An Elementary Study of Chemistry, William McPherson [icecream ebook reader txt] 📗
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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:
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
is reversible and has been used as a means of separating oxygen from the air. Treated with acids, barium peroxide yields hydrogen peroxide:
Barium chloride (BaCl2·2H2O). 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:
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.
RADIUMHistorical. In 1896 the French scientist Becquerel observed that the mineral pitchblende possesses certain remarkable properties. It affects photographic plates even in complete darkness, and discharges a 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 700°. 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.
EXERCISES1. 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?
CHAPTER XXV THE MAGNESIUM FAMILYThe 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-earth metals, their hydroxides are nearly insoluble in water. Most of their compounds dissociate in such a way as to give a simple, colorless, metallic ion.
MAGNESIUMOccurrence. 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:
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,—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.
Magnesium cement. A paste of magnesium hydroxide and water slowly absorbs carbon dioxide from the air and becomes very hard. The hardness of the product is increased by the presence of a considerable amount of magnesium chloride in the paste. The hydroxide, with or without the chloride, is used in the preparation of cements for some purposes.
Magnesium carbonate (MgCO3). Magnesium carbonate is a very abundant mineral. It occurs in a number of localities as magnesite, which is usually amorphous, but sometimes forms pure crystals resembling calcite. More commonly it is found associated with calcium carbonate. The mineral dolomite has the composition CaCO3·MgCO3. Limestone containing smaller amounts of magnesium carbonate is known as dolomitic limestone. Dolomite is one of the most common rocks, forming whole mountain masses. It is harder and less readily attacked by acids than limestone. It is valuable as a building stone and as ballast for roadbeds and foundations. Like calcium carbonate, magnesium carbonate is insoluble in water, though easily dissolved by acids.
Basic carbonate of magnesium. We should expect to find magnesium carbonate precipitated when a soluble magnesium salt and a soluble carbonate are brought together:
Instead of this, some carbon dioxide escapes and the product is found to be a basic carbonate. The most common basic carbonate of magnesium has the formula 4MgCO3·Mg(OH)2, and is sometimes called magnesia alba. This compound is formed by the partial hydrolysis of the normal carbonate at first precipitated:
Magnesium chloride (MgCl2·6H2O). Magnesium chloride is found in many natural waters and in many salt deposits (see Stassfurt salts). It is obtained as a by-product in the manufacture of potassium chloride from carnallite. As there is no very important use for it, large quantities annually go to waste. When heated to drive off the water of crystallization the chloride is decomposed as shown in the equation
Owing to the abundance of magnesium chloride, this reaction is being used to some extent in the preparation of both magnesium oxide and hydrochloric acid.
Boiler scale. When water which contains certain salts in solution is evaporated in steam boilers, a hard insoluble material called scale deposits in the boiler. The formation of this scale may be due to several distinct causes.
1. To the deposit of calcium sulphate. This salt, while sparingly soluble in cold water, is almost completely insoluble in superheated water. Consequently it is precipitated when water containing it is heated in a boiler.
2. To decomposition of acid carbonates. As we have seen, calcium and magnesium acid carbonates are decomposed on heating, forming insoluble normal carbonates:
Ca(HCO3)2 = CaCO3 + H2O + CO2.
3. To hydrolysis of magnesium salts. Magnesium chloride, and to some extent magnesium sulphate, undergo hydrolysis when superheated in solution, and the magnesium hydroxide, being sparingly soluble, precipitates:
MgCl2 + 2H2O <--> Mg(OH)2 + 2HCl.
This scale adheres tightly to the boiler in compact layers and, being a non-conductor of heat, causes much waste of fuel. It is very difficult to remove, owing to its hardness and resistance to reagents. Thick scale sometimes cracks, and the water coming in contact with the overheated
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