 |
ZnCO_{3} = ZnO + CO_{2}, ZnS + 3O = ZnO + SO_{2}.
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 deg., 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 deg. 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 deg. 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 small amounts 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) (ZnSO{4}.7H{2}O). This salt is readily crystallized from strong solutions in transparent colorless crystals. It is prepared commercially by careful roasting of the sulphide:
ZnS + 4O = ZnSO_{4}.
Zinc chloride (ZnCl{2}.H{2}O). When a solution of zinc chloride is slowly evaporated a salt of the composition ZnCl{2}.H{2}O 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:
ZnCl{2}.H{2}O = 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.
CADMIUM
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 (CdCl_{2}.2H_{2}O), the sulphate (3CdSO_{4}.8H_{2}O), and the nitrate (Cd(NO_{3})_{2}.4H_{2}O). 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.
EXERCISES
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.
CHAPTER XXVI
THE ALUMINIUM FAMILY
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 the earth 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.
ALUMINIUM
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 essential constituent of all soils and of most rocks excepting limestone and sandstone. Cryolite (Na{3}AlF{6}), 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.
Preparation. Aluminium was first prepared by Woehler, in 1827, by heating anhydrous aluminium chloride with potassium:
AlCl_{3} + 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 box A (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 rods B dip 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 hole C. 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.
Physical properties. Aluminium is a tin-white metal which melts at 640 deg. 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 + 6H_{2}SO_{4} = Al_{2}(SO_{4})_{3} + 6H_{2}O + 3SO_{2}.
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 the oxygen 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 deg. 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 of thermite.
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 heap B of a mixture of sodium peroxide and aluminium, into which is stuck a piece of magnesium ribbon C. Powdered fluorspar D is 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 (Al{2}O{3}). 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 very hard, 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} = Al{2}O{3} + 3H{2}O.
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 = AlCl{3} + 3H{2}O, Al(OH){3} + 3NaOH = Al(ONa){3} + 3H{2}O.
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 + H{2}O.
This substance, the formula of which is frequently written HAlO_{2}, 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(AlO_{2})_{2} is 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 a mordant, and aluminium salts, particularly the acetate, are used in this way.
Aluminium chloride (AlCl{3}.6 H{2}O). 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(AlCl_{3}.6 H_{2}O) = Al_{2}O_{3} + 6HCl + 9H_{2}O.
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 called alums. Thus, with potassium sulphate the reaction is expressed by the equation
K_{2}SO_{4} + Al_{2}(SO_{4})_{3} + 24H_{2}O = 2(KAl(SO_{4})_{2}.12H_{2}O).
Under similar conditions ammonium sulphate yields ammonium alum:
(NH_{4})_{2}SO_{4} + Al_{2}(SO_{4})_{3} + 24H_{2}O = 2(NH_{4}Al(SO_{4})_{2}.12H_{2}O).
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 alum KAl(SO{4}){2}.12H{2}O. Ammonium alum NH{4}Al(SO{4}){2}.12H{2}O. Ammonium iron alum NH{4}Fe(SO{4}){2}.12H{2}O. Potassium chrome alum KCr(SO{4}){2}.12H{2}O.
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.
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 are brought together in solution we should expect to have aluminium carbonate precipitated according to the equation
3Na{2}CO{3} + 2AlCl{3} = Al{2}(CO{3}){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,
Al_{2}(CO_{3})_{3} + 6H_{2}O = 2Al(OH)_{3} + 3H_{2}CO_{3}.
Similarly a soluble sulphide, instead of precipitating aluminium sulphide (Al{2}S{3}), 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:
3Na_{2}S + 2AlCl_{3} + 6H_{2}O = 2Al(OH)_{3} + 6NaCl + 3H_{2}S.
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 (KAlSi{3}O{8}), a mixed salt of potassium and aluminium with the polysilicic acid (H{4}Si{3}O{8}). 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 Al{2}Si{2}O{7}.2H{2}O. In relatively pure condition it is called kaolin; in the impure state, mixed with sand and other substances, it forms common clay. Mica is another very abundant mineral, having varying composition, but being essentially of the formula KAlSiO{4}. 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 vapor the 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.
EXERCISES
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?
CHAPTER XXVII
THE IRON FAMILY
=================================================================== APPROXIMATE SYMBOL ATOMIC DENSITY MELTING OXIDES WEIGHT POINT Iron Fe 55.9 7.93 1800 deg. FeO, Fe{2}O{3} Cobalt Co 59.0 8.55 1800 deg. CoO, Co{2}O{3} Nickel Ni 58.7 8.9 1600 deg. NiO, Ni{2}O{3} ===================================================================
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.
IRON
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:
Hematite Fe{2}O{3}. Magnetite Fe{3}O{4}. Siderite FeCO{3}. Limonite 2Fe{2}O{2}.3H{2}O.
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 fuel is 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 the flux; the liquid produced from the flux and the ore is called slag.
(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.
Process. The reduction of iron is carried out in large towers called blast furnaces. The blast furnace (Fig. 85) is usually about 80 ft. high and 20 ft. in internal diameter at its widest part, narrowing somewhat both toward the top and toward the bottom. The walls are built of steel and lined with fire-brick. The base is provided with a number of pipes T, called tuyers, through which hot air can be forced into the furnace. The tuyers are supplied from a large pipe S, which circles the furnace as a girdle. The base has also an opening M, through which the liquid metal can be drawn off from time to time, and a second opening P, somewhat above the first, through which the excess of slag overflows. The top is closed by a movable trap C and C, called the cone, and through this the materials to be used are introduced. The gases produced by the combustion of the fuel and the reduction of the ore, together with the nitrogen of the air forced in through the tuyers, escape through pipes D, called downcomer pipes, which leave the furnace near the top. These gases are very hot and contain combustible substances, principally carbon monoxide; they are therefore utilized as fuel for the engines and also to heat the blast admitted through the tuyers. The lower part of the furnace is often furnished with a water jacket. This consists of a series of pipes W built into the walls, through which water can be circulated to reduce their temperature.
Charges consisting of coke (or anthracite coal), ore, and flux in proper proportions are introduced into the furnace at intervals through the trap top. The coke burns fiercely in the hot-air blast, giving an intense heat and forming carbon monoxide. The ore, working down in the furnace as the coke burns, becomes very hot, and by the combined reducing action of the carbon and carbon monoxide is finally reduced to metal and collects as a liquid in the bottom of the furnace, the slag floating on the molten iron. After a considerable amount of the iron has collected the slag is drawn off through the opening P. The molten iron is then drawn off into large ladles and taken to the converters for the manufacture of steel, or it is run out into sand molds, forming the bars or ingots called "pigs." The process is a continuous one, and when once started it is kept in operation for months or even years without interruption.
It seems probable that the first product of combustion of the carbon, at the point where the tuyers enter the furnace, is carbon dioxide. This is at once reduced to carbon monoxide by the intensely heated carbon present, so that no carbon dioxide can be found at that point. For practical purposes, therefore, we may consider that carbon monoxide is the first product of combustion.
Varieties of iron. The iron of commerce is never pure, but contains varying amounts of other elements, such as carbon, silicon, phosphorus, sulphur, and manganese. These elements may either be alloyed with the iron or may be combined with it in the form of definite chemical compounds. In some instances, as in the case of graphite, the mixture may be merely mechanical.
The properties of iron are very much modified by the presence of these elements and by the form of the combination between them and the iron; the way in which the metal is treated during its preparation has also a marked influence on its properties. Owing to these facts many kinds of iron are recognized in commerce, the chief varieties being cast iron, wrought iron, and steel.
Cast iron. The product of the blast furnace, prepared as just described, is called cast iron. It varies considerably in composition, usually containing from 90 to 95% iron, the remainder being largely carbon and silicon with smaller amounts of phosphorus and sulphur. When the melted metal from the blast furnace is allowed to cool rapidly most of the carbon remains in chemical combination with the iron, and the product is called white cast iron. If the cooling goes on slowly, the carbon partially separates as flakes of graphite which remain scattered through the metal. This product is softer and darker in color and is called gray cast iron.
Properties of cast iron. Cast iron is hard, brittle, and rather easily melted (melting point about 1100 deg.). It cannot be welded or forged into shape, but is easily cast in sand molds. It is strong and rigid but not elastic. It is used for making castings and in the manufacture of other kinds of iron. Cast iron, which contains the metal manganese up to the extent of 20%, together with about 3% carbon, is called spiegel iron; when more than this amount of manganese is present the product is called ferromanganese. The ferromanganese may contain as much as 80% manganese. These varieties of cast iron are much used in the manufacture of steel.
Wrought iron. Wrought iron is made by burning out from cast iron most of the carbon, silicon, phosphorus, and sulphur which it contains. The process is called puddling, and is carried out in a furnace constructed as represented in Fig. 86. The floor of the furnace F is somewhat concave and is made of iron covered with a layer of iron oxide. A long flame produced by burning fuel upon the grate G is directed downward upon the materials placed upon the floor, and the draught is maintained by the stack S. A is the ash box and T a trap to catch the solid particles carried into the stack by the draught. Upon the floor of the furnace is placed the charge of cast iron, together with a small amount of material to make a slag. The iron is soon melted by the flame directed upon it, and the sulphur, phosphorus, and silicon are oxidized by the iron oxide, forming oxides which are anhydrides of acids. These combine with the flux, which is basic in character, or with the iron oxide, to form a slag. The carbon is also oxidized and escapes as carbon dioxide. As the iron is freed from other elements it becomes pasty, owing to the higher melting point of the purer iron, and in this condition forms small lumps which are raked together into a larger one. The large lump is then removed from the furnace and rolled or hammered into bars, the slag; being squeezed out in this process. The product has a stranded or fibrous structure. The product of a puddling furnace is called wrought iron.
Properties of wrought iron. Wrought iron is nearly pure iron, usually containing about 0.3% of other substances, chiefly carbon. It is tough, malleable, and fibrous in structure. It is easily bent and is not elastic, so it will not sustain pressure as well as cast iron. It can be drawn out into wire of great tensile strength, and can also be rolled into thin sheets (sheet iron). It melts at a high temperature (about 1600 deg.) and is therefore forged into shape rather than cast. If melted, it would lose its fibrous structure and be changed into a low carbon steel.
Steel. Steel, like wrought iron, is made by burning out from cast iron a part of the carbon, silicon, phosphorus, and sulphur which it contains; but the process is carried out in a very different way, and usually, though not always, more carbon is found in steel than in wrought iron. A number of processes are in use, but nearly all the steel of commerce is made by one of the two following methods.
1. Bessemer process. This process, invented about 1860, is by far the most important. It is carried out in great egg-shaped crucibles called converters (Fig. 87), each one of which will hold as much as 15 tons of steel. The converter is built of steel and lined with silica. It is mounted on trunnions T, so that it can be tipped over on its side for filling and emptying. One of the trunnions is hollow and a pipe P connects it with an air chamber A, which forms a false bottom to the converter. The true bottom is perforated, so that air can be forced in by an air blast admitted through the trunnion and the air chamber.
White-hot, liquid cast iron from a blast furnace is run into the converter through its open necklike top O, the converter being tipped over to receive it; the air blast is then turned on and the converter rotated to a nearly vertical position. The elements in the iron are rapidly oxidized, the silicon first and then the carbon. The heat liberated in the oxidation, largely due to the combustion of silicon, keeps the iron in a molten condition. When the carbon is practically all burned out cast iron or spiegel iron, containing a known percentage of carbon, is added and allowed to mix thoroughly with the fluid. The steel is then run into molds, and the ingots so formed are hammered or rolled into rails or other forms. By this process any desired percentage of carbon can be added to the steel. Low carbon steel, which does not differ much from wrought iron in composition, is now made in this way and is replacing the more expensive wrought iron for many purposes.
The basic lining process. When the cast iron contains phosphorus and sulphur in appreciable quantities, the lining of the converter is made of dolomite. The silicon and carbon burn, followed by the phosphorus and sulphur, and the anhydrides of acids so formed combine with the basic oxides of the lining, forming a slag. This is known as the basic lining process.
2. Open-hearth process. In this process a furnace very similar to a puddling furnace is used, but it is lined with silica or dolomite instead of iron oxide. A charge consisting in part of old scrap iron of any kind and in part of cast iron is melted in the furnace by a gas flame. The silicon and carbon are slowly burned away, and when a test shows that the desired percentage of carbon is present the steel is run out of the furnace. Steel may therefore be defined as the product of the Bessemer or open-hearth processes.
Properties of steel. Bessemer and open-hearth steel usually contain only a few tenths of a per cent of carbon, less than 0.1% silicon, and a very much smaller quantity of phosphorus and sulphur. Any considerable amount of the latter elements makes the steel brittle, the sulphur affecting it when hot, and the phosphorus when cold. This kind of steel is used for structural purposes, for rails, and for nearly all large steel articles. It is hard, malleable, ductile, and melts at a lower temperature than wrought iron. It can be forged into shape, rolled into sheets, or cast in molds.
Relation of the three varieties of iron. It will be seen that wrought iron is usually very nearly pure iron, while steel contains an appreciable amount of alloy material, chiefly carbon, and cast iron still more of the same substances. It is impossible, however, to assign a given sample of iron to one of these three classes on the basis of its chemical composition alone. A low carbon steel, for example, may contain less carbon than a given sample of wrought iron. The real distinction between the three is the process by which they are made. The product of the blast furnace is cast iron; that of the puddling furnace is wrought iron; that of the Bessemer and open-hearth methods is steel.
Tool steel. Steel designed for use in the manufacture of edged tools and similar articles should be relatively free from silicon and phosphorus, but should contain from 0.5 to 1.5% carbon. The percentage of carbon should be regulated by the exact use to which the steel is to be put. Steel of this character is usually made in small lots from either Bessemer or open-hearth steel in the following way.
A charge of melted steel is placed in a large crucible and the calculated quantity of pure carbon is added. The carbon dissolves in the steel, and when the solution is complete the metal is poured out of the crucible. This is sometimes called crucible steel.
Tempering of steel. Steel containing from 0.5 to 1.5% carbon is characterized by the property of "taking temper." When the hot steel is suddenly cooled by plunging it into water or oil it becomes very hard and brittle. On carefully reheating this hard form it gradually becomes less brittle and softer, so that by regulating the temperature to which steel is reheated in tempering almost any condition of temper demanded for a given purpose, such as for making springs or cutting tools, can be obtained.
Steel alloys. It has been found that small quantities of a number of different elements when alloyed with steel very much improve its quality for certain purposes, each element having a somewhat different effect. Among the elements most used in this connection are manganese, silicon, chromium, nickel, tungsten, and molybdenum.
The usual method for adding these elements to the steel is to first prepare a very rich alloy of iron with the element to be added, and then add enough of this alloy to a large quantity of the steel to bring it to the desired composition. A rich alloy of iron with manganese or silicon can be prepared directly in a blast furnace, and is called ferromanganese or ferrosilicon. Similar alloys of iron with the other elements mentioned are made in an electric furnace by reducing the mixed oxides with carbon.
Pure iron. Perfectly pure iron is rarely prepared and is not adapted to commercial uses. It can be made by reducing pure oxide of iron in a current of hydrogen at a high temperature. Prepared in this way it forms a black powder; when melted it forms a tin-white metal which is less fusible and more malleable than wrought iron. It is easily acted upon by moist air.
Compounds of iron. Iron differs from the metals so far studied in that it is able to form two series of compounds in which the iron has two different valences. In the one series the iron is divalent and forms compounds which in formulas and many chemical properties are similar to the corresponding zinc compounds. It can also act as a trivalent metal, and in this condition forms salts similar to those of aluminium. Those compounds in which the iron is divalent are known as ferrous compounds, while those in which it is trivalent are known as ferric.
Oxides of iron. Iron forms several oxides. Ferrous oxide (FeO) is not found in nature, but can be prepared artificially in the form of a black powder which easily takes up oxygen, forming ferric oxide:
2FeO + O = Fe{2}O{3}.
Ferric oxide is the most abundant ore of iron and occurs in great deposits, especially in the Lake Superior region. It is found in many mineral varieties which vary in density and color, the most abundant being hematite, which ranges in color from red to nearly black. When prepared by chemical processes it forms a red powder which is used as a paint pigment (Venetian red) and as a polishing powder (rouge).
Magnetite has the formula Fe{3}O{4} and is a combination of FeO and Fe{2}O{3}. It is a very valuable ore, but is less abundant than hematite. It is sometimes called magnetic oxide of iron, or lodestone, since it is a natural magnet.
Ferrous salts. These salts are obtained by dissolving iron in the appropriate acid, or, when insoluble, by precipitation. They are usually light green in color and crystallize well. In chemical reactions they are quite similar to the salts of magnesium and zinc, but differ from them in one important respect, namely, that they are easily changed into compounds in which the metal is trivalent. Thus ferrous chloride treated with chlorine or aqua regia is changed into ferric chloride:
FeCl{2} + Cl = FeCl{3}.
Ferrous hydroxide exposed to moist air is rapidly changed into ferric hydroxide:
2Fe(OH)_{2} + H_{2}O + O = 2Fe(OH)_{3}.
Ferrous sulphate (copperas, green vitriol) (FeSO{4}.7H{2}O). Ferrous sulphate is the most familiar ferrous compound. It is prepared commercially as a by-product in the steel-plate mills. Steel plates are cleaned by the action of dilute sulphuric acid upon them, and in the process some of the iron dissolves. The liquors are concentrated and the green vitriol separates from them.
Ferrous sulphide (FeS). Ferrous sulphide is sometimes found in nature as a golden-yellow crystalline mineral. It is formed as a black precipitate when a soluble sulphide and an iron salt are brought together in solution:
FeSO{4} + Na{2}S = FeS + Na{2}SO{4}.
It can also be made as a heavy dark-brown solid by fusing together the requisite quantities of sulphur and iron. It is obtained as a by-product in the metallurgy of lead:
PbS + Fe = FeS + Pb.
It is used in the laboratory in the preparation of hydrosulphuric acid:
FeS + 2HCl = FeCl{2} + H{2}S.
Iron disulphide _(pyrites)_ (FeS_{2}). This substance bears the same relation to ferrous sulphide that hydrogen dioxide does to water. It occurs abundantly in nature in the form of brass-yellow cubical crystals and in compact masses. Sometimes the name "fool's gold" is applied to it from its superficial resemblance to the precious metal. It is used in very large quantities as a source of sulphur dioxide in the manufacture of sulphuric acid, since it burns readily in the air, forming ferric oxide and sulphur dioxide:
2FeS{2} + 11O = Fe{2}O{3} + 4SO{2}.
Ferrous carbonate (FeCO_{3}). This compound occurs in nature as siderite, and is a valuable ore. It will dissolve to some extent in water containing carbon dioxide, just as will calcium carbonate, and waters containing it are called chalybeate waters. These chalybeate waters are supposed to possess certain medicinal virtues and form an important class of mineral waters.
Ferric salts. Ferric salts are usually obtained by treating an acidified solution of a ferrous salt with an oxidizing agent:
2FeCl_{2} + 2HCl + O = 2FeCl_{3} + H_{2}O,
2FeSO_{4} + H_{2}SO_{4} + O = Fe_{2}(SO_{4})_{3} + H_{2}O.
They are usually yellow or violet in color, are quite soluble, and as a rule do not crystallize well. Heated with water in the absence of free acid, they hydrolyze even more readily than the salts of aluminium. The most familiar ferric salts are the chloride and the sulphate.
Ferric chloride (FeCl_{3}). This salt can be obtained most conveniently by dissolving iron in hydrochloric acid and then passing chlorine into the solution:
Fe + 2HCl = FeCl_{2} + 2H,
FeCl{2} + Cl = FeCl{3}.
When the pure salt is heated with water it is partly hydrolyzed:
FeCl_{3} + 3 H_{2}O Fe(OH)_{3} + 3HCl.
This is a reversible reaction, however, and hydrolysis can therefore be prevented by first adding a considerable amount of the soluble product of the reaction, namely, hydrochloric acid.
Ferric sulphate (Fe{2}(SO{4}){3}). This compound can be made by treating an acid solution of green vitriol with an oxidizing agent. It is difficult to crystallize and hard to obtain in pure condition. When an alkali sulphate in proper quantity is added to ferric sulphate in solution an iron alum is formed, and is easily obtained in large crystals. The best known iron alums have the formulas KFe(SO{4}){2}.12H{2}O and NH{4}Fe(SO{4}){2}.12H{2}O. They are commonly used when a pure ferric salt is required.
Ferric hydroxide (Fe(OH)_{3}). When solutions of ferric salts are treated with ammonium hydroxide, ferric hydroxide is formed as a rusty-red precipitate, insoluble in water.
Iron cyanides. A large number of complex cyanides containing iron are known, the most important being potassium ferrocyanide, or yellow prussiate of potash (K{4}FeC{6}N{6}), and potassium ferricyanide, or red prussiate of potash (K{3}FeC{6}N{6}). These compounds are the potassium salts of the complex acids of the formulas H{4}FeC{6}N{6} and H{3}FeC{6}N{6}.
Oxidation of ferrous salts. It has just been seen that when a ferrous salt is treated with an oxidizing agent in the presence of a free acid a ferric salt is formed:
2FeSO_{4} + H_{2}SO_{4} + O = Fe_{2}(SO_{4})_{3} + H_{2}O.
In this reaction oxygen is used up, and the valence of the iron is changed from 2 to 3. The same equation may be written
2Fe^{+}, 2SO{4}^{—} + 2H^{+}, SO{4}^{—} + O = 2Fe^{+}, 3SO{4}^{—} + H{2}O.
Hydrogen ions have been oxidized to water, while the charge of each iron ion has been increased from 2 to 3.
In a similar way the conversion of ferrous chloride into ferric chloride may be written
Fe^{}, 2Cl^{-} + Cl = Fe^{}, + 3Cl^{-}.
Here again the valence of the iron and the charge on the iron ion has been increased from 2 to 3, though no oxygen has entered into the reaction. As a rule, however, changes of this kind are brought about by the use of an oxidizing agent, and are called oxidations.
The term "oxidation" is applied to all reactions in which the valence of the metal of a compound is increased, or, in other words, to all reactions in which the charge of a cation is increased.
Reduction of ferric salts. The changes which take place when a ferric salt is converted into a ferrous salt are the reverse of the ones just described. This is seen in the equation
FeCl{3} + H = FeCl{2} + HCl
In this reaction the valence of the iron has been changed from 3 to 2. The same equation may be written
Fe^{+}, 3Cl{-} + H = Fe^{+}, + H^{+} + 3Cl{-}
It will be seen that the charge of the iron ions has been diminished from 3 to 2. Since these changes are the reverse of the oxidation changes just considered, they are called reduction reactions. The term "reduction" is applied to all processes in which the valence of the metal of a compound is diminished, or, in other words, to all processes in which the charge on the cations is diminished.
NICKEL AND COBALT
These elements occur sparingly in nature, usually combined with arsenic or with arsenic and sulphur. Both elements have been found in the free state in meteorites. Like iron they form two series of compounds, but the salts corresponding to the ferrous salts are the most common, the ones corresponding to the ferric salts being difficult to obtain. Thus we have the chlorides NiCl{2}.6H{2}O and CoCl{2}.6H{2}O; the sulphates NiSO{4}.7H{2}O and CoSO{4}.7H{2}O; the nitrates Ni(NO{3}){2}.6H{2}O and Co(NO{3}){2}.6H{2}O.
Nickel is largely used as an alloy with other metals. Alloyed with copper it forms coin metal from which five-cent pieces are made, with copper and zinc it forms German silver, and when added to steel in small quantities nickel steel is formed which is much superior to common steel for certain purposes. When deposited by electrolysis upon the surface of other metals such as iron, it forms a covering which will take a high polish and protects the metal from rust, nickel not being acted upon by moist air. Salts of nickel are usually green.
Compounds of cobalt fused with glass give it an intensely blue color. In powdered form such glass is sometimes used as a pigment called smalt. Cobalt salts, which contain water of crystallization, are usually cherry red in color; when dehydrated they become blue.
EXERCISES
1. In the manufacture of cast iron, why is the air heated before being forced into the furnace?
2. Write the equations showing how each of the following compounds of iron could be obtained from the metal itself: ferrous chloride, ferrous hydroxide, ferrous sulphate, ferrous sulphide, ferrous carbonate, ferric chloride, ferric sulphate, ferric hydroxide.
3. Account for the fact that a solution of sodium carbonate, when added to a solution of a ferric salt, precipitates an hydroxide and not a carbonate.
4. Calculate the percentage of iron in each of the common iron ores.
5. One ton of steel prepared by the Bessemer process is found by analysis to contain 0.2% carbon. What is the minimum weight of carbon which must be added in order that the steel may be made to take a temper?
CHAPTER XXVIII
COPPER, MERCURY, AND SILVER
================================================================== FORMULAS OF OXIDES SYMBOL ATOMIC DENSITY MELTING ____ WEIGHT POINT "ous" "ic" __ __ __ __ __ __ __ Copper Cu 63.6 8.89 1084 deg. Cu_{2}O CuO Mercury Hg 200.00 13.596 -39.5 deg. Hg_{2}O HgO Silver Ag 107.93 10.5 960 deg. Ag_{2}O AgO ==================================================================
The family. By referring to the periodic arrangement of the elements (page 168), it will be seen that mercury is not included in the same family with copper and silver. Since the metallurgy of the three elements is so similar, however, and since they resemble each other so closely in chemical properties, it is convenient to class them together for study.
1. Occurrence. The three elements occur in nature to some extent in the free state, but are usually found as sulphides. Their ores are easy to reduce.
2. Properties. They are heavy metals of high luster and are especially good conductors of heat and electricity. They are not very active chemically. Neither hydrochloric nor dilute sulphuric acid has any appreciable action upon them. Concentrated sulphuric acid attacks all three, forming metallic sulphates and evolving sulphur dioxide, while nitric acid, both dilute and concentrated, converts them into nitrates with the evolution of oxides of nitrogen.
3. _Two series of salts._ Copper and mercury form oxides of the types M_{2}O and MO, as well as two series of salts. In one series the metals are univalent and the salts have formulas like those of the sodium salts. They are called cuprous and mercurous salts. In the other series the metals are divalent and resemble magnesium salts in formulas. These are called cupric and mercuric salts. Silver forms only one series of salts, being always a univalent metal.
COPPER
Occurrence. The element copper has been used for various purposes since the earliest days of history. It is often found in the metallic state in nature, large masses of it occurring pure in the Lake Superior region and in other places to a smaller extent. The most valuable ores are the following:
Cuprite Cu_{2}O. Chalcocite Cu_{2}S. Chalcopyrite CuFeS_{2}. Bornite Cu_{3}FeS_{3}. Malachite CuCO_{3}.Cu(OH)_{2}. Azurite 2CuCO_{3}.Cu(OH)_{2}.
Metallurgy of copper. Ores containing little or no sulphur are easy to reduce. They are first crushed and the earthy impurities washed away. The concentrated ore is then mixed with carbon and heated in a furnace, metallic copper resulting from the reduction of the copper oxide by the hot carbon.
Metallurgy of sulphide ores. Much of the copper of commerce is made from chalcopyrite and bornite, and these ores are more difficult to work. They are first roasted in the air, by which treatment much of the sulphur is burned to sulphur dioxide. The roasted ore is then melted in a small blast furnace or in an open one like a puddling furnace. In melting, part of the iron combines with silica to form a slag of iron silicate. The product, called crude matte, contains about 50% copper together with sulphur and iron. Further purification is commonly carried on by a process very similar to the Bessemer process for steel. The converter is lined with silica, and a charge of matte from the melting furnace, together with sand, is introduced, and air is blown into the mass. By this means the sulphur is practically all burned out by the air, and the remaining iron combines with silica and goes off as slag. The copper is poured out of the converter and molded into anode plates for refining.
Refining of copper. Impure copper is purified by electrolysis. A large plate of it, serving as an anode, is suspended in a tank facing a thin plate of pure copper, which is the cathode. The tank is filled with a solution of copper sulphate and sulphuric acid to serve as the electrolyte. A current from a dynamo passes from the anode to the cathode, and the copper, dissolving from the anode, is deposited upon the cathode in pure form, while the impurities collect on the bottom of the tank. Electrolytic copper is one of the purest of commercial metals and is very nearly pure copper.
Recovery of gold and silver. Gold and silver are often present in small quantities in copper ores, and in electrolytic refining these metals collect in the muddy deposit on the bottom of the tank. The mud is carefully worked over from time to time and the precious metals extracted from it. A surprising amount of gold and silver is obtained in this way.
Properties of copper. Copper is a rather heavy metal of density 8.9, and has a characteristic reddish color. It is rather soft and is very malleable, ductile, and flexible, yet tough and strong; it melts at 1084 deg.. As a conductor of heat and electrical energy it is second only to silver.
Hydrochloric acid, dilute sulphuric acid, and fused alkalis are almost without action upon it; nitric acid and hot, concentrated sulphuric acid, however, readily dissolve it. In moist air it slowly becomes covered with a thin layer of green basic carbonate; heated in the air it is easily oxidized to black copper oxide (CuO).
Uses. Copper is extensively used for electrical purposes, for roofs and cornices, for sheathing the bottom of ships, and for making alloys. In the following table the composition of some of these alloys is indicated:
COMPOSITION OF ALLOYS OF COPPER IN PERCENTAGES
Aluminium bronze copper (90 to 97%), aluminium (3 to 10%). Brass copper (63 to 73%), zinc (27 to 37%). Bronze copper (70 to 95%), zinc (1 to 25%), tin (1 to 18%). German silver copper (56 to 60%), zinc (20%), nickel (20 to 25%). Gold coin copper (10%), gold (90%). Gun metal copper (90%), tin (10%). Nickel coin copper (75%), nickel (25%) Silver coin copper (10%), silver (90%).
Electrotyping. Matter is often printed from electrotype plates which are prepared as follows. The matter is set up in type and wax is firmly pressed down upon the face of it until a clear impression is obtained. The impressed side of the wax is coated with graphite and the impression is made the cathode in an electrolytic cell containing a copper salt in solution. When connected with a current the copper is deposited as a thin sheet upon the letters in wax, and when detached is a perfect copy of the type, the under part of the letters being hollow. The sheet is strengthened by pouring on the under surface a suitable amount of molten metal (commercial lead is used). The sheet so strengthened is then used in printing.
Two series of copper compounds. Copper, like iron, forms two series of compounds: in the cuprous compounds it is univalent; in the cupric it is divalent. The cupric salts are much the more common of the two, since the cuprous salts pass readily into cupric by oxidation.
Cuprous compounds. The most important cuprous compound is the oxide (Cu_{2}O), which occurs in nature as ruby copper or cuprite. It is a bright red substance and can easily be prepared by heating copper to a high temperature in a limited supply of air. It is used for imparting a ruby color to glass.
By treating cuprous oxide with different acids a number of cuprous salts can be made. Many of these are insoluble in water, the chloride (CuCl) being the best known. When suspended in dilute hydrochloric acid it is changed into cupric chloride, the oxygen taking part in the reaction being absorbed from the air:
2CuCl + 2HCl + O = 2CuCl{2} + H{2}O.
Cupric compounds. Cupric salts are easily made by dissolving cupric oxide in acids, or, when insoluble, by precipitation. Most of them are blue or green in color, and the soluble ones crystallize well. Since they are so much more familiar than the cuprous salts, they are frequently called merely copper salts.
Cupric oxide (CuO). This is a black insoluble substance obtained by heating copper in excess of air, or by igniting the hydroxide or nitrate. It is used as an oxidizing agent.
Cupric hydroxide (Cu(OH)_{2}). The hydroxide prepared by treating a solution of a copper salt with sodium hydroxide is a light blue insoluble substance which easily loses water and changes into the oxide. Heat applied to the liquid containing the hydroxide suspended in it serves to bring about the reaction represented by the equation
Cu(OH){2} = CuO + H{2}O.
Cupric sulphate (blue vitriol) (CuSO{4}.5H{2}O). This substance, called blue vitriol or bluestone, is obtained as a by-product in a number of processes and is produced in very large quantities. It forms large blue crystals, which lose water when heated and crumble to a white powder. The salt finds many uses, especially in electrotyping and in making electrical batteries.
Cupric sulphide (CuS). The insoluble black sulphide (CuS) is easily prepared by the action of hydrosulphuric acid upon a solution of a copper salt:
CuSO{4} + H{2}S = CuS + H{2}SO{4}.
It is insoluble in water and dilute acids.
MERCURY
Occurrence. Mercury occurs in nature chiefly as the sulphide (HgS) called cinnabar, and in globules of metal inclosed in the cinnabar. The mercury mines of Spain have long been famous, California being the next largest producer.
Metallurgy. Mercury is a volatile metal which has but little affinity for oxygen. Sulphur, on the other hand, readily combines with oxygen. These facts make the metallurgy of mercury very simple. The crushed ore, mixed with a small amount of carbon to reduce any oxide or sulphate that might be formed, is roasted in a current of air. The sulphur burns to sulphur dioxide, while the mercury is converted into vapor and is condensed in a series of condensing vessels. The metal is purified by distillation.
Properties. Mercury is a heavy silvery liquid with a density of 13.596. It boils at 357 deg. and solidifies at -39.5 deg.. Small quantities of many metals dissolve in it, forming liquid alloys, while with larger quantities it forms solid alloys. The alloys of mercury are called amalgams.
Toward acids mercury conducts itself very much like copper; it is easily attacked by nitric and hot, concentrated sulphuric acids, while cold sulphuric and hydrochloric acids have no effect on it.
Uses. Mercury is extensively used in the construction of scientific instruments, such as the thermometer and barometer, and as a liquid over which to collect gases which are soluble in water. The readiness with which it alloys with silver and gold makes it very useful in the extraction of these elements.
Compounds of mercury. Like copper, mercury forms two series of compounds: the mercurous, of which mercurous chloride (HgCl) is an example; and the mercuric, represented by mercuric chloride (HgCl_{2}).
Mercuric oxide (HgO). Mercuric oxide can be obtained either as a brick-red or as a yellow substance. When mercuric nitrate is heated carefully the red modification is formed in accordance with the equation
Hg(NO_{3})_{2} = HgO + 2NO_{2} + O.
The yellow modification is prepared by adding a solution of a mercuric salt to a solution of sodium or potassium hydroxide:
Hg(NO{3}){2} + 2NaOH = 2NaNO{3} + Hg(OH){2},
Hg(OH){2} = HgO + H{2}O.
When heated the oxide darkens until it becomes almost black; at a higher temperature it decomposes into mercury and oxygen. It was by this reaction that oxygen was discovered.
Mercurous chloride (calomel) (HgCl). Being insoluble, mercurous chloride is precipitated as a white solid when a soluble chloride is added to a solution of mercurous nitrate:
HgNO{3} + NaCl = HgCl + NaNO{3}.
Commercially it is manufactured by heating a mixture of mercuric chloride and mercury. When exposed to the light it slowly changes into mercuric chloride and mercury:
2HgCl = HgCl_{2} + Hg.
It is therefore protected from the light by the use of colored bottles. It is used in medicine.
Most mercurous salts are insoluble in water, the principal soluble one being the nitrate, which is made by the action of cold, dilute nitric acid on mercury.
Mercuric chloride (_corrosive sublimate_) (HgCl_{2}). This substance can be made by dissolving mercuric oxide in hydrochloric acid. On a commercial scale it is made by subliming a mixture of common salt and mercuric sulphate:
2NaCl + HgSO{4} = HgCl{2} + Na{2}SO{4}.
The mercuric chloride, being readily volatile, vaporizes and is condensed again in cool vessels. Like mercurous chloride it is a white solid, but differs from it in that it is soluble in water. It is extremely poisonous and in dilute solutions is used as an antiseptic in dressing wounds.
Mercuric sulphide (HgS). As cinnabar this substance forms the chief native compound of mercury, occurring in red crystalline masses. By passing hydrosulphuric acid into a solution of a mercuric salt it is precipitated as a black powder, insoluble in water and acids. By other means it can be prepared as a brilliant red powder known as vermilion, which is used as a pigment in fine paints.
The iodides of mercury. If a solution of potassium iodide is added to solutions of a mercurous and a mercuric salt respectively, the corresponding iodides are precipitated. Mercuric iodide is the more important of the two, and as prepared above is a red powder which changes to yellow on heating to 150 deg.. The yellow form on cooling changes back again to the red form, or may be made to do so by rubbing it with a knife blade or some other hard object.
SILVER
Occurrence. Silver is found in small quantities in the uncombined state; usually, however, it occurs in combination with sulphur, either as the sulphide (Ag_{2}S) or as a small constituent of other sulphides, especially those of lead and copper. It is also found alloyed with gold.
Metallurgy. Parkes's process. Silver is usually smelted in connection with lead. The ores are worked over together, as described under lead, and the lead and silver obtained as an alloy, the silver being present in small quantity. The alloy is melted and metallic zinc is stirred in. Zinc will alloy with silver but not with lead, and it is found that the silver leaves the lead and, in the form of an alloy with zinc, forms as a crust upon the lead and is skimmed off. This crust, which, of course, contains lead adhering to it, is partially melted and the most of the lead drained off. The zinc is removed by distillation, and the residue is melted on an open hearth in a current of air; by this means the zinc and lead remaining with the silver are changed into oxides and the silver remains behind unaltered.
Amalgamation process. In some localities the old amalgamation process is used. The silver ore is treated with common salt and ferrous compounds, which process converts the silver first into chloride and then into metallic silver. Mercury is then added and thoroughly mixed with the mass, forming an amalgam with the silver. After some days the earthy materials are washed away and the heavier amalgam is recovered. The mercury is distilled off and the silver left in impure form.
Refining silver. The silver obtained by either of the above processes may still contain copper, gold, and iron, and is refined by "parting" with sulphuric acid. The metal is heated with strong sulphuric acid which dissolves the silver, copper, and iron present, but not the gold. In the solution of silver sulphate so obtained copper plates are suspended, upon which the pure silver precipitates, the copper going into solution as sulphate, as shown in the equation
Ag_{2}SO_{4} + Cu = 2Ag + CuSO_{4}.
The solution obtained as a by-product in this process furnishes most of the blue vitriol of commerce. Silver is also refined by electrolytic methods similar to those used in refining copper.
Properties of silver. Silver is a heavy, rather soft, white metal, very ductile and malleable and capable of taking a high polish. It surpasses all other metals as a conductor of heat and electricity, but is too costly to find extensive use for such purposes. It melts at a little lower temperature than copper (961 deg.). It alloys readily with other heavy metals, and when it is to be used for coinage a small amount of copper—from 8 to 10%—is nearly always melted with it to give it hardness.
It is not acted upon by water or air, but is quickly tarnished when in contact with sulphur compounds, turning quite black in time. Hydrochloric acid and fused alkalis do not act upon it, but nitric acid and hot, concentrated sulphuric acid dissolve it with ease.
Electroplating. Since silver is not acted upon by water or air, and has a pleasing appearance, it is used to coat various articles made of cheaper metals. Such articles are said to be silver plated. The process by which this is done is called electroplating. It is carried on as follows: The object to be plated (such as a spoon) is attached to a wire and dipped into a solution of a silver salt. Electrical connection is made in such a way that the article to be plated serves as the cathode, while the anode is made up of one or more plates of silver (Fig. 88, A). When a current is passed through the electrolyte silver dissolves from the anode plate and deposits on the cathode in the form of a closely adhering layer. By making the proper change in the electrolyte and anode plate objects may be plated with gold and other metals.
Compounds of silver. Silver forms two oxides but only one series of salts, namely, the one which corresponds to the mercurous and cuprous series.
Silver nitrate (_lunar caustic_) (AgNO_{3}). This salt is easily prepared by dissolving silver in nitric acid and evaporating the resulting solution. It crystallizes in flat plates, and when heated carefully can be melted without decomposition. When cast into sticks it is called lunar caustic, for it has a very corrosive action on flesh, and is sometimes used in surgery to burn away abnormal growths.
The alchemists designated the metals by the names of the heavenly bodies. The moon (luna) was the symbol for silver; hence the name "lunar caustic."
Silver sulphide (Ag_{2}S). This occurs in nature and constitutes one of the principal ores of silver. It can be obtained in the form of a black solid by passing hydrosulphuric acid through a solution of silver nitrate.
Compounds of silver with the halogens. The chloride, bromide, and iodide of silver are insoluble in water and acids, and are therefore precipitated by bringing together a soluble halogen salt with silver nitrate:
AgNO{3} + KCl = AgCl + KNO{3}.
They are remarkable for the fact that they are very sensitive to the action of light, undergoing a change of color and chemical composition when exposed to sunlight, especially if in contact with organic matter such as gelatin.
Photography. The art of photography is based on the fact that the halogen compounds of silver are affected by the light, particularly in the presence of organic matter. From a chemical standpoint the processes involved may be described under two heads: (1) the preparation of the negative; (2) the preparation of the print.
1. Preparation of the negative. The plate used in the preparation of the negative is made by spreading a thin layer of gelatin, in which silver bromide is suspended (silver iodide is sometimes added also), over a glass plate or celluloid film and allowing it to dry. When the plate so prepared is placed in a camera and the image of some object is focused upon it, the silver salt undergoes a change which is proportional at each point to the intensity of the light falling upon it. In this way an image of the object photographed is produced upon the plate, which is, however, invisible and is therefore called "latent." It can be made visible by the process of developing.
To develop the image the exposed plate is immersed in a solution of some reducing agent called the developer. The developer reduces that portion of the silver salt which has been affected by the light, depositing it in the form of black metallic silver which closely adheres to the plate.
The unaffected silver salt, upon which the developer has no action, must now be removed from the plate. This is done by immersing the plate in a solution of sodium thiosulphate (hypo). After the silver salt has been dissolved off, the plate is washed with water and dried. The plate so prepared is called the negative because it is a picture of the object photographed, with the lights exactly reversed. This is called fixing the negative.
2. Preparation of the print. The print is made from paper which is prepared in the same way as the negative plate. The negative is placed upon this paper and exposed to the light in such a way that the light must pass through the negative before striking the paper. If the paper is coated with silver chloride, a visible image is produced, in which case a developer is not needed. The proofs are made in this way. In order to make them permanent the unchanged silver chloride must be dissolved off with sodium thiosulphate. The print is then toned by dipping it into a solution of gold or platinum salts. The silver on the print passes into solution, while the gold or platinum takes its place. These metals give a characteristic color or tone to the print, the gold making it reddish brown, while the platinum gives it a steel-gray tone. If a silver bromide paper is used in making the print, a latent image is produced which must be developed as in the case of the negative itself. The silver bromide is much more sensitive than the chloride, so that the printing can be done in artificial light. Since the darkest places on the negative cut off the most light, it is evident that the lights of the print will be the reverse of those of the negative, and will therefore correspond to those of the object photographed. The print is therefore called the positive.
EXERCISES
1. Account for the fact that copper has been used for so long a time.
2. Write equations for the action of concentrated sulphuric and nitric acids upon the metals of this family.
3. How would you account for the fact that normal copper sulphate is slightly acid to litmus?
4. Contrast the action of heat on cupric nitrate and mercuric nitrate.
5. State reasons why mercury is adapted for use in thermometers and barometers.
6. How could you distinguish between mercurous chloride and mercuric chloride?
7. Write equations for the preparation of mercuric and mercurous iodides.
8. How would you account for the fact that solutions of the different salts of a metal usually have the same color?
9. Crude silver usually contains iron and lead. What would become of these metals in refining by parting with sulphuric acid?
10. In the amalgamation process for extracting silver, how does ferrous chloride convert silver chloride into silver? Write equation. Why is the silver sulphide first changed into silver chloride?
11. What impurities would you expect to find in the copper sulphate prepared from the refining of silver?
12. How could you prepare pure silver chloride from a silver coin?
13. Mercuric nitrate and silver nitrate are both white solids soluble in water. How could you distinguish between them?
14. Account for the fact that sulphur waters turn a silver coin black; also for the fact that a silver spoon is blackened by foods (eggs, for example) containing sulphur.
15. When a solution of silver nitrate is added to a solution of potassium chlorate no precipitate forms. How do you account for the fact that a precipitate of silver chloride is not formed?
CHAPTER XXIX
TIN AND LEAD
==================================================================== SYMBOL ATOMIC DENSITY MELTING COMMON OXIDES WEIGHT POINT Tin Sn 119.0 7.35 235 deg. SnO SnO{2} Lead Pb 206.9 11.38 327 deg. PbO Pb{3}O{4} PbO{2} ====================================================================
The family. Tin and lead, together with silicon and germanium, form a family in Group IV of the periodic table. Silicon has been discussed along with the non-metals, while germanium, on account of its rarity, needs only to be mentioned.
The other family of Group IV includes carbon, already described, and a number of rare elements.
TIN
Occurrence. Tin is found in nature chiefly as the oxide (SnO_{2}), called cassiterite or tinstone. The most famous mines are those of Cornwall in England, and of the Malay Peninsula and East India Islands; in small amounts tinstone is found in many other localities.
Metallurgy. The metallurgy of tin is very simple. The ore, separated as far as possible from earthy materials, is mixed with carbon and heated in a furnace, the reduction taking place readily. The equation is
SnO{2} + C = Sn + CO{2}.
The metal is often purified by carefully heating it until it is partly melted; the pure tin melts first and can be drained away from the impurities.
Properties. Pure tin, called block tin, is a soft white metal with a silver-like appearance and luster; it melts readily (235 deg.) and is somewhat lighter than copper, having a density of 7.3. It is quite malleable and can be rolled out into very thin sheets, forming tin foil; most tin foil, however, contains a good deal of lead.
Under ordinary conditions it is quite unchanged by air or moisture, but at a high temperature it burns in air, forming the oxide SnO_{2}. Dilute acids have no effect upon it, but concentrated acids attack it readily. Concentrated hydrochloric acid changes it into the chloride
Sn + 2HCl = SnCl_{2} + 2H.
With sulphuric acid tin sulphate and sulphur dioxide are formed:
Sn + 2H_{2}SO_{4} = SnSO_{4} + SO_{2} + 2H_{2}O
Concentrated nitric acid oxidizes it, forming a white insoluble compound of the formula H{2}SnO{3}, called metastannic acid:
3Sn + 4HNO{3} + H{2}O = 3H{2}SnO{3} + 4NO.
Uses of tin. A great deal of tin is made into tin plate by dipping thin steel sheets into the melted metal. Owing to the way in which tin resists the action of air and dilute acids, tin plate is used in many ways, such as in roofing, and in the manufacture of tin cans, cooking vessels, and similar articles.
Many useful alloys contain tin, some of which have been mentioned in connection with copper. When tin is alloyed with other metals of low melting point, soft, easily melted alloys are formed which are used for friction bearings in machinery; tin, antimony, lead, and bismuth are the chief constituents of these alloys. Pewter and soft solder are alloys of tin and lead.
Compounds of tin. Tin forms two series of compounds: the stannous, in which the tin is divalent, illustrated in the compounds SnO, SnS, SnCl_{2}; the stannic, in which it is tetravalent as shown in the compounds SnO_{2}, SnS_{2}. There is also an acid, H_{2}SnO_{3}, called stannic acid, which forms a series of salts called stannates. While this acid has the same composition as metastannic acid, the two are quite different in their chemical properties. This difference is probably due to the different arrangement of the atoms in the molecules of the two substances. Only a few compounds of tin need be mentioned.
Stannic oxide (SnO_{2}). Stannic oxide is of interest, since it is the chief compound of tin found in nature. It is sometimes found in good-sized crystals, but as prepared in the laboratory is a white powder. When fused with potassium hydroxide it forms potassium stannate, acting very much like silicon dioxide:
SnO{2} + 2KOH = K{2}SnO{3} + H{2}O.
Chlorides of tin. Stannous chloride is prepared by dissolving tin in concentrated hydrochloric acid and evaporating the solution to crystallization. The crystals which are obtained have the composition SnCl{2}.2H{2}O, and are known as tin crystals. By treating a solution of stannous chloride with aqua regia, stannic chloride is formed:
SnCl{2} + 2Cl = SnCl{4}.
The salt which crystallizes from such a solution has the composition SnCl_{4}.5H_{2}O, and is known commercially as oxymuriate of tin. If metallic tin is heated in a current of dry chlorine, the anhydrous chloride (SnCl_{4}) is obtained as a heavy colorless liquid which fumes strongly on exposure to air.
The ease with which stannous chloride takes up chlorine to form stannic chloride makes it a good reducing agent in many reactions, changing the higher chlorides of metals to lower ones. Thus mercuric chloride is changed into mercurous chloride:
SnCl_{2} + 2HgCl_{2} = SnCl_{4} + 2HgCl.
If the stannous chloride is in excess, the reaction may go further, producing metallic mercury:
SnCl{2} + 2HgCl = SnCl{4} + 2Hg.
Ferric chloride is in like manner reduced to ferrous chloride:
SnCl{3} + 2FeCl{3} = SnCl{4} + 2FeCl{2}.
The chlorides of tin, as well as the alkali stannates, are much used as mordants in dyeing processes. The hydroxides of tin and free stannic acid, which are easily liberated from these compounds, possess in very marked degree the power of fixing dyes upon fibers, as explained under aluminium.
LEAD
Occurrence. Lead is found in nature chiefly as the sulphide (PbS), called galena; to a much smaller extent it occurs as carbonate, sulphate, chromate, and in a few other forms. Practically all the lead of commerce is made from galena, two general methods of metallurgy being in use.
Metallurgy. 1. The sulphide is melted with scrap iron, when iron sulphide and metallic lead are formed; the liquid lead, being the heavier, sinks to the bottom of the vessel and can be drawn off:
PbS + Fe = Pb + FeS.
2. The sulphide is roasted in the air until a part of it has been changed into oxide and sulphate. The air is then shut off and the heating continued, the reactions indicated in the following equations taking place:
2PbO + PbS = 3Pb + SO_{2},
PbSO{4} + PbS = 2Pb + 2SO{2}.
The lead so prepared usually contains small amounts of silver, arsenic, antimony, copper, and other metals. The silver is removed by Parkes's method, as described under silver, and the other metals in various ways. The lead of commerce is one of the purest commercial metals, containing as a rule only a few tenths per cent of impurities.
Properties. Lead is a heavy metal (den. = 11.33) which has a brilliant silvery luster on a freshly cut surface, but which soon tarnishes to a dull blue-gray color. It is soft, easily fused (melting at 327 deg.), and quite malleable, but has little toughness or strength.
It is not acted upon to any great extent by the oxygen of the air under ordinary conditions, but is changed into oxide at a high temperature. With the exception of hydrochloric and sulphuric acids, most acids, even very weak ones, act upon it, forming soluble lead salts. Hot, concentrated hydrochloric and sulphuric acids also attack it to a slight extent.
Uses. Lead is employed in the manufacture of lead pipes and in large storage batteries. In the form of sheet lead it is used in lining the chambers of sulphuric acid works and in the preparation of paint pigments. Some alloys of lead, such as solder and pewter (lead and tin), shot (lead and arsenic), and soft bearing metals, are widely used. Type metal consists of lead, antimony, and sometimes tin. Compounds of lead form several important pigments.
Compounds of lead. In nearly all its compounds lead has a valence of 2, but a few corresponding to stannic compounds have a valence of 4.
Lead oxides. Lead forms a number of oxides, the most important of which are litharge, red lead or minium, and lead peroxide.
1. Litharge (PbO). This oxide forms when lead is oxidized at a rather low temperature, and is obtained as a by-product in silver refining. It is a pale yellow powder, and has a number of commercial uses. It is easily soluble in nitric acid:
PbO + 2HNO{3} = Pb(NO{3}){2} + H{2}O.
2. Red lead, or minium (Pb{3}O{4}). Minium is prepared by heating lead (or litharge) to a high temperature in the air. It is a heavy powder of a beautiful red color, and is much used as a pigment.
3. _Lead peroxide_ (PbO_{2}). This is left as a residue when minium is heated with nitric acid:
Pb_{3}O_{4} + 4HNO_{3} = 2Pb(NO_{3})_{2} + PbO_{2} + 2H_{2}O.
It is a brown powder which easily gives up a part of its oxygen and, like manganese dioxide and barium dioxide, is a good oxidizing agent.
Soluble salts of lead. The soluble salts of lead can be made by dissolving (Pb(C_{2}H_{3}O_{2})_{2}.3H_{2}O), litharge in acids. Lead acetate called sugar of lead, and lead nitrate (Pb(NO_{3})_{2}) are the most familiar examples. They are while crystalline solids and are poisonous in character.
Insoluble salts of lead; lead carbonate. While the normal carbonate of lead (PbCO_{3}) is found to some extent, in nature and can be prepared in the laboratory, basic carbonates of varying composition are much more easy to obtain. One of the simplest of these has the composition 2PbCO_{3}.Pb(OH)_{2}. A mixture of such carbonates is called white lead. This is prepared on a large scale as a paint pigment and as a body for paints which are to be colored with other substances.
White lead. White lead is an amorphous white substance which, when mixed with oil, has great covering power, that is, it spreads out in an even waxy film, free from streaks and lumps, and covers the entire surface upon which it is spread. Its disadvantage as a pigment lies in the fact that it gradually blackens when exposed to sulphur compounds, which are often present in the air, forming black lead sulphide (PbS).
Technical preparation of white lead. Different methods are used in the preparation of white lead, but the old one known as the Dutch process is still the principal one employed. In this process, earthenware pots about ten inches high and of the shape shown in Fig. 89 are used. In the bottom A is placed a 3% solution of acetic acid (vinegar answers the purpose very well). The space above this is filled with thin, perforated, circular pieces of lead, supported by the flange B of the pot. These pots are placed close together on a bed of tan bark on the floor of a room known as the corroding room. They are covered over with boards, upon which tan bark is placed, and another row of pots is placed on this. In this way the room is filled. The white lead is formed by the fumes of the acetic acid, together with the carbon dioxide set free in the fermentation of the tan bark acting on the lead. About three months are required to complete the process.
Lead sulphide (PbS). In nature this compound occurs in highly crystalline condition, the crystals having much the same luster as pure lead. It is readily prepared in the laboratory as a black precipitate, by the action of hydrosulphuric acid upon soluble lead salts:
Pb(NO{3}){2} + H{2}S = PbS + 2HNO{3}.
It is insoluble both in water and in dilute acids.
Other insoluble salts. Lead chromate (PbCrO_{4}) is a yellow substance produced by the action of a soluble lead salt upon a soluble chromate, thus:
K{2}CrO{4} + Pb(NO{3}){2} = PbCrO{4} + 2 KNO{3}.
It is used as a yellow pigment. Lead sulphate (PbSO{4}) is a white substance sometimes found in nature and easily prepared by precipitation. Lead chloride (PbCl{2}) is likewise a white substance nearly insoluble in cold water, but readily soluble in boiling water.
Thorium and cerium. These elements are found in a few rare minerals, especially in the monazite sand of the Carolinas and Brazil. The oxides of these elements are used in the preparation of the Welsbach mantles for gas lights, because of the intense light given out when a mixture of the oxides is heated. These mantles contain the oxides of cerium and thorium in the ratio of about 1% of the former to 99% of the latter. Compounds of thorium, like those of radium, are found to possess radio-activity, but in a less degree.
EXERCISES
1. How could you detect lead if present in tin foil?
2. Stannous chloride reduces gold chloride (AuCl_{3}) to gold. Give equation.
3. What are the products of hydrolysis when stannic chloride is used as a mordant?
4. How could you detect arsenic, antimony, or copper in lead?
5. Why is lead so extensively used for making water pipes?
6. What sulphates other than lead are insoluble?
7. Could lead nitrate be used in place of barium chloride in testing for sulphates?
8. How much lead peroxide could be obtained from 1 kg. of minium?
9. The purity of white lead is usually determined by observing the volume of carbon dioxide given off when it is treated with an acid. What acid should be used? On the supposition that it has the formula 2PbCO{3}.Pb(OH){2}, how nearly pure was a sample if 1 g. gave 30 cc. of carbon dioxide at 20 deg. and 750 mm.?
10. Silicon belongs in the same family with tin and lead. In what respects are these elements similar?
11. What weight of tin could be obtained by the reduction of 1 ton of cassiterite?
12. What reaction would you expect to take place when lead peroxide is treated with hydrochloric acid?
13. White lead is often adulterated with barytes. Suggest a method for detecting it, if present, in a given example of white lead.
CHAPTER XXX
MANGANESE AND CHROMIUM
==================================================================== SYMBOL ATOMIC DENSITY MELTING FORMULAS OF ACIDS WEIGHT POINT Manganese Mn 55.0 8.01 1900 deg. H{2}MnO{4} and HMnO{4} Chromium Cr 52.1 7.3 3000 deg. H{2}CrO{4} and H{2}Cr{2}O{7} ====================================================================
General. Manganese and chromium, while belonging to different families, have so many features in common in their chemical conduct that they may be studied together with advantage. They differ from most of the elements so far studied in that they can act either as acid-forming or base-forming elements. As base-forming elements each of the metals forms two series of salts. In the one series, designated by the suffix "ous," the metal is divalent; in the other series, designated by the suffix "ic," the metal is trivalent. Only the manganous and the chromic salts, however, are of importance. The acids in which these elements play the part of a non-metal are unstable, but their salts are usually stable, and some of them are important compounds.
MANGANESE
Occurrence. Manganese is found in nature chiefly as the dioxide MnO{2}, called pyrolusite. In smaller amounts it occurs as the oxides Mn{2}O{3} and Mn{3}O{4}, and as the carbonate MnCO{3}. Some iron ores also contain manganese.
Preparation and properties. The element is difficult to prepare in pure condition and has no commercial applications. It can be prepared, however, by reducing the oxide with aluminium powder or by the use of the electric furnace, with carbon as the reducing agent. The metal somewhat resembles iron in appearance, but is harder, less fusible, and more readily acted upon by air and moisture. Acids readily dissolve it, forming manganous salts.
Oxides of manganese. The following oxides of manganese are known: MnO, Mn_{2}O_{3}, Mn_{3}O_{4}, MnO_{2}, and Mn_{2}O_{7}. Only one of these, the dioxide, needs special mention.
Manganese dioxide (_pyrolusite_) (MnO_{2}). This substance is the most abundant manganese compound found in nature, and is the ore from which all other compounds of manganese are made. It is a hard, brittle, black substance which is valuable as an oxidizing agent. It will be recalled that it is used in the preparation of chlorine and oxygen, in decolorizing glass which contains iron, and in the manufacture of ferromanganese.
Compounds containing manganese as a base-forming element. As has been stated previously, manganese forms two series of salts. The most important of these salts, all of which belong to the manganous series, are the following:
Manganous chloride MnCl{2}.4H{2}O. Manganous sulphide MnS. Manganous sulphate MnSO{4}.4H{2}O. Manganous carbonate MnCO{3}. Manganous hydroxide Mn(OH){2}.
The chloride and sulphate may be prepared by heating the dioxide with hydrochloric and sulphuric acids respectively:
MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + 2Cl,
MnO_{2} + H_{2}SO_{4} = MnSO_{4} + H_{2}O + O.
The sulphide, carbonate, and hydroxide, being insoluble, may be prepared from a solution of the chloride or sulphate by precipitation with the appropriate reagents. Most of the manganous salts are rose colored. They not only have formulas similar to the ferrous salts, but resemble them in many of their chemical properties. |
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