An Elementary Study of Chemistry, William McPherson [icecream ebook reader txt] 📗
- Author: William McPherson
- Performer: -
Book online «An Elementary Study of Chemistry, William McPherson [icecream ebook reader txt] 📗». Author William McPherson
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 cæsium were discovered in this way, rubidium having bright red lines and cæsium a very intense blue line. Lithium colors the flame deep red, and has a bright red line in its spectrum.
EXERCISES1. 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:
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.
Invented many lecture-room and laboratory appliances (Bunsen burner); invented the spectroscope and with it discovered rubidium and cæsium; 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?
CHAPTER XXIV THE ALKALINE-EARTH FAMILYThe 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
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.
CALCIUMOccurrence. 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 compounds containing the clement are known, the most important of which are the following:
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:
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,—hence the name quicklime, or live lime,—the process being called slaking. The equation is
Lime readily absorbs moisture from the air, and is used to dry moist gases, especially ammonia, which cannot be dried 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 at B. The fire is built on the grate F and the hot products of combustion are drawn up through the stack, decomposing the limestone. The kiln is charged at C, 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 at D. 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.
Calcium hydroxide (slaked lime) (Ca(OH)2). Pure calcium hydroxide is a light white powder. It is sparingly soluble in water, forming a solution called limewater, 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 the industries 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.
Mortar is 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, or sets. At the same time carbon dioxide is slowly absorbed from the air, forming hard calcium carbonate:
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
Comments (0)