THE ALKALINE-EARTH FAMILY
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| | | | |
| | | | MILLIGRAMS SOL- |
| | | | UBLE IN 1 L. |
| | | | OF WATER AT 18° |
| SYMBOL | ATOMIC | DENSITY |__________________| CARBONATE
| | WEIGHT | | | | DECOMPOSES
| | | | SULPHATE| HYDROX-|
| | | | | IDE |
__________|________|________|_________|_________|________|_________________
| | | | | |
Calcium | Ca | 40.1 | 1.54 | 2070.00 | 1670. | At dull red heat
Strontium | Sr | 87.6 | 2.50 | 170.00 | 7460. | At white heat
Barium | Ba | 137.4 | 3.75 | 2.29 | 36300. | Scarcely at all
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~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 + H_{2}O = 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.
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) CaCO_{3}.
Phosphorite Ca_{3}(PO_{4})_{2}.
Fluorspar CaF_{2}.
Wollastonite CaSiO_{3}.
Gypsum CaSO_{4}·2H_{2}O.
Anhydrite CaSO_{4}.
~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:
CaCO_{3} = CaO + CO_{2}.
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 + H_{2}O = 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 + CO_{2} = CaCO_{3}.
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.
[Illustration: 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} + CO_{2} = CaCO_{3} + H_{2}O.
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~ (CaCO_{3}). 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:
Na_{2}CO_{3} + CaCl_{2} = CaCO_{3} + 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:
CaCO_{3} + H_{2}CO_{3} = Ca(HCO_{3})_{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~ (CaCl_{2}). 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
CaCl_{2}·6H_{2}O. 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~ (CaOCl_{2}). When chlorine acts upon a solution of
calcium hydroxide the reaction is similar to that which occurs between
chlorine and potassium hydroxide:
2 Ca(OH)_{2} + 4 Cl = CaCl_{2} + Ca(ClO)_{2} + 2 H_{2}O.
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 CaOCl_{2}. 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 (CaCl_{2}) 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 + H_{2}SO_{4} = CaSO_{4} + HCl + HClO,
\Cl
HCl + HClO = H_{2}O + 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:
CaOCl_{2} + CO_{2} = CaCO_{3} + 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:
6CaOCl_{2} = 5CaCl_{2} + Ca(ClO_{3})_{2}.
This reaction is used in the preparation of potassium chlorate.
~Calcium fluoride~ (_fluorspar_) (CaF_{2}). 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_) (CaSO_{4}·2H_{2}O). 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(CaSO_{4}·2H_{2}O) = 2CaSO_{4}·H_{2}O + 2H_{2}O.
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(HCO_{3})_{2} = CaCO_{3} + H_{2}O + CO_{2}.
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(HCO_{3})_{2} + Ca(OH)_{2} = 2CaCO_{3} + 2H_{2}O.
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:
CaSO_{4} + Na_{2}CO_{3} = CaCO_{3} + Na_{2}SO_{4}.
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~ (CaC_{2}). This substance is made by heating well-dried
coke and lime in an electrical furnace. The equation is
CaO + 3C = CaC_{2} + 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.
[Illustration: Fig. 81]
~Calcium phosphate~ (Ca_{3}(PO_{4})_{2}). This important substance
occurs abundantly in nature as a constituent of apatite
(3Ca_{3}(PO_{4})_{2}·CaF_{2}), 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.
STRONTIUM
~Occurrence.~ Strontium occurs sparingly in nature, usually as
strontianite (SrCO_{3}) and as celestite (SrSO_{4}). 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)_{2}·8H_{2}O). 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(NO_{3})_{2}·4H_{2}O). 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
Barium is somewhat more abundant than strontium, occurring in nature
largely as barytes, or heavy spar (BaSO_{4}), and witherite (BaCO_{3}).
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(NO_{3})_{2} = BaO + 2NO_{2} + O.
Heated to a low red heat in the air, the oxide combines with oxygen,
forming the peroxide (BaO_{2}). 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
BaO_{2} <--> 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:
BaO_{2} + 2HCl = BaCl_{2} + H_{2}O_{2}.
~Barium chloride~ (BaCl_{2}·2H_{2}O). 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)_ (BaSO_{4}). 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:
BaCl_{2} + H_{2}SO_{4} = BaSO_{4} + 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(NO_{3})_{2})
is used in the manufacture of green lights. Soluble barium compounds are
poisonous.
RADIUM
~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 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.
EXERCISES
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?