CALCIUM

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 compounds containing the clement are known, the most important of which are the following:

Calcite (marble)	CaCO3.
Phosphorite	Ca3(PO4)2.
Fluorspar	CaF2.
Wollastonite	CaSiO3.
Gypsum	CaSO4·2H2O.
Anhydrite	CaSO4.

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,—hence the name quicklime, or live lime,—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 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.

Fig. 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 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:

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 to go 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 CaCl₂·6H₂O. 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

/ClO
Ca
\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:

/ClO
Ca + 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) (CaSO₄·2H₂O). 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 called gypsite. 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 115° it loses a portion of its water of crystallization in accordance with the equation

2(CaSO₄·2H₂O) = 2CaSO₄·H₂O + 2H₂O.

The product is a fine white powder called plaster 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,—(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 odor due 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 carbon A, which serves as one of the electrodes. The other electrodes B, 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 top C, and in the lower part of the furnace this mixture becomes intensely heated, forming liquid carbide. This is drawn off through the taphole D.

The carbon monoxide formed in the reaction escapes through the pipes E and is led back into the furnace. The pipes F supply air, so that the monoxide burns as it reënters the furnace and assists in heating the charge. The carbon dioxide so formed, together with the nitrogen entering as air, escape at G. An alternating current is used.

Fig. 81

Calcium phosphate (Ca₃(PO₄)₂). This important substance occurs abundantly in nature as a constituent of apatite (3 Ca₃(PO₄)₂·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.