Courtney Ryley Cooper - "The White Desert"

Theodor Storm - "Bulemanns Haus"

Various - "Punch, or the London Charivari, Volume 104, February 4, 1893"

Hamlin Garland - "Prairie Folks"

Annual Bibliography of Commonwealth Literature 2007
This paper argues that discourses of love in Ghanaian market literature for youth offer a view into complex negotiations of agency and empowerment. Drawing on Deborah Durham's notion of youth as "social `shifters'" and Francis Nyamnjoh's conception of the "interconnectedness" of agency, I take Ghanaian market literature as one specific case of how African literature for youth foregrounds questions of continuity and change as African societies enter into increasingly complex global relations. In this literature for youth, received notions of love, often constructed out of impressions from American pop and hip hop music, carry new notions of agency that compete with existing "domesticated" forms. Authors like Ike Tandoh and Evelyn Tay employ discourses of love to offer youth alternative avenues for empowerment in a context of socio-economic disenfranchizement. In a creative process of "straddling", this writing both reveals and reproduces the contradictions that obtain in youth configurations of agency.

An Elementary Study of Chemistry

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




_Varieties of glass._ The ingredients mentioned above make a soft,
easily fusible glass. If potassium carbonate is substituted for the
sodium carbonate, the glass is much harder and less easily fused;
increasing the amount of sand has somewhat the same effect. Potassium
glass is largely used in making chemical glassware, since it resists the
action of reagents better than the softer sodium glass. If lead oxide is
substituted for the whole or a part of the lime, the glass is very soft,
but has a high index of refraction and is valuable for making optical
instruments and artificial jewels.

[Illustration: Fig. 75]

_Coloring of glass._ Various substances fused along with the glass
mixture give characteristic colors. The amber color of common bottles is
due to iron compounds in the glass; in other cases iron colors the glass
green. Cobalt compounds color it deep blue; those of manganese give it
an amethyst tint and uranium compounds impart a peculiar yellowish green
color. Since iron is nearly always present in the ingredients, glass is
usually slightly yellow. This color can be removed by adding the proper
amount of manganese dioxide, for the amethyst color of manganese and the
yellow of iron together produce white light.

_Nature of glass._ Glass is not a definite chemical compound and its
composition varies between wide limits. Fused glass is really a solution
of various silicates, such as those of calcium and lead, in fused sodium
or potassium silicate. A certain amount of silicon dioxide is also
present. This solution is then allowed to solidify under such conditions
of cooling that the dissolved substances do not separate from the
solvent. The compounds which are used to color the glass are sometimes
converted into silicates, which then dissolve in the glass, giving it a
uniform color. In other cases, as in the milky glasses which resemble
porcelain in appearance, the color or opaqueness is due to the finely
divided color material evenly distributed throughout the glass, but not
dissolved in it. Milky glass is made by mixing calcium fluoride, tin
oxide, or some other insoluble substance in the melted glass. Copper or
gold in metallic form scattered through glass gives it shades of red.


TITANIUM

Titanium is a very widely distributed element in nature, being
found in almost all soils, in many rocks, and even in plant and
animal tissues. It is not very abundant in any one locality,
and it possesses little commercial value save in connection
with the iron industry. Its most common ore is rutile
(TiO_{2}), which resembles silica in many respects.

In both physical and chemical properties titanium resembles
silicon, though it is somewhat more metallic in character. This
resemblance is most marked in the acids of titanium. It not
only forms metatitanic and orthotitanic acids but a great
variety of polytitanic acids as well.


BORON

~Occurrence.~ Boron is never found free in nature. It occurs as boric acid
(H_{3}BO_{3}), and in salts of polyboric acids, which usually have very
complicated formulas.

~Preparation and properties.~ Boron can be prepared from its oxide by
reduction with magnesium, exactly as in the case of silicon. It
resembles silicon very strikingly in its properties. It occurs in
several allotropic forms, is very hard when crystallized, and is rather
inactive toward reagents. It forms a hydride, BH_{3}, and combines
directly with the elements of the chlorine family. Boron fluoride
(BF_{3}) is very similar to silicon fluoride in its mode of formation
and chemical properties.

~Boric oxide~ (B_{2}O_{3}). Boron forms one well-known oxide, B_{2}O_{3},
called boric anhydride. It is formed as a glassy mass by heating boric
acid to a high temperature. It absorbs water very readily, uniting with
it to form boric acid again:

B_{2}O_{3} + 3H_{2}O = 2H_{3}BO_{3}.

In this respect it differs from silicon dioxide, which will not combine
directly with water.

~Boric acid~ (H_{3}BO_{3}). This is found in nature in considerable
quantities and forms one of the chief sources of boron compounds. It is
found dissolved in the water of hot springs in some localities,
particularly in Italy. Being volatile with steam, the vapor which
escapes from these springs has some boric acid in it. It is easily
obtained from these sources by condensation and evaporation, the
necessary heat being supplied by other hot springs.

Boric acid crystallizes in pearly flakes, which are greasy to the touch.
In the laboratory it is easily prepared by treating a strong, hot
solution of borax with sulphuric acid. Boric acid being sparingly
soluble in water crystallizes out on cooling:

Na_{2}B_{4}O_{7} + 5H_{2}O + H_{2}SO_{4} = Na_{2}SO_{4} + 4H_{3}BO_{3}.

The substance is a mild antiseptic, and on this account is often used in
medicine and as a preservative for canned foods and milk.

~Metaboric and polyboric acids.~ When boric acid is gently heated it is
converted into metaboric acid (HBO_{2}):

H_{3}BO_{3} = HBO_{2} + H_{2}O.

On heating metaboric acid to a somewhat higher temperature tetraboric
acid (H_{2}B_{4}O_{7}) is formed:

4HBO_{2} = H_{2}B_{4}O_{7} + H_{2}O.

Many other complex acids of boron are known.

~Borax.~ Borax is the sodium salt of tetraboric acid, having the formula
Na_{2}B_{4}O_{7}.10 H_{2}O. It is found in some arid countries, as
southern California and Tibet, but is now made commercially from the
mineral colemanite, which is the calcium salt of a complex boric acid.
When this is treated with a solution of sodium carbonate, calcium
carbonate is precipitated and borax crystallizes from the solution.

When heated borax at first swells up greatly, owing to the expulsion of
the water of crystallization, and then melts to a clear glass. This
glass has the property of easily dissolving many metallic oxides, and on
this account borax is used as a flux in soldering, for the purpose of
removing from the metallic surfaces to be soldered the film of oxide
with which they are likely to be covered. These oxides often give a
characteristic color to the clear borax glass, and borax beads are
therefore often used in testing for the presence of metals, instead of
the metaphosphoric acid bead already described.

The reason that metallic oxides dissolve in borax is that borax
contains an excess of acid anhydride, as can be more easily
seen if its formula is written 2NaBO_{2} + B_{2}O_{3}. The
metallic oxide combines with this excess of acid anhydride,
forming a mixed salt of metaboric acid.

Borax is extensively used as a constituent of enamels and glazes for
both metal ware and pottery. It is also used as a flux in soldering and
brazing, and in domestic ways it serves as a mild alkali, as a
preservative for meats, and in a great variety of less important
applications.


EXERCISES

1. Account for the fact that a solution of borax in water is alkaline.

2. What weight of water of crystallization does 1 kg. of borax contain?

3. When a concentrated solution of borax acts on silver nitrate a borate
of silver is formed. If the solution of borax is dilute, however, an
hydroxide of silver forms. Account for this difference in behavior.




CHAPTER XXII

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.




CHAPTER XXIII

THE ALKALI METALS


=================================================================
| | | | |
| SYMBOL | ATOMIC | DENSITY | MELTING | FIRST PREPARED
| | WEIGHT | | POINT |
__________|________|________|_________|_________|________________
| | | | |
Lithium | Li | 7.03 | 0.59 | 186. deg. | Davy 1820
Sodium | Na | 23.05 | 0.97 | 97.6 deg. | " 1807
Potassium | K | 39.15 | 0.87 | 62.5 deg. | " 1807
Rubidium | Rb | 85.5 | 1.52 | 38.5 deg. | Bunsen 1861
Caesium | Cs | 132.9 | 1.88 | 26.5 deg. | " 1860
=================================================================

~The family.~ The metals listed in the above table constitute the even
family in Group I in the periodic arrangement of the elements, and
therefore form a natural family. The name alkali metals is commonly
applied to the family for the reason that the hydroxides of the most
familiar members of the family, namely sodium and potassium, have long
been called alkalis.

1. _Occurrence._ While none of these metals occur free in nature, their
compounds are very widely distributed, being especially abundant in sea
and mineral waters, in salt beds, and in many rocks. Only sodium and
potassium occur in abundance, the others being rarely found in any
considerable quantity.

2. _Preparation._ The metals are most conveniently prepared by the
electrolysis of their fused hydroxides or chlorides, though it is
possible to prepare them by reducing their oxides or carbonates with
carbon.

3. _Properties._ They are soft, light metals, having low melting points
and small densities, as is indicated in the table. Their melting points
vary inversely with their atomic weights, while their densities (sodium
excepted) vary directly with these. The pure metals have a silvery
luster but tarnish at once when exposed to the air, owing to the
formation of a film of oxide upon the surface of the metal. They are
therefore preserved in some liquid, such as coal oil, which contains no
oxygen. Because of their strong affinity for oxygen they decompose water
with great ease, forming hydroxides and liberating hydrogen in
accordance with the equation

M + H_{2}O = MOH + H,

where M stands for any one of these metals. These hydroxides are white
solids; they are readily soluble in water and possess very strong basic
properties. These bases are nearly equal in strength, that is, they all
dissociate in water to about the same extent.

4. _Compounds._ The alkali metals almost always act as univalent
elements in the formation of compounds, the composition of which can be
represented by such formulas as MH, MCl, MNO_{3}, M_{2}SO_{4},
M_{3}PO_{4}. These compounds, when dissolved in water, dissociate in
such a way as to form simple, univalent metallic ions which are
colorless. With the exception of lithium these metals form very few
insoluble compounds, so that it is not often that precipitates
containing them are obtained. Only sodium and potassium will be studied
in detail, since the other metals of the family are of relatively small
importance.

The compounds of sodium and potassium are so similar in properties that
they can be used interchangeably for most purposes. Other things being
equal, the sodium compounds are prepared in preference to those of
potassium, since they are cheaper. When a given sodium compound is
deliquescent, or is so soluble that it is difficult to purify, the
corresponding potassium compound is prepared in its stead, provided its
properties are more desirable in these respects.


SODIUM

~Occurrence in nature.~ Large deposits of sodium chloride have been found
in various parts of the world, and the water of the ocean and of many
lakes and springs contains notable quantities of it. The element also
occurs as a constituent of many rocks and is therefore present in the
soil formed by their disintegration. The mineral cryolite
(Na_{3}AlF_{6}) is an important substance, and the nitrate, carbonate,
and borate also occur in nature.

~Preparation.~ In 1807 Sir Humphry Davy succeeded in preparing very small
quantities of metallic sodium by the electrolysis of the fused
hydroxide. On account of the cost of electrical energy it was for many
years found more economical to prepare it by reducing the carbonate with
carbon in accordance with the following equation:

Na_{2}CO_{3} + 2C = 2Na + 3CO.

The cost of generating the electric current has been diminished to such
an extent, however, that it is now more economical to prepare sodium by
Davy's original method, namely, by the electrolysis of the fused
hydroxide or chloride. When the chloride is used the process is
difficult to manage, owing to the higher temperature required to keep
the electrolyte fused, and because of the corroding action of the fused
chloride upon the containing vessel.

[Illustration: SIR HUMPHRY DAVY (English) (1778-1829)

Isolated sodium, lithium, potassium, barium, strontium, and calcium by
means of electrolysis; demonstrated the elementary nature of chlorine;
invented the safety lamp; discovered the stupefying effects of nitrous
oxide]

~Technical preparation.~ The sodium hydroxide is melted in a
cylindrical iron vessel (Fig. 76) through the bottom of which
rises the cathode K. The anodes A, several in number, are
suspended around the cathode from above. A cylindrical vessel
C floats in the fused alkali directly over the cathode, and
under this cap the sodium and hydrogen liberated at the cathode
collect. The hydrogen escapes by lifting the cover, and the
sodium, protected from the air by the hydrogen, is skimmed or
drained off from time to time. Oxygen is set free upon the
anode and escapes into the air through the openings O without
coming into contact with the sodium or hydrogen. This process
is carried on extensively at Niagara Falls.

[Illustration: Fig. 76]

~Properties.~ Sodium is a silver-white metal about as heavy as water, and
so soft that it can be molded easily by the fingers or pressed into
wire. It is very active chemically, combining with most of the
non-metallic elements, such as oxygen and chlorine, with great energy.
It will often withdraw these elements from combination with other
elements, and is thus able to decompose water and the oxides and
chlorides of many metals.

~Sodium peroxide~ (NaO). Since sodium is a univalent element we should
expect it to form an oxide of the formula Na_{2}O. While such an oxide
can be prepared, the peroxide (NaO) is much better known. It is a
yellowish-white powder made by burning sodium in air. Its chief use is
as an oxidizing agent. When heated with oxidizable substances it gives
up a part of its oxygen, as shown in the equation

2NaO = Na_{2}O + O.

Water decomposes it in accordance with the equation

2NaO + 2H_{2}O = 2NaOH + H_{2}O_{2}.

Acids act readily upon it, forming a sodium salt and hydrogen peroxide:

2NaO + 2HCl = 2NaCl + H_{2}O_{2}.

In these last two reactions the hydrogen dioxide formed may decompose
into water and oxygen if the temperature is allowed to rise:

H_{2}O_{2} = H_{2}O + O.

~Peroxides.~ It will be remembered that barium dioxide (BaO_{2})
yields hydrogen dioxide when treated with acids, and that
manganese dioxide gives up oxygen when heated with sulphuric
acid. Oxides which yield either hydrogen dioxide or oxygen when
treated with water or an acid are called peroxides.

Copyright (c) 2007. topboookz.com. All rights reserved.