Magnesium the BANTAMWEIGHT METAL (Aug, 1946)
Magnesium the BANTAMWEIGHT METAL
How Chemists Have Put It to Work as a Jack-of-All-Trades.
By KENNETH M. SWEZEY
DURING the war magnesium was extensively used as a lightweight structural metal for aircraft parts and as pyrotechnic material for star shells, signal flares, tracer bullets, and flash and incendiary bombs. Strong, silvery white, and only two thirds as heavy as aluminum, it is the lightest of all construction metals. In the form of powder, thin sheets, or wire, it burns with a dazzling flame that water or even carbon dioxide will not put out. Never found alone in nature, magnesium is made on a tremendous scale by the electrolysis of its compounds. These compounds are among the most plentiful substances in the crust of the earth. Whole mountain ranges consist of dolomite, a double carbonate of magnesium and calcium. Asbestos, talc, and meerschaum are magnesium silicates. Epsom salts, named after the springs at Epsom, England, where they were first isolated in 1695, are magnesium sulphate. In the form of its chloride, there are nearly 6,000,000 tons of magnesium in every cubic mile of-the sea, a vast storehouse of supply.
Less spectacular, perhaps, than the metal, the compounds of magnesium are just as important. Asbestos and magnesium oxide are among our most valuable insulators against heat. Magnesium oxychloride forms a superior artificial stone and flooring material. Magnesium carbonate is used for insulation and for making dentifrices, talcum powder, other magnesium chemicals, and Pyrex glass. Epsom, salts, citrate of magnesia, and milk of magnesia are three of the many magnesium compounds that are employed in medicine.
With only a box of Epsom salts as a starting material, you can make many of these compounds in your home laboratory. With a few inches of magnesium ribbon, you can likewise test some of the exciting properties of the metal itself.
One property that helps make the metal such an important incendiary material for wartime use is its ability to steal oxygen from such ordinarily stable compounds as water and carbon dioxide. During the war, magnesium fires generally were extinguished by smothering with sand. Water helped only when applied in quantities sufficient to cool the metal below the point of combustion. This was rarely possible.
As a demonstration of magnesium as an oxygen grabber, boil some water in a flask, and then with tongs lower a short length of lighted magnesium ribbon into the steam.
(Put away the rest of the ribbon before lighting the piece, and always handle with caution.) Instead of going out, the magnesium continues to burn brightly, getting oxygen by decomposing the steam.
Carbon dioxide, usually one of the best fire-extinguishing materials, is as helpless as steam against burning magnesium. Fill a beaker with this gas by pouring 1/2″ of water into it and adding a little baking soda (sodium bicarbonate) and vinegar or other common acid. As soon as the bubbling has stopped, test for carbon dioxide by lowering a lighted match into the beaker. The match will go out at once. Now lower a piece of burning magnesium into the glass and see the difference. The metal continues to burn furiously. In the process, it changes into magnesium oxide, and black specks of carbon, wrested from the carbon dioxide gas, are flung to the sides and bottom of the beaker.
Even nitrogen, which ordinarily is one of the most inert gases, will unite with hot magnesium when conditions are right, forming magnesium nitride. To show this, put a few short pieces of magnesium ribbon on the center of an upturned can cover and heat the cover over a gas flame until the metal catches fire. Then remove the cover from the flame and allow it to cool until you can touch it with your hand. Now put several drops of water on the warm substance that remains and hold a bit of cotton wool moistened with hydrochloric acid above it. White smoke of ammonium chloride immediately rises. On burning, the magnesium united with oxygen and nitrogen from the air, forming the oxide and nitride. When water was added, the nitride decomposed into ammonia gas and magnesium hydroxide.
Most magnesium compounds can be produced from the carbonate or the hydroxide. The carbonate occurs naturally as magnesite and, mixed with calcium carbonate, in certain forms of marble and limestone. It can be made artificially by mixing hot solutions of magnesium sulphate (Epsom salts) and sodium carbonate (washing soda). Since the carbonate is not soluble in water, it is precipitated as a fine white powder. When dried, it can be used as a polishing agent and for heat insulation.
Magnesium hydroxide also is made by precipitation. Again you can start with your Epsom salts, this time adding a solution of sodium hydroxide (ordinary lye). When dissolved, both of these solids produce clear solutions. Mixed together, they form a white precipitate. A suspension of this, in pure form, is the drug-store “milk of magnesia.” By strongly heating either your hydroxide or carbonate, you produce magnesium oxide (magnesia), used widely for heat insulating, for the lining of high-temperature furnaces, and for making oxychloride cement, widely used for flooring and imitative stone.
By dissolving either the carbonate or hydroxide in hydrochloric acid you get magnesium chloride, the compound found in sea water from which magnesium metal is made in vast quantities. To obtain crystals of this chemical, evaporate the solution over a water bath until nearly dry. Then complete the drying in a warm dry place in the open air.
If the heating is continued too long, the crystals will lose some of their water content and partly decompose, giving up hydrochloric acid.
This decomposition of magnesium chloride by heat provides one method for the manufacture of hydrochloric acid. As a demonstration, put some crystals of magnesium chloride in a test tube and fit the tube with a stopper having a glass delivery tube long enough to reach the bottom of a second test tube, which is half filled with water. Now gently heat the crystals. They will first melt, and then vapors will go through the delivery tube and bubble up through the water. Part will be water vapor and part hydrogen chloride. The latter will dissolve in the water, changing it into hydrochloric acid.