THE METALS
~The metals.~ The elements which remain to be considered are known
collectively as the metals. They are also called the base-forming
elements, since their hydroxides are bases. A metal may therefore be
defined as an element whose hydroxide is a base. When a base dissolves
in water the hydroxyl groups form the anions, while the metallic element
forms the cations. From this standpoint a metal can be defined as an
element capable of forming simple cations in solution.
The distinction between a metal and a non-metal is not a very sharp one,
since the hydroxides of a number of elements act as bases under some
conditions and as acids under others. We have seen that antimony is an
element of this kind.
~Occurrence of metals in nature.~ A few of the metals are found in nature
in the free state. Among these are gold, platinum, and frequently
copper. They are usually found combined with other elements in the form
of oxides or salts of various acids. Silicates, carbonates, sulphides,
and sulphates are the most abundant salts. All inorganic substances
occurring in nature, whether they contain a metal or not, are called
_minerals_. Those minerals from which a useful substance can be
extracted are called _ores_ of the substance. These two terms are most
frequently used in connection with the metals.
~Extraction of metals,--metallurgy.~ The process of extracting a metal
from its ores is called the metallurgy of the metal. The metallurgy of
each metal presents peculiarities of its own, but there are several
methods of general application which are very frequently employed.
1. _Reduction of an oxide with carbon._ Many of the metals occur in
nature in the form of oxides. When these oxides are heated to a high
temperature with carbon the oxygen combines with it and the metal is set
free. Iron, for example, occurs largely in the form of the oxide
Fe_{2}O_{3}. When this is heated with carbon the reaction expressed in
the following equation takes place:
Fe_{2}O_{3} + 3 C = 2 Fe + 3 CO.
Many ores other than oxides may be changed into oxides which can then be
reduced by carbon. The conversion of such ores into oxides is generally
accomplished by heating, and this process is called _roasting_. Many
carbonates and hydroxides decompose directly into the oxide on heating.
Sulphides, on the other hand, must be heated in a current of air, the
oxygen of the air entering into the reaction. The following equations
will serve to illustrate these changes in the case of the ores of iron:
FeCO_{3} = FeO + CO_{2},
2Fe(OH)_{3} = Fe_{2}O_{3} + 3H_{2}O,
2FeS_{2} + 11O = Fe_{2}O_{3} + 4SO_{2}.
2. _Reduction of an oxide with aluminium._ Not all oxides, however, can
be reduced by carbon. In such cases aluminium may be used. Thus chromium
may be obtained in accordance with the following equation:
Cr_{2}O_{3} + 2 Al = 2 Cr + Al_{2}O_{3}.
This method is a comparatively new one, having been brought into use by
the German chemist Goldschmidt; hence it is sometimes called the
Goldschmidt method.
3. _Electrolysis._ In recent years increasing use is being made of the
electric current in the preparation of metals. In some cases the
separation of the metal from its compounds is accomplished by passing
the current through a solution of a suitable salt of the metal, the
metal usually being deposited upon the cathode. In other cases the
current is passed through a fused salt of the metal, the chloride being
best adapted to this purpose.
~Electro-chemical industries.~ Most of the electro-chemical industries of
the country are carried on where water power is abundant, since this
furnishes the cheapest means for the generation of electrical energy.
Niagara Falls is the most important locality in this country for such
industries, and many different electro-chemical products are
manufactured there. Some industries depend upon electrolytic processes,
while in others the electrical energy is used merely as a source of heat
in electric furnaces.
~Preparation of compounds of the metals.~ Since the compounds of the
metals are so numerous and varied in character, there are many ways of
preparing them. In many cases the properties of the substance to be
prepared, or the material available for its preparation, suggest a
rather unusual way. There are, however, a number of general principles
which are constantly applied in the preparation of the compounds of the
metals, and a clear understanding of them will save much time and effort
in remembering the details in any given case. The most important of
these general methods for the preparation of compounds are the
following:
1. _By direct union of two elements._ This is usually accomplished by
heating the two elements together. Thus the sulphides, chlorides, and
oxides of a metal can generally be obtained in this way. The following
equations serve as examples of this method:
Fe + S = FeS,
Mg + O = MgO,
Cu + 2Cl = CuCl_{2}.
2. _By the decomposition of a compound._ This decomposition may be
brought about either by heat alone or by the combined action of heat and
a reducing agent. Thus when the nitrate of a metal is heated the oxide
of the metal is usually obtained. Copper nitrate, for example,
decomposes as follows:
Cu(NO_{3})_{2} = CuO + 2NO_{2} + O.
Similarly the carbonates of the metals yield oxides, thus:
CaCO_{3} = CaO + CO_{2}.
Most of the hydroxides form an oxide and water when heated:
2Al(OH)_{3} = Al_{2}O_{3} + 3H_{2}O.
When heated with carbon, sulphates are reduced to sulphides, thus:
BaSO_{4} + 2C = BaS + 2CO_{2}.
3. _Methods based on equilibrium in solution._ In the preparation of
compounds the first requisite is that the reactions chosen shall be of
such a kind as will go on to completion. In the chapter on chemical
equilibrium it was shown that reactions in solution may become complete
in either of three ways: (1) a gas may be formed which escapes from
solution; (2) an insoluble solid may be formed which precipitates; (3)
two different ions may combine to form undissociated molecules. By the
judicious selection of materials these principles may be applied to the
preparation of a great variety of compounds, and illustrations of such
methods will very frequently be found in the subsequent pages.
4. _By fusion methods._ It sometimes happens that substances which are
insoluble in water and in acids, and which cannot therefore be brought
into double decomposition in the usual way, are soluble in other
liquids, and when dissolved in them can be decomposed and converted into
other desired compounds. Thus barium sulphate is not soluble in water,
and sulphuric acid, being less volatile than most other acids, cannot
easily be driven out from this salt When brought into contact with
melted sodium carbonate, however, it dissolves in it, and since barium
carbonate is insoluble in melted sodium carbonate, double decomposition
takes place:
Na_{2}CO_{3} + BaSO_{4} = BaCO_{3} + Na_{2}SO_{4}.
On dissolving the cooled mixture in water the sodium sulphate formed in
the reaction, together with any excess of sodium carbonate which may be
present, dissolves. The barium carbonate can then be filtered off and
converted into any desired salt by the processes already described.
5. _By the action of metals on salts of other metals._ When a strip of
zinc is placed in a solution of a copper salt the copper is precipitated
and an equivalent quantity of zinc passes into solution:
Zn + CuSO_{4} = Cu + ZnSO_{4}.
In like manner copper will precipitate silver from its salts:
Cu + Ag_{2}SO_{4} = 2Ag + CuSO_{4}.
It is possible to tabulate the metals in such a way that any one of them
in the table will precipitate any one following it from its salts. The
following is a list of some of the commoner metals arranged in this way:
Zinc
Iron
Tin
Lead
Copper
Bismuth
Mercury
Silver
Gold
According to this table copper will precipitate bismuth, mercury,
silver, or gold from their salts, and will in turn be precipitated by
zinc, iron, tin, or lead. Advantage is taken of this principle in the
purification of some of the metals, and occasionally in the preparation
of metals and their compounds.
~Important insoluble compounds.~ Since precipitates play so important a
part in the reactions which substances undergo, as well as in the
preparation of many chemical compounds, it is important to know what
substances are insoluble. Knowing this, we can in many cases predict
reactions under certain conditions, and are assisted in devising ways to
prepare desired compounds. While there is no general rule which will
enable one to foretell the solubility of any given compound,
nevertheless a few general statements can be made which will be of much
assistance.
1. _Hydroxides._ All hydroxides are insoluble save those of ammonium,
sodium, potassium, calcium, barium, and strontium.
2. _Nitrates._ All nitrates are soluble in water.
3. _Chlorides._ All chlorides are soluble save silver and mercurous
chlorides. (Lead chloride is but slightly soluble.)
4. _Sulphates._ All sulphates are soluble save those of barium,
strontium, and lead. (Sulphates of silver and calcium are only
moderately soluble.)
5. _Sulphides._ All sulphides are insoluble save those of ammonium,
sodium, and potassium. The sulphides of calcium, barium, strontium, and
magnesium are insoluble in water, but are changed by hydrolysis into
acid sulphides which are soluble. On this account they cannot be
prepared by precipitation.
6. _Carbonates, phosphates, and silicates._ All normal carbonates,
phosphates, and silicates are insoluble save those of ammonium, sodium
and potassium.
EXERCISES
1. Write equations representing four different ways for preparing
Cu(NO_{3})_{2}.
2. Write equations representing six different ways for preparing
ZnSO_{4}.
3. Write equations for two reactions to illustrate each of the three
ways in which reactions in solutions may become complete.
4. Give one or more methods for preparing each of the following
compounds: CaCl_{2}, PbCl_{2}, BaSO_{4}, CaCO_{3}, (NH_{4})_{2}S,
Ag_{2}S, PbO, Cu(OH)_{2} (for solubilities, see last paragraph of
chapter). State in each case the general principle involved in the
method of preparation chosen.