EXERCISES

(NH4)2CO3= NH4HCO3+ NH3.

The acid carbonate, or bicarbonate, is prepared by saturating a solution of ammonium hydroxide with carbon dioxide:

NH4OH + CO2= NH4HCO3.

It is a well-crystallized stable substance.

Ammonium sulphide((NH4)2S). Ammonium sulphide is prepared by the action of hydrosulphuric acid upon ammonium hydroxide:

2NH4OH + H2S = (NH4)2S + 2H2O.

If the action is allowed to continue until no more hydrosulphuric acid is absorbed, the product is the acid sulphide, sometimes called the hydrosulphide:

NH4OH + H2S = NH4HS + H2O.

If equal amounts of ammonium hydroxide and ammonium acid sulphide are brought together, the normal sulphide is formed:

NH4OH + NH4HS = (NH4)2S + H2O

It has been obtained in the solid state, but only with great difficulty. As used in the laboratory it is always in the form of a solution. It is much used in the process of chemical analysis because it is a soluble sulphide and easily prepared. On exposure to the air ammonium sulphide slowly decomposes, being converted into ammonia, water, and sulphur:

(NH4)2S + O = 2NH3+ H2O + S.

As fast as the sulphur is liberated it combines with the unchanged sulphide to form several different ammonium sulphides in which there are from two to five sulphur atoms in the molecule, thus: (NH4)2S2, (NH4)2S3, (NH4)2S5. These sulphides in turn decompose by further action of oxygen, so that the final products of the reaction are those given in the equation. A solution of these compounds is yellow and is sometimes calledyellow ammonium sulphide.

FLAME REACTIONâ€”SPECTROSCOPEWhen compounds of either sodium or potassium are brought into the non-luminous flame of a Bunsen burner the flame becomes colored. Sodium compounds color it intensely yellow, while those of potassium color it pale violet. When only one of these elements is present it iseasy to identify it by this simple test, but when both are present the intense color of the sodium flame entirely conceals the pale tint characteristic of potassium compounds.It is possible to detect the potassium flame in such cases, however, in the following way. When light is allowed to shine through a very small hole or slit in some kind of a screen, such as a piece of metal, upon a triangular prism of glass, the light is bent or refracted out of its course instead of passing straight through the glass. It thus comes out of the prism at some angle to the line at which it entered. Yellow light is bent more than red, and violet more than yellow. When light made up of the yellow of sodium and the violet of potassium shines through a slit upon such a prism, the yellow and the violet lights come out at somewhat different angles, and so two colored lines of lightâ€”a yellow line and a violet lineâ€”are seen on looking into the prism in the proper direction. The instrument used for separating the rays of light in this way is called aspectroscope(Fig. 79). The material to be tested is placed on a platinum wire and held in the colorless Bunsen flame. The resulting light passes through the slit in the end of tubeB, and then throughBto the prism. The resulting lines of light are seen by looking into the tubeA, which contains a magnifying lens. Most elements give more than one image of the slit, each having a different color, and the series of colored lines due to an element is called its spectrum.

FLAME REACTIONâ€”SPECTROSCOPE

When compounds of either sodium or potassium are brought into the non-luminous flame of a Bunsen burner the flame becomes colored. Sodium compounds color it intensely yellow, while those of potassium color it pale violet. When only one of these elements is present it iseasy to identify it by this simple test, but when both are present the intense color of the sodium flame entirely conceals the pale tint characteristic of potassium compounds.

It is possible to detect the potassium flame in such cases, however, in the following way. When light is allowed to shine through a very small hole or slit in some kind of a screen, such as a piece of metal, upon a triangular prism of glass, the light is bent or refracted out of its course instead of passing straight through the glass. It thus comes out of the prism at some angle to the line at which it entered. Yellow light is bent more than red, and violet more than yellow. When light made up of the yellow of sodium and the violet of potassium shines through a slit upon such a prism, the yellow and the violet lights come out at somewhat different angles, and so two colored lines of lightâ€”a yellow line and a violet lineâ€”are seen on looking into the prism in the proper direction. The instrument used for separating the rays of light in this way is called aspectroscope(Fig. 79). The material to be tested is placed on a platinum wire and held in the colorless Bunsen flame. The resulting light passes through the slit in the end of tubeB, and then throughBto the prism. The resulting lines of light are seen by looking into the tubeA, which contains a magnifying lens. Most elements give more than one image of the slit, each having a different color, and the series of colored lines due to an element is called its spectrum.

Fig. 79Fig. 79

The spectra of the known elements have been carefully studied, and any element which imparts a characteristic color to a flame, or has a spectrum of its own, can be identified even when other elements are present. Through the spectroscopic examination of certain minerals a number of elements have been discovered by the observation of lines which did not belong to any known element. A study of the substance then brought to light the new element. Rubidium and cÃ¦sium were discovered in this way, rubidium having bright red lines and cÃ¦sium a very intense blue line. Lithium colors the flame deep red, and has a bright red line in its spectrum.

1.What is an alkali? Can a metal itself be an alkali?

2.Write equations showing how the following changes may be brought about, giving the general principle involved in each change: NaCl --> Na2SO3, Na2SO3--> NaCl, NaCl --> NaBr, Na2SO4--> NaNO3, NaNO3--> NaHCO3.

3.What carbonates are soluble?

4.State the conditions under which the reaction represented by the following equation can be made to go in either direction:

Na2CO3+ H2O + CO2<--> 2 NaHCO3.

5.Account for the fact that solutions of sodium carbonate and potassium carbonate are alkaline.

6.What non-metallic element is obtained from the deposits of Chili saltpeter?

7.Supposing concentrated hydrochloric acid (den. = 1.2) to be worth six cents a pound, what is the value of the acid generated in the preparation of 1 ton of sodium carbonate by the Le Blanc process?

8.What weight of sodium carbonate crystals will 1 kg. of the anhydrous salt yield?

9.Write equations for the preparation of potassium hydroxide by three different methods.

10.What would take place if a bit of potassium hydroxide were left exposed to the air?

11.Write the equations for the reactions between sodium hydroxide and bromine; between potassium hydroxide and iodine.

12.Write equations for the preparation of potassium sulphate; of potassium acid carbonate.

ROBERT WILHELM BUNSEN (German) (1811-1899) Invented many lecture-room and laboratory appliances (Bunsen burner); invented the spectroscope and with it discovered rubidium and cÃ¦sium; greatly perfected methods of electrolysis, inventing a new battery; made many investigations among metallic and organic substancesROBERT WILHELM BUNSEN (German) (1811-1899)Invented many lecture-room and laboratory appliances (Bunsen burner); invented the spectroscope and with it discovered rubidium and cÃ¦sium; greatly perfected methods of electrolysis, inventing a new battery; made many investigations among metallic and organic substances

13.What weight of carnallite would be necessary in the preparation of 1 ton of potassium carbonate?

14.Write the equations showing how ammonium chloride, ammonium sulphate, ammonium carbonate, and ammonium nitrate may be prepared from ammonium hydroxide.

15.Write an equation to represent the reaction involved in the preparation of ammonia from ammonium chloride.

16.What substances already studied are prepared from the following compounds? ammonium chloride; ammonium nitrate; ammonium nitrite; sodium nitrate; sodium chloride.

17.How could you prove that the water in crystals of common salt is not water of crystallization?

18.How could you distinguish between potassium chloride and potassium iodide? between sodium chloride and ammonium chloride? between sodium nitrate and potassium nitrate?

SYMBOLATOMIC WEIGHTDENSITYMILLIGRAMS SOLUBLE IN 1 L OF WATER AT 18Â°CARBONATE DECOMPOSESSULPHATEHYDROXIDECalciumCa40.11.542070.001670.At dull red heatStrontiumSr87.62.50170.007460.At white heatBariumBa137.43.752.2936300.Scarcely at all

The family.The alkaline-earth family consists of the very abundant element calcium and the much rarer elements strontium and barium. They are called the alkaline-earth metals because their properties are between those of the alkali metals and the earth metals. The earth metals will be discussed in a later chapter. The family is also frequently called the calcium family.

1.Occurrence.These elements do not occur free in nature. Their most abundant compounds are the carbonates and sulphates; calcium also occurs in large quantities as the phosphate and silicate.

2.Preparation.The metals were first prepared by Davy in 1808 by electrolysis. This method has again come into use in recent years. Strontium and barium have as yet been obtained only in small quantities and in the impure state, and many of their physical properties,such as their densities and melting points, are therefore imperfectly known.

3.Properties.The three metals resemble each other very closely. They are silvery-white in color and are about as hard as lead. Their densities increase with their atomic weights, as is shown in the table on opposite page. Like the alkali metals they have a strong affinity for oxygen, tarnishing in the air through oxidation. They decompose water at ordinary temperatures, forming hydroxides and liberating hydrogen. When ignited in the air they burn with brilliancy, forming oxides of the general formula MO. These oxides readily combine with water, according to the equation

MO + H2O = M(OH)2.

Each of the elements has a characteristic spectrum, and the presence of the metals can easily be detected by the spectroscope.

4.Compounds.The elements are divalent in almost all of their compounds, and these compounds in solution give simple, divalent, colorless ions. The corresponding salts of the three elements are very similar to each other and show a regular variation in properties in passing from calcium to strontium and from strontium to barium. This is seen in the solubility of the sulphate and hydroxide, and in the ease of decomposition of the carbonates, as given in the table. Unlike the alkali metals, their normal carbonates and phosphates are insoluble in water.

Occurrence.The compounds of calcium are very abundant in nature, so that the total amount of calcium in the earth's crust is very large. A great many different compoundscontaining the clement are known, the most important of which are the following:

Calcite (marble)CaCO3.PhosphoriteCa3(PO4)2.FluorsparCaF2.WollastoniteCaSiO3.GypsumCaSO4Â·2H2O.AnhydriteCaSO4.

Preparation.Calcium is now prepared by the electrolysis of the melted chloride, the metal depositing in solid condition on the cathode. It is a gray metal, considerably heavier and harder than sodium. It acts upon water, forming calcium hydroxide and hydrogen, but the action does not evolve sufficient heat to melt the metal. It promises to become a useful substance, though no commercial applications for it have as yet been found.

Calcium oxide(lime, quicklime) (CaO). Lime is prepared by strongly heating calcium carbonate (limestone) in large furnaces called kilns:

CaCO3= CaO + CO2.

When pure, lime is a white amorphous substance. Heated intensely, as in the oxyhydrogen flame, it gives a brilliant light called the lime light. Although it is a very difficultly fusible substance, yet in the electric furnace it can be made to melt and even boil. Water acts upon lime with the evolution of a great deal of heat,â€”hence the name quicklime, or live lime,â€”the process being called slaking. The equation is

CaO + H2O = Ca(OH)2.

Lime readily absorbs moisture from the air, and is used to dry moist gases, especially ammonia, which cannot bedried by the usual desiccating agents. It also absorbs carbon dioxide, forming the carbonate

CaO + CO2= CaCO3.

Lime exposed to air is therefore gradually converted into hydroxide and carbonate, and will no longer slake with water. It is then said to be air-slaked.

Limekilns.The older kiln, still in common use, consists of a large cylindrical stack in which the limestone is loosely packed. A fire is built at the base of the stack, and when the burning is complete it is allowed to die out and the lime is removed from the kiln. The newer kilns are constructed as shown in Fig. 80. A number of fire boxes are built around the lower part of the kiln, one of which is shown atB. The fire is built on the grateFand the hot products of combustion are drawn up through the stack, decomposing the limestone. The kiln is charged atC, and sometimes fuel is added with the limestone to cause combustion throughout the contents of the kiln. The burned lime is raked out through openings in the bottom of the stack, one of which is shown atD.The advantage of this kind of a kiln over the older form is that the process is continuous, limestone being charged in at the top as fast as the lime is removed at the bottom.

Fig. 80Fig. 80

Calcium hydroxide(slaked lime) (Ca(OH)2). Pure calcium hydroxide is a light white powder. It is sparingly soluble in water, forming a solution calledlimewater, which is often used in medicine as a mild alkali. Chemically, calcium hydroxide is a moderately strong base, though not so strong as sodium hydroxide. Owing to its cheapness it is much used in theindustries whenever an alkali is desired. A number of its uses have already been mentioned. It is used in the preparation of ammonia, bleaching powder, and potassium hydroxide. It is also used to remove carbon dioxide and sulphur compounds from coal gas, to remove the hair from hides in the tanneries (this recalls the caustic or corrosive properties of sodium hydroxide), and for making mortar.

Mortaris a mixture of calcium hydroxide and sand. When it is exposed to the air or spread upon porous materials moisture is removed from it partly by absorption in the porous materials and partly by evaporation, and the mortar becomes firm, orsets. At the same time carbon dioxide is slowly absorbed from the air, forming hard calcium carbonate:

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

By this combined action the mortar becomes very hard and adheres firmly to the surface upon which it is spread. The sand serves to give body to the mortar and makes it porous, so that the change into carbonate can take place throughout the mass. It also prevents too much shrinkage.

Cement.When limestone to which clay and sand have been added in certain proportions is burned until it is partly fused (some natural marl is already of about the right composition), and the clinker so produced is ground to powder, the product is called cement. When this material is moistened it sets to a hard stone-like mass which retains its hardness even when exposed to the continued action of water. It can be used for under-water work, such as bridge piers, where mortar would quickly soften. Several varieties of cement are made, the best known of which is Portland cement.

Growing importance of cement.Cement is rapidly coming into use for a great variety of purposes. It is often used in place of mortar in the construction of brick buildings. Mixed with crushed stone and sand it forms concrete which is used in foundation work. It is also used in making artificial stone, terra-cotta trimmings for buildings, artificial stone walks and floors, and the like. It is being used more and more for making many articles which were formerly made of wood or stone, and the entire walls of buildings are sometimes made of cement blocks or of concrete.

Calcium carbonate(CaCO3). This substance is found in a great many natural forms to which various names have been given. They may be classified under three heads:

1.Amorphous carbonate.This includes those forms which are not markedly crystalline. Limestone is the most familiar of these and is a grayish rock usually found in hard stratified masses. Whole mountain ranges are sometimes made up of this material. It is always impure, usually containing magnesium carbonate, clay, silica, iron and aluminium compounds, and frequently fossil remains. Marl is a mixture of limestone and clay. Pearls, chalk, coral, and shells are largely calcium carbonate.

2.Hexagonal carbonate.Calcium carbonate crystallizes in the form of rhomb-shaped crystals which belong to the hexagonal system. When very pure and transparent the substance is called Iceland spar. Calcite is a similar form, but somewhat opaque or clouded. Mexican onyx is a massive variety, streaked or banded with colors due to impurities. Marble when pure is made up of minute calcite crystals. Stalactites and stalagmites are icicle-like forms sometimes found in caves.

3.Rhombic carbonate.Calcium carbonate sometimes crystallizes in needle-shaped crystals belonging to the rhombic system. This is the unstable form and tends togo over into the other variety. Aragonite is the most familiar example of this form.

Preparation and uses of calcium carbonate.In the laboratory pure calcium carbonate can be prepared by treating a soluble calcium salt with a soluble carbonate:

Na2CO3+ CaCl2= CaCO3+ 2NaCl.

When prepared in this way it is a soft white powder often called precipitated chalk, and is much used as a polishing powder. It is insoluble in water, but dissolves in water saturated with carbon dioxide, owing to the formation of the acid calcium carbonate which is slightly soluble:

CaCO3+ H2CO3= Ca(HCO3)2.

The natural varieties of calcium carbonate find many uses, such as in the preparation of lime and carbon dioxide; in metallurgical operations, especially in the blast furnaces; in the manufacture of soda, glass, and crayon (which, in addition to chalk, usually contains clay and calcium sulphate); for building stone and ballast for roads.

Calcium chloride(CaCl2). This salt occurs in considerable quantity in sea water. It is obtained as a by-product in many technical processes, as in the Solvay soda process. When crystallized from its saturated solutions it forms colorless needles of the composition CaCl2Â·6H2O. By evaporating a solution to dryness and heating to a moderate temperature calcium chloride is obtained anhydrous as a white porous mass. In this condition it absorbs water with great energy and is a valuable drying agent.

Bleaching powder(CaOCl2). When chlorine acts upon a solution of calcium hydroxide the reaction is similar to that which occurs between chlorine and potassium hydroxide:

2Ca(OH)2+ 4Cl = CaCl2+ Ca(ClO)2+ 2H2O.

If, however, chlorine is conducted over calcium hydroxide in the form of a dry powder, it is absorbed and a substance is formed which appears to have the composition represented in the formula CaOCl2. This substance is called bleaching powder, or hypochlorite of lime. It is probably the calcium salt of both hydrochloric and hypochlorous acids, so that its structure is represented by the formula

/ClOCa\Cl.

In solution this substance acts exactly like a mixture of calcium chloride (CaCl2) and calcium hypochlorite (Ca(ClO)2), since it dissociates to form the ions Ca++, Cl-, and ClO-.

Bleaching powder undergoes a number of reactions which make it an important substance.

1. When treated with an acid it evolves chlorine:

/ClOCa Â Â Â Â Â + H2SO4= CaSO4+ HCl + HClO,\ClHCl + HClO = H2O + 2Cl.

/ClOCa Â Â Â Â Â + H2SO4= CaSO4+ HCl + HClO,\Cl

HCl + HClO = H2O + 2Cl.

This reaction can be employed in the preparation of chlorine, or the nascent chlorine may be used as a bleaching agent.

2. It is slowly decomposed by the carbon dioxide of the air, yielding calcium carbonate and chlorine:

CaOCl2+ CO2= CaCO3+ 2Cl.

Owing to this slow action the substance is a good disinfectant.

3. When its solution is boiled the substance breaks down into calcium chloride and chlorate:

6CaOCl2= 5CaCl2+ Ca(ClO3)2.

This reaction is used in the preparation of potassium chlorate.

Calcium fluoride(fluorspar) (CaF2). Fluorspar has already been mentioned as the chief natural compound of fluorine. It is found in large quantities in a number of localities, and is often crystallized in perfect cubes of a light green or amethyst color. It can be melted easily in a furnace, and is sometimes used in the fused condition in metallurgical operations to protect a metal from the action of the air during its reduction. It is used as the chief source of fluorine compounds, especially hydrofluoric acid.

Calcium sulphate(gypsum) (CaSO4Â·2H2O). This abundant substance occurs in very perfectly formed crystals or in massive deposits. It is often found in solution in natural waters and in the sea water. Salts deposited from sea water are therefore likely to contain this substance (see Stassfurt salts).

It is very sparingly soluble in water, and is thrown down as a fine white precipitate when any considerable amounts of a calcium salt and a soluble sulphate (or sulphuric acid) are brought together in solution. Its chief use is in the manufacture of plaster of Paris and of hollow tiles for fireproof walls. Such material is calledgypsite. It is also used as a fertilizer.

Calcium sulphate, like the carbonate, occurs in many forms in nature. Gypsum is a name given to all common varieties. Granular or massive specimens are called alabaster, while all those which are well crystallized are called selenite. Satin spar is still another variety often seen in mineral collections.

Plaster of Paris.When gypsum is heated to about 115Â° it loses a portion of its water of crystallization in accordance with the equation

2(CaSO4Â·2H2O) = 2CaSO4Â·H2O + 2H2O.

The product is a fine white powder calledplaster of Paris. On being moistened it again takes up this water, and in so doing first forms a plastic mass, which soon becomes very firm and hard and regains its crystalline structure. These properties make it very valuable as a material for forming casts and stucco work, for cementing glass to metals, and for other similar purposes. If overheated so that all water is driven off, the process of taking up water is so slow that the material is worthless. Such material is said to be dead burned. Plaster of Paris is very extensively used as the finishing coat for plastered walls.

Hard water.Waters containing compounds of calcium and magnesium in solution are called hard waters because they feel harsh to the touch. The hardness of water may be of two kinds,â€”(1) temporary hardness and (2) permanent hardness.

1.Temporary hardness.We have seen that when water charged with carbon dioxide comes in contact with limestone a certain amount of the latter dissolves, owing to the formation of the soluble acid carbonate of calcium. The hardness of such waters is said to be temporary, since it may be removed by boiling. The heat changes the acid carbonate into the insoluble normal carbonate which then precipitates, rendering the water soft:

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

Such waters may also be softened by the addition of sufficient lime or calcium hydroxide to convert the acid carbonate of calcium into the normal carbonate. The equation representing the reaction is

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

2.Permanent hardness.The hardness of water may also be due to the presence of calcium and magnesium sulphates or chlorides. Boiling the water does not affect these salts; hence such waters are said to have permanent hardness. They may be softened, however, by the addition of sodium carbonate, which precipitates the calcium and magnesium as insoluble carbonates:

CaSO4+ Na2CO3= CaCO3+ Na2SO4.

This process is sometimes called "breaking" the water.

Commercial methods for softening water.The average water of a city supply contains not only the acid carbonates of calcium and magnesium but also the sulphates and chlorides of these metals, together with other salts in smaller quantities. Such waters are softened on a commercial scale by the addition of the proper quantities of calcium hydroxide and sodium carbonate. The calcium hydroxide is added first to precipitate all the acid carbonates. After a short time the sodium carbonate is added to precipitate the other soluble salts of calcium and magnesium, together with any excess of calcium hydroxide which may have been added. The quantity of calcium hydroxide and sodium carbonate required is calculated from a chemical analysis of the water. It will be noticed that the water softened in this way will contain sodium sulphate and chloride, but the presence of these salts is not objectionable.

Calcium carbide(CaC2). This substance is made by heating well-dried coke and lime in an electrical furnace. The equation is

CaO + 3C = CaC2+ CO.

The pure carbide is a colorless, transparent, crystalline substance. In contact with water it is decomposed with the evolution of pure acetylene gas, having a pleasant ethereal odor. The commercial article is a dull gray porous substance which contains many impurities. The acetylene prepared from this substance has a very characteristic odordue to impurities, the chief of these being phosphine. It is used in considerable quantities as a source of acetylene gas for illuminating purposes.

Technical preparation.Fig. 81 represents a recent type of a carbide furnace. The base of the furnace is provided with a large block of carbonA, which serves as one of the electrodes. The other electrodesB, several in number, are arranged horizontally at some distance above this. A mixture of coal and lime is fed into the furnace through the trap topC, and in the lower part of the furnace this mixture becomes intensely heated, forming liquid carbide. This is drawn off through the tapholeD.The carbon monoxide formed in the reaction escapes through the pipesEand is led back into the furnace. The pipesFsupply air, so that the monoxide burns as it reÃ«nters the furnace and assists in heating the charge. The carbon dioxide so formed, together with the nitrogen entering as air, escape atG. An alternating current is used.

The carbon monoxide formed in the reaction escapes through the pipesEand is led back into the furnace. The pipesFsupply air, so that the monoxide burns as it reÃ«nters the furnace and assists in heating the charge. The carbon dioxide so formed, together with the nitrogen entering as air, escape atG. An alternating current is used.

Fig. 81Fig. 81

Calcium phosphate(Ca3(PO4)2). This important substance occurs abundantly in nature as a constituent of apatite (3 Ca3(PO4)2Â·CaF2), in phosphate rock, and as the chief mineral constituent of bones. Bone ash is therefore nearly pure calcium phosphate. It is a white powder, insoluble in water, although it readily dissolves in acids, being decomposed by them and converted into soluble acid phosphates, as explained in connection with the acids of phosphorus.

Occurrence.Strontium occurs sparingly in nature, usually as strontianite (SrCO3) and as celestite (SrSO4). Both minerals form beautiful colorless crystals, though celestite is sometimes colored a faint blue. Only a few of the compounds of strontium have any commercial applications.

Strontium hydroxide(Sr(OH)2Â·8H2O). The method of preparation of strontium hydroxide is analogous to that of calcium hydroxide. The substance has the property of forming an insoluble compound with sugar, which can easily be separated again into its constituents. It is therefore sometimes used in the sugar refineries to extract sugar from impure mother liquors from which the sugar will not crystallize.

Strontium nitrate(Sr(NO3)2Â·4H2O). This salt is prepared by treating the native carbonate with nitric acid. When ignited with combustible materials it imparts a brilliant crimson color to the flame, and because of this property it is used in the manufacture of red lights.

Barium is somewhat more abundant than strontium, occurring in nature largely as barytes, or heavy spar (BaSO4), and witherite (BaCO3). Like strontium, it closely resembles calcium both in the properties of the metal and in the compounds which it forms.

Oxides of barium.Barium oxide (BaO) can be obtained by strongly heating the nitrate:

Ba(NO3)2= BaO + 2NO2+ O.

Heated to a low red heat in the air, the oxide combines with oxygen,forming the peroxide (BaO2). If the temperature is raised still higher, or the pressure is reduced, oxygen is given off and the oxide is once more formed. The reaction

BaO2<--> BaO + O

is reversible and has been used as a means of separating oxygen from the air. Treated with acids, barium peroxide yields hydrogen peroxide:

BaO2+ 2HCl = BaCl2+ H2O2.

Barium chloride(BaCl2Â·2H2O). Barium chloride is a white well-crystallized substance which is easily prepared from the native carbonate. It is largely used in the laboratory as a reagent to detect the presence of sulphuric acid or soluble sulphates.

Barium sulphate(barytes)(BaSO4). Barium sulphate occurs in nature in the form of heavy white crystals. It is precipitated as a crystalline powder when a barium salt is added to a solution of a sulphate or sulphuric acid:

BaCl2+ H2SO4= BaSO4+ 2HCl.

This precipitate is used, as are also the finely ground native sulphate and carbonate, as a pigment in paints. On account of its low cost it is sometimes used as an adulterant of white lead, which is also a heavy white substance.

Barium compounds color the flame green, and the nitrate (Ba(NO3)2) is used in the manufacture of green lights. Soluble barium compounds are poisonous.

Historical.In 1896 the French scientist Becquerel observed that the mineral pitchblende possesses certain remarkable properties. It affects photographic plates even in complete darkness, and dischargesa gold-leaf electroscope when brought close to it. In 1898 Madam Curie made a careful study of pitchblende to see if these properties belong to it or to some unknown substance contained in it. She succeeded in extracting from it a very small quantity of a substance containing a new element which she named radium.

In 1910 Madam Curie succeeded in obtaining radium itself by the electrolysis of radium chloride. It is a silver-white metal melting at about 700Â°. It blackens in the air, forming a nitride, and decomposes water. Its atomic weight is about 226.5.

Properties.Compounds of radium affect a photographic plate or electroscope even through layers of paper or sheets of metal. They also bring about chemical changes in substances placed near them. Investigation of these strange properties has suggested that the radium atoms are unstable and undergo a decomposition. As a result of this decomposition very minute bodies, to which the name corpuscles has been given, are projected from the radium atom with exceedingly great velocity. It is to these corpuscles that the strange properties of radium are due. It seems probable that the gas helium is in some way formed during the decomposition of radium.

Two or three other elements, particularly uranium and thorium, have been found to possess many of the properties of radium in smaller degree.

Radium and the atomic theory.If these views in regard to radium should prove to be well founded, it will be necessary to modify in some respects the conception of the atom as developed in a former chapter. The atom would have to be regarded as a compound unit made up of several parts. In a few cases, as in radium and uranium, it would appear that this unit is unstable and undergoes transformation into more stable combinations. This modification would not, in any essential way, be at variance with the atomic theory as propounded by Dalton.

1.What properties have the alkaline-earth metals in common with the alkali metals? In what respects do they differ?

2.Write the equation for the reaction between calcium carbide and water.

3.For what is calcium chlorate used?

4.Could limestone be completely decomposed if heated in a closed vessel?

5.Caves often occur in limestone. Account for their formation.

6.What is the significance of the term fluorspar? (Consult dictionary.)

7.Could calcium chloride be used in place of barium chloride in testing for sulphates?

8.What weight of water is necessary to slake the lime obtained from 1 ton of pure calcium carbonate?

9.What weight of gypsum is necessary in the preparation of 1 ton of plaster of Paris?

10.Write equations to represent the reactions involved in the preparation of strontium hydroxide and strontium nitrate from strontianite.

11.Write equations to represent the reactions involved in the preparation of barium chloride from heavy spar.

12.Could barium hydroxide be used in place of calcium hydroxide in testing for carbon dioxide?

SYMBOLATOMIC WEIGHTDENSITYMELTING POINTBOILING POINTOXIDEMagnesiumMg24.361.75750Â°920Â°MgOZincZn65.47.00420Â°950Â°ZnOCadmiumCd112.48.67320Â°778Â°CdO

The family.In the magnesium family are included the four elements: magnesium, zinc, cadmium, and mercury. Between the first three of these metals there is a close family resemblance, such as has been traced between the members of the two preceding families. Mercury in some respects is more similar to copper and will be studied in connection with that metal.

1.Properties.When heated to a high temperature in the air each of these metals combines with oxygen to form an oxide of the general formula MO, in which M represents the metal. Magnesium decomposes boiling water slowly, while zinc and cadmium have but little action on it.

2.Compounds.The members of this group are divalent in nearly all their compounds, so that the formulas of their salts resemble those of the alkaline-earth metals. Like the alkaline-earth metals, their carbonates and phosphates are insoluble in water. Their sulphates, however, are readily soluble. Unlike both the alkali and alkaline-earthmetals, their hydroxides are nearly insoluble in water. Most of their compounds dissociate in such a way as to give a simple, colorless, metallic ion.

Occurrence.Magnesium is a very abundant element in nature, ranking a little below calcium in this respect. Like calcium, it is a constituent of many rocks and also occurs in the form of soluble salts.

Preparation.The metal magnesium, like most metals whose oxides are difficult to reduce with carbon, was formerly prepared by heating the anhydrous chloride with sodium:

MgCl2+ 2Na = 2NaCl + Mg.

It is now made by electrolysis, but instead of using as the electrolyte the melted anhydrous chloride, which is difficult to obtain, the natural mineral carnallite is used. This is melted in an iron pot which also serves as the cathode in the electrolysis. A rod of carbon dipping into the melted salt serves as the anode. The apparatus is very similar to the one employed in the preparation of sodium.

Properties.Magnesium is a rather tough silvery-white metal of small density. Air does not act rapidly upon it, but a thin film of oxide forms upon its surface, dimming its bright luster. The common acids dissolve it with the formation of the corresponding salts. It can be ignited readily and in burning liberates much heat and gives a brilliant white light. This light is very rich in the rays which affect photographic plates, and the metal in the form of fine powder is extensively used in the production of flash lights and for white lights in pyrotechnic displays.

Magnesium oxide(magnesia) (MgO). Magnesium oxide, sometimes called magnesia or magnesia usta, resembles lime in many respects. It is much more easily formed than lime and can be made in the same way,â€”by igniting the carbonate. It is a white powder, very soft and light, and is unchanged by heat even at very high temperatures. For this reason it is used in the manufacture of crucibles, for lining furnaces, and for other purposes where a refractory substance is needed. It combines with water to form magnesium hydroxide, but much more slowly and with the production of much less heat than in the case of calcium oxide.

Magnesium hydroxide(Mg(OH)2). The hydroxide formed in this way is very slightly soluble in water, but enough dissolves to give the water an alkaline reaction. Magnesium hydroxide is therefore a fairly strong base. It is an amorphous white substance. Neither magnesia nor magnesium salts have a very marked effect upon the system; and for this reason magnesia is a very suitable antidote for poisoning by strong acids, since any excess introduced into the system will have no injurious effect.

Back to Index Next