THE MAGNESIUM FAMILY
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|SYMBOL |ATOMIC |DENSITY |MELTING |BOILING | OXIDE
| |WEIGHT | | POINT | POINT |
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Magnesium | Mg | 24.36 | 1.75 | 750° | 920° | MgO
Zinc | Zn | 65.4 | 7.00 | 420° | 950° | ZnO
Cadmium | Cd |112.4 | 8.67 | 320° | 778° | CdO
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~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°, 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° 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° 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).
~Compounds of zinc.~ In general, the compounds of zinc are similar in
formula and appearance to those of magnesium, but in other properties
they often differ markedly. A number of them have value in commercial
ways.
~Zinc oxide~ (_zinc white_) (ZnO). Zinc oxide occurs in impure form in
nature, being colored red by manganese and iron compounds. It can be
prepared just like magnesium oxide, but is more often made by burning
the metal.
Zinc oxide is a pure white powder which becomes yellow on heating and
regains its white color when cold. It is much used as a white pigment in
paints, under the name of zinc white, and has the advantage over white
lead in that it is not changed in color by sulphur compounds, while lead
turns black. It is also used in the manufacture of rubber goods.
~Commercial preparation of zinc oxide.~ Commercially it is often
made from franklinite in the following way. The franklinite is
mixed with coal and heated to a high temperature in a furnace,
by which process the zinc is set free and converted into vapor.
As the vapor leaves the furnace through a conduit it meets a
current of air and takes fire in it, forming zinc oxide. The
oxide passes on and is filtered from the air through canvas
bags, which allow the air to pass but retain the oxide. It is
thus made by burning the metal, though the metal is not
actually isolated in the process.
~Soluble salts.~ The soluble salts of zinc can be made by dissolving the
metal or the oxide in the appropriate acid. They are all somewhat
poisonous. The sulphate and chloride are the most familiar.
~Zinc sulphate~ (_white vitriol_) (ZnSO_{4}·7H_{2}O). This salt is readily
crystallized from strong solutions in transparent colorless crystals. It
is prepared commercially by careful roasting of the sulphide:
ZnS + 4O = ZnSO_{4}.
~Zinc chloride~ (ZnCl_{2}·H_{2}O). When a solution of zinc chloride is
slowly evaporated a salt of the composition ZnCl_{2}·H_{2}O crystallizes
out. If the water is completely expelled by heat and the residue
distilled, the anhydrous chloride is obtained and may be cast into
sticks or broken into lumps. In this distillation, just as in heating
magnesium chloride, some of the chloride is decomposed:
ZnCl_{2}·H_{2}O = ZnO + 2HCl.
The anhydrous chloride has a great affinity for water, and is used as a
dehydrating agent. It is also a germicide, and wood which is to be
exposed to conditions which favor decay, as, for example, railroad ties,
is often soaked in solutions of this salt.
~Insoluble compounds.~ The insoluble compounds of zinc can be prepared by
precipitation. The most important are the sulphide, carbonate, and
hydroxide.
~Zinc sulphide~ (ZnS). This substance occurs as the mineral sphalerite,
and is one of the most valued ores of zinc. Very large deposits occur in
southwestern Missouri. The natural mineral is found in large crystals or
masses, resembling resin in color and luster. When prepared by
precipitation the sulphide is white.
CADMIUM
~The element.~ This element occurs in small quantities in some zinc ores.
In the course of the metallurgy of zinc the cadmium compounds undergo
chemical changes quite similar to those of the zinc compounds, and the
cadmium distills along with the zinc. Being more volatile, it comes over
with the first of the zinc and is prepared from the first portions of
the distillate by special methods of purification. The element very
closely resembles zinc in most respects. Some of its alloys are
characterized by having low melting points.
~Compounds of cadmium.~ Among the compounds of cadmium may be mentioned
the chloride (CdCl_{2}·2H_{2}O), the sulphate (3CdSO_{4}·8H_{2}O), and
the nitrate (Cd(NO_{3})_{2}·4H_{2}O). These are white solids soluble in
water. The sulphide (CdS) is a bright yellow substance which is
insoluble in water and in dilute acids. It is valuable as a pigment in
fine paints.
EXERCISES
1. What properties have the metals of the magnesium family in common
with the alkali metals; with the alkaline-earth metals?
2. Compare the action of the metals of the magnesium group on water with
that of the other metals studied.
3. What metals already studied are prepared by electrolysis?
4. Write the equations representing the reactions between magnesium and
hydrochloric acid; between magnesium and dilute sulphuric acid.
5. What property of magnesium was taken advantage of in the isolation of
argon?
6. With phosphoric acid magnesium forms salts similar to those of
calcium. Write the names and formulas of the corresponding magnesium
salts.
7. How could you distinguish between magnesium chloride and magnesium
sulphate? between Glauber's salts and Epsom salts?
8. What weight of carnallite is necessary in the preparation of 500 g.
of magnesium?
9. Account for the fact that paints made of zinc oxide are not colored
by hydrosulphuric acid.
10. What hydroxide studied, other than zinc hydroxide, has both acid and
basic properties?
11. Write equations showing how the following compounds of zinc may be
obtained from metallic zinc: the oxide, chloride, nitrate, carbonate,
sulphate, sulphide, hydroxide.