metal from them in fairly pure condition. It occurs in nature in many forms of combination,—in large deposits as oxides, sulphides, and carbonates, and in smaller quantities in a great variety of minerals. Indeed, very few rocks or soils are free from small amounts of iron, and it is assimilated by plants and animals playing an important part in life processes.

Metallurgy.It will be convenient to treat of the metallurgy of iron under two heads,—Materials Used and Process.

Materials used.Four distinct materials are used in the metallurgy of iron:

1.Iron ore.The ores most frequently used in the metallurgy of iron are the following:HematiteFe2O3.MagnetiteFe3O4.SideriteFeCO3.Limonite2Fe2O2·3H2O.These ores always contain impurities, such as silica, sulphides, and earthy materials. All ores, with the exception of the oxides, are first roasted to expel any water and carbon dioxide present and to convert any sulphide into oxide.2.Carbon.Carbon in some form is necessary both as a fuel and as a reducing agent. In former times wood charcoal was used to supply the carbon, but now anthracite coal or coke is almost universally used.3.Hot air.To maintain the high temperature required for the reduction of iron a very active combustion of fuelis necessary. This is secured by forcing a strong blast of hot air into the lower part of the furnace during the reduction process.4.Flux.(a)Purpose of the flux.All the materials which enter the furnace must leave it again either in the form of gases or as liquids. The iron is drawn off as the liquid metal after its reduction. To secure the removal of the earthy matter charged into the furnace along with the ore, materials are added to the charge which will, at the high temperature of the furnace, combine with the impurities in the ore, forming a liquid. The material added for this purpose is called theflux; the liquid produced from the flux and the ore is calledslag.(b)Function of the slag.While the main purpose of adding flux to the charge is to remove from the furnace in the form of liquid slag the impurities originally present in the ore, the slag thus produced serves several other functions. It keeps the contents of the furnace in a state of fusion, thus preventing clogging, and makes it possible for the small globules of iron to run together with greater ease into one large liquid mass.(c)Character of the slag.The slag is really a kind of readily fusible glass, being essentially a calcium-aluminium silicate. The ore usually contains silica and some aluminium compounds, so that limestone (which also contains some silica and aluminium) is added to furnish the calcium required for the slag. If the ore and the limestone do not contain a sufficient amount of silica and aluminium for the formation of the slag, these ingredients are added in the form of sand and feldspar. In the formation of slag from these materials the ore is freed from the silica and aluminium which it contained.

1.Iron ore.The ores most frequently used in the metallurgy of iron are the following:

HematiteFe2O3.MagnetiteFe3O4.SideriteFeCO3.Limonite2Fe2O2·3H2O.

These ores always contain impurities, such as silica, sulphides, and earthy materials. All ores, with the exception of the oxides, are first roasted to expel any water and carbon dioxide present and to convert any sulphide into oxide.

2.Carbon.Carbon in some form is necessary both as a fuel and as a reducing agent. In former times wood charcoal was used to supply the carbon, but now anthracite coal or coke is almost universally used.

3.Hot air.To maintain the high temperature required for the reduction of iron a very active combustion of fuelis necessary. This is secured by forcing a strong blast of hot air into the lower part of the furnace during the reduction process.

4.Flux.(a)Purpose of the flux.All the materials which enter the furnace must leave it again either in the form of gases or as liquids. The iron is drawn off as the liquid metal after its reduction. To secure the removal of the earthy matter charged into the furnace along with the ore, materials are added to the charge which will, at the high temperature of the furnace, combine with the impurities in the ore, forming a liquid. The material added for this purpose is called theflux; the liquid produced from the flux and the ore is calledslag.

(b)Function of the slag.While the main purpose of adding flux to the charge is to remove from the furnace in the form of liquid slag the impurities originally present in the ore, the slag thus produced serves several other functions. It keeps the contents of the furnace in a state of fusion, thus preventing clogging, and makes it possible for the small globules of iron to run together with greater ease into one large liquid mass.

(c)Character of the slag.The slag is really a kind of readily fusible glass, being essentially a calcium-aluminium silicate. The ore usually contains silica and some aluminium compounds, so that limestone (which also contains some silica and aluminium) is added to furnish the calcium required for the slag. If the ore and the limestone do not contain a sufficient amount of silica and aluminium for the formation of the slag, these ingredients are added in the form of sand and feldspar. In the formation of slag from these materials the ore is freed from the silica and aluminium which it contained.

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