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An Elementary Study of Chemistry

W >> William McPherson >> An Elementary Study of Chemistry




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 deg. 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 reenters 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 deg.. 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?




CHAPTER XXV

THE MAGNESIUM FAMILY


===========================================================================
|SYMBOL |ATOMIC |DENSITY |MELTING |BOILING | OXIDE
| |WEIGHT | | POINT | POINT |
---------------------------------------------------------------------------
Magnesium | Mg | 24.36 | 1.75 | 750 deg. | 920 deg. | MgO
Zinc | Zn | 65.4 | 7.00 | 420 deg. | 950 deg. | ZnO
Cadmium | Cd |112.4 | 8.67 | 320 deg. | 778 deg. | CdO
===========================================================================

~The 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.


MAGNESIUM

~Occurrence.~ 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:

MgCl_{2} + 2Na = 2NaCl + Mg.

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~ (MgCO_{3}). 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 CaCO_{3}.MgCO_{3}. 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:

Na_{2}CO_{3} + MgCl_{2} = MgCO_{3} + 2NaCl.

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 4MgCO_{3}.Mg(OH)_{2}, and is sometimes called magnesia alba.
This compound is formed by the partial hydrolysis of the normal
carbonate at first precipitated:

5MgCO_{3} + 2H_{2}O = 4MgCO_{3}.Mg(OH)_{2} + H_{2}CO_{3}.

~Magnesium chloride~ (MgCl_{2}.6H_{2}O). 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

MgCl_{2}.6H_{2}O = MgO + 2HCl + 5H_{2}O.

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(HCO_{3})_{2} = CaCO_{3} + H_{2}O + CO_{2}.

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:

MgCl_{2} + 2H_{2}O <--> 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 iron occasions an explosion.
Moreover, the acids set free in the hydrolysis of the magnesium
salts attack the iron tubes and rapidly corrode them. These
causes combine to make the formation of scale a matter which
occasions much trouble in cases where hard water is used in
steam boilers. Water containing such salts should be softened,
therefore, before being used in boilers.

~Magnesium sulphate~ (_Epsom salt_) (MgSO_{4}.7H_{2}O). Like the chloride,
magnesium sulphate is found rather commonly in springs and in salt
deposits. A very large deposit of the almost pure salt has been found in
Wyoming. Its name was given to it because of its abundant occurrence in
the waters of the Epsom springs in England.

Magnesium sulphate has many uses in the industries. It is used to a
small extent in the preparation of sodium and potassium sulphates, as a
coating for cotton cloth, in the dye industry, in tanning, and in the
manufacture of paints and laundry soaps. To some extent it is used in
medicine.

~Magnesium silicates.~ Many silicates containing magnesium are known and
some of them are important substances. Serpentine, asbestos, talc, and
meerschaum are examples of such substances.


ZINC

~Occurrence.~ Zinc never occurs free in nature. Its compounds have been
found in many different countries, but it is not a constituent of common
rocks and minerals, and its occurrence is rather local and confined to
definite deposits or pockets. It occurs chiefly in the following ores:

Sphalerite (zinc blende) ZnS.
Zincite ZnO.
Smithsonite ZnCO_{3}.
Willemite Zn_{2}SiO_{4}.
Franklinite ZnO.Fe_{2}O_{3}.

One fourth of the world's output of zinc comes from the United States,
Missouri being the largest producer.

~Metallurgy.~ The ores employed in the preparation of zinc are chiefly the
sulphide, oxide, and carbonate. They are first roasted in the air, by
which process they are changed into oxide:

ZnCO_{3} = ZnO + CO_{2},
ZnS + 3O = ZnO + SO_{2}.

The oxide is then mixed with coal dust, and the mixture is heated in
earthenware muffles or retorts, natural gas being used as fuel in many
cases. The oxide is reduced by this means to the metallic state, and the
zinc, being volatile at the high temperature reached, distills and is
collected in suitable receivers. At first the zinc collects in the form
of fine powder, called zinc dust or flowers of zinc, recalling the
formation under similar conditions of flowers of sulphur. Later, when
the whole apparatus has become warm, the zinc condenses to a liquid in
the receiver, from which it is drawn off into molds. Commercial zinc
often contains a number of impurities, especially carbon, arsenic, and
iron.

~Physical properties.~ Pure zinc is a rather heavy bluish-white metal with
a high luster. It melts at about 420 deg., and if heated much above this
temperature in the air takes fire and burns with a very bright bluish
flame. It boils at about 950 deg. and can therefore be purified by
distillation.

Many of the physical properties of zinc are much influenced by the
temperature and previous treatment of the metal. When cast into ingots
from the liquid state it becomes at ordinary temperatures quite hard,
brittle, and highly crystalline. At 150 deg. it is malleable and can be
rolled into thin sheets; at higher temperatures it again becomes very
brittle. When once rolled into sheets it retains its softness and
malleability at ordinary temperatures. When melted and poured into water
it forms thin brittle flakes, and in this condition is called granulated
or mossy zinc.

~Chemical properties.~ Zinc is tarnished superficially by moist air, but
beyond this is not affected by it. It does not decompose even boiling
water. When the metal is quite pure, sulphuric and hydrochloric acids
have scarcely any action upon it; when, however, it contains small
amounts of other metals such as magnesium or arsenic, or when it is
merely in contact with metallic platinum, brisk action takes place and
hydrogen is evolved. For this reason, when pure zinc is used in the
preparation of hydrogen a few drops of platinum chloride are often added
to the solution to assist the chemical action. Nitric acid dissolves the
metal readily, with the formation of zinc nitrate and various reduction
products of nitric acid. The strong alkalis act upon zinc and liberate
hydrogen:

Zn + 2KOH = Zn(OK)_{2} + 2H.

The product of this reaction, potassium zincate, is a salt of zinc
hydroxide, which is thus seen to have acid properties, though it usually
acts as a base.

~Uses of zinc.~ The metal has many familiar uses. Rolled into sheets, it
is used as a lining for vessels which are to contain water. As a thin
film upon the surface of iron (galvanized iron) it protects the iron
from rust. Iron is usually galvanized by dipping it into a bath of
melted zinc, but electrical methods are also employed. Zinc plates are
used in many forms of electrical batteries. In the laboratory zinc is
used in the preparation of hydrogen, and in the form of zinc dust as a
reducing agent.

One of the largest uses of zinc is in the manufacture of alloys. Brass,
an alloy of zinc and copper, is the most important of these; German
silver, consisting of copper, zinc, and nickel, has many uses; various
bronzes, coin metals, and bearing metals also contain zinc. Its ability
to alloy with silver finds application in the separation of silver from
lead (see silver).

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