Oscar Wilde - "The Picture of Dorian Gray"

J.G. Adams - "Hymns for Christian Devotion"

Various - "Scientific American Supplement, No. 358, November 11, 1882"

Susanna Moodie - "The Monctons: A Novel, Volume I"

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




The conditions which prevail in the laboratory are never the standard
conditions. It becomes necessary, therefore, to find a way to calculate
the volume which a gas will occupy under standard conditions from the
volume which it occupies under any other conditions. This may be done in
accordance with the following laws.

~Law of Charles.~ This law expresses the effect which a change in the
temperature of a gas has upon its volume. It may be stated as follows:
_For every degree the temperature of a gas rises above zero the volume
of the gas is increased by 1/273 of the volume which it occupies at
zero; likewise for every degree the temperature of the gas falls below
zero the volume of the gas is decreased by 1/273 of the volume which it
occupies at zero, provided in both cases that the pressure to which the
gas is subjected remains constant._

If V represents the volume of gas at 0 deg., then the volume at 1 deg. will be
V + 1/273 V; at 2 deg. it will be V + 2/273 V; or, in general, the
volume v, at the temperature t, will be expressed by the formula

(1) v = V + t/273 V,

or (2) v = V(1 + (t/273)).

Since 1/273 = 0.00366, the formula may be written

(3) v = V(1 + 0.00366t).

Since the value of V (volume under standard conditions) is the one
usually sought, it is convenient to transpose the equation to the
following form:

(4) V = v/(1 + 0.00366t).

The following problem will serve as an illustration of the application
of this equation.

The volume of a gas at 20 deg. is 750 cc.; find the volume it will occupy at
0 deg., the pressure remaining constant.

In this case, v = 750 cc. and t = 20. By substituting these values,
equation (4) becomes

V = 750/(1 + 0.00366 x 20) = 698.9 cc.

~Law of Boyle.~ This law expresses the relation between the volume
occupied by a gas and the pressure to which it is subjected. It may be
stated as follows: _The volume of a gas is inversely proportional to the
pressure under which it is measured, provided the temperature of the gas
remains constant._

If V represents the volume when subjected to a pressure P and v
represents its volume when the pressure is changed to p, then, in
accordance with the above law, V : v :: p : P, or VP = vp.
In other words, for a given weight of a gas the product of the numbers
representing its volume and the pressure to which it is subjected is a
constant.

Since the pressure of the atmosphere at any point is indicated by the
barometric reading, it is convenient in the solution of the problems to
substitute the latter for the pressure measured in grams per square
centimeter. The average reading of the barometer at the sea level is 760
mm., which corresponds to a pressure of 1033.3 g. per square centimeter.
The following problem will serve as an illustration of the application
of Boyle's law.

A gas occupies a volume of 500 cc. in a laboratory where the barometric
reading is 740 mm. What volume would it occupy if the atmospheric
pressure changed so that the reading became 750 mm.?

Substituting the values in the equation VP = vp, we have 500 x 740 =
v x 750, or v = 493.3 cc.

~Variations in the volume of a gas due to changes both in temperature and
pressure.~ Inasmuch as corrections must be made as a rule for both
temperature and pressure, it is convenient to combine the equations
given above for the corrections for each, so that the two corrections
may be made in one operation. The following equation is thus obtained:

(5) V_{s} = vp/(760(1 + 0.00366t)),

in which V_{s} represents the volume of a gas under standard
conditions and v, p, and t the volume, pressure, and temperature
respectively at which the gas was actually measured.

The following problem will serve to illustrate the application of this
equation.

A gas having a temperature of 20 deg. occupies a volume of 500 cc. when
subjected to a pressure indicated by a barometric reading of 740 mm.
What volume would this gas occupy under standard conditions?

In this problem v = 500, p = 740, and t = 20. Substituting these
values in the above equation, we get

V_{s} = (500 x 740)/(760 (1 + 0.00366 x 20)) = 453.6 cc.

[Illustration: Fig. 8]

~Variations in the volume of a gas due to the pressure of aqueous vapor.~
In many cases gases are collected over water, as explained under the
preparation of oxygen. In such cases there is present in the gas a
certain amount of water vapor. This vapor exerts a definite pressure,
which acts in opposition to the atmospheric pressure and which therefore
must be subtracted from the latter in determining the effective pressure
upon the gas. Thus, suppose we wish to determine the pressure to which
the gas in tube A (Fig. 8) is subjected. The tube is raised or lowered
until the level of the water inside and outside the tube is the same.
The atmosphere presses down upon the surface of the water (as indicated
by the arrows), thus forcing the water upward within the tube with a
pressure equal to the atmospheric pressure. The full force of this
upward pressure, however, is not spent in compressing the gas within the
tube, for since it is collected over water it contains a certain amount
of water vapor. This water vapor exerts a pressure (as indicated by the
arrow within the tube) in opposition to the upward pressure. It is
plain, therefore, that the effective pressure upon the gas is equal to
the atmospheric pressure less the pressure exerted by the aqueous vapor.
The pressure exerted by the aqueous vapor increases with the
temperature. The figures representing the extent of this pressure (often
called the _tension of aqueous vapor_) are given in the Appendix. They
express the pressure or tension in millimeters of mercury, just as the
atmospheric pressure is expressed in millimeters of mercury.
Representing the pressure of the aqueous vapor by a, formula (5)
becomes

(6) V_{s} = v(p - a)/(760(1 + 0.00366t)).

The following problem will serve to illustrate the method of applying
the correction for the pressure of the aqueous vapor.

The volume of a gas measured over water in a laboratory where the
temperature is 20 deg. and the barometric reading is 740 mm. is 500 cc. What
volume would this occupy under standard conditions?

The pressure exerted by the aqueous vapor at 20 deg. (see table in Appendix)
is equal to the pressure exerted by a column of mercury 17.4 mm. in
height. Substituting the values of v, t, p, and a in formula
(6), we have

(6) V_{s} = 500(740 - 17.4)/(760(1 + 0.00366 x 20)) = 442.9 cc.

~Adjustment of tubes before reading gas volumes.~ In measuring the volumes
of gases collected in graduated tubes or other receivers, over a liquid
as illustrated in Fig. 8, the reading should be taken after raising or
lowering the tube containing the gas until the level of the liquid
inside and outside the tube is the same; for it is only under these
conditions that the upward pressure within the tube is the same as the
atmospheric pressure.


EXERCISES

1. What is the meaning of the following words? phlogiston, ozone,
phosphorus. (Consult dictionary.)

2. Can combustion take place without the emission of light?

3. Is the evolution of light always produced by combustion?

4. (a) What weight of oxygen can be obtained from 100 g. of water?
(b) What volume would this occupy under standard conditions?

5. (a) What weight of oxygen can be obtained from 500g. of mercuric
oxide? (b) What volume would this occupy under standard conditions?

6. What weight of each of the following compounds is necessary to
prepare 50 l. of oxygen? (a) water; (b) mercuric oxide; (c)
potassium chlorate.

7. Reduce the following volumes to 0 deg., the pressure remaining constant:
(a) 150 cc. at 10 deg.; (b) 840 cc. at 273 deg..

8. A certain volume of gas is measured when the temperature is 20 deg.. At
what temperature will its volume be doubled?

9. Reduce the following volumes to standard conditions of pressure, the
temperature remaining constant: (a) 200 cc. at 740 mm.; (b) 500 l.
at 380 mm.

10. What is the weight of 1 l. of oxygen when the pressure is 750 mm.
and the temperature 0 deg.?

11. Reduce the following volumes to standard conditions of temperature
and pressure: (a) 340 cc. at 12 deg. and 753 mm; (b) 500 cc. at 15 deg. and
740 mm.

12. What weight of potassium chlorate is necessary to prepare 250 l. of
oxygen at 20 deg. and 750 mm.?

13. Assuming the cost of potassium chlorate and mercuric oxide to be
respectively $0.50 and $1.50 per kilogram, calculate the cost of
materials necessary for the preparation of 50 l. of oxygen from each of
the above compounds.

14. 100 g. of potassium chlorate and 25 g. of manganese dioxide were
heated in the preparation of oxygen. What products were left in the
flask, and how much of each was present?




CHAPTER III

HYDROGEN


~Historical.~ The element hydrogen was first clearly recognized as a
distinct substance by the English investigator Cavendish, who in 1766
obtained it in a pure state, and showed it to be different from the
other inflammable airs or gases which had long been known. Lavoisier
gave it the name hydrogen, signifying water former, since it had been
found to be a constituent of water.

~Occurrence.~ In the free state hydrogen is found in the atmosphere, but
only in traces. In the combined state it is widely distributed, being a
constituent of water as well as of all living organisms, and the
products derived from them, such as starch and sugar. About 10% of the
human body is hydrogen. Combined with carbon, it forms the substances
which constitute petroleum and natural gas.

It is an interesting fact that while hydrogen in the free state
occurs only in traces on the earth, it occurs in enormous
quantities in the gaseous matter surrounding the sun and
certain other stars.

~Preparation from water.~ Hydrogen can be prepared from water by several
methods, the most important of which are the following.

1. _By the electric current._ As has been indicated in the preparation
of oxygen, water is easily separated into its constituents, hydrogen and
oxygen, by passing an electric current through it under certain
conditions.

2. _By the action of certain metals._ When brought into contact with
certain metals under appropriate conditions, water gives up a portion
or the whole of its hydrogen, its place being taken by the metal. In the
case of a few of the metals this change occurs at ordinary temperatures.
Thus, if a bit of sodium is thrown on water, an action is seen to take
place at once, sufficient heat being generated to melt the sodium, which
runs about on the surface of the water. The change which takes place
consists in the displacement of one half of the hydrogen of the water by
the sodium, and may be represented as follows:

_ _ _ _
| hydrogen | | sodium |
sodium + | hydrogen |(water) = | hydrogen |(sodium hydroxide) + hydrogen
|_oxygen _| |_oxygen _|

The sodium hydroxide formed is a white solid which remains dissolved in
the undecomposed water, and may be obtained by evaporating the solution
to dryness. The hydrogen is evolved as a gas and may be collected by
suitable apparatus.

Other metals, such as magnesium and iron, decompose water rapidly, but
only at higher temperatures. When steam is passed over hot iron, for
example, the iron combines with the oxygen of the steam, thus displacing
the hydrogen. Experiments show that the change may be represented as
follows:

_ _
| hydrogen | _ _ _ _
iron + | hydrogen |(water) = | iron |(iron oxide) + | hydrogen |
|_oxygen _| |_oxygen _| |_hydrogen_|

The iron oxide formed is a reddish-black compound, identical with that
obtained by the combustion of iron in oxygen.

~Directions for preparing hydrogen by the action of steam on
iron.~ The apparatus used in the preparation of hydrogen from
iron and steam is shown in Fig. 9. A porcelain or iron tube
B, about 50 cm. in length and 2 cm. or 3 cm. in diameter, is
partially filled with fine iron wire or tacks and connected as
shown in the figure. The tube B is heated, slowly at first,
until the iron is red-hot. Steam is then conducted through the
tube by boiling the water in the flask A. The hot iron
combines with the oxygen in the steam, setting free the
hydrogen, which is collected over water. The gas which first
passes over is mixed with the air previously contained in the
flask and tube, and is allowed to escape, _since a mixture of
hydrogen with oxygen or air explodes violently when brought in
contact with a flame_. It is evident that the flask A must be
disconnected from the tube before the heat is withdrawn.

That the gas obtained is different from air and oxygen may be
shown by holding a bottle of it mouth downward and bringing a
lighted splint into it. The hydrogen is ignited and burns with
an almost colorless flame.

[Illustration Fig. 9]

~Preparation from acids~ (_usual laboratory method_). While hydrogen can
be prepared from water, either by the action of the electric current or
by the action of certain metals, these methods are not economical and
are therefore but little used. In the laboratory hydrogen is generally
prepared from compounds known as acids, all of which contain hydrogen.
When acids are brought in contact with certain metals, the metals
dissolve and set free the hydrogen of the acid. Although this reaction
is a quite general one, it has been found most convenient in preparing
hydrogen by this method to use either zinc or iron as the metal and
either hydrochloric or sulphuric acid as the acid. Hydrochloric acid is
a compound consisting of 2.77% hydrogen and 97.23% chlorine, while
sulphuric acid consists of 2.05% hydrogen, 32.70% sulphur, and 65.25%
oxygen.

The changes which take place in the preparation of hydrogen from zinc
and sulphuric acid (diluted with water) may be represented as follows:

_ _ _ _
| hydrogen |(sulphuric | zinc |(zinc
zinc + | sulphur | acid) = | sulphur | sulphate) + hydrogen
|_oxygen _| |_oxygen _|

In other words, the zinc has taken the place of the hydrogen in
sulphuric acid. The resulting compound contains zinc, sulphur, and
oxygen, and is known as zinc sulphate. This remains dissolved in the
water present in the acid. It may be obtained in the form of a white
solid by evaporating the liquid left after the metal has passed into
solution.

When zinc and hydrochloric acid are used the following changes take
place:

_ _ _ _
| hydrogen |(hydrochloric | zinc |(zinc
zinc + |_chlorine_| acid) = |_chlorine_| chloride) + hydrogen

When iron is used the changes which take place are exactly similar to
those just given for zinc.

[Illustration Fig. 10.]

~Directions for preparing hydrogen from acids.~ The preparation
of hydrogen from acids is carried out in the laboratory as
follows: The metal is placed in a flask or wide-mouthed bottle
A (Fig. 10) and the acid is added slowly through the funnel
tube B. The metal dissolves in the acid, while the hydrogen
which is liberated escapes through the exit tube C and is
collected over water. It is evident that the hydrogen which
passes over first is mixed with the air from the bottle A.
Hence care must be taken not to bring a flame near the exit
tube, since, as has been stated previously, such a mixture
explodes with great violence when brought in contact with a
flame.

~Precautions.~ Both sulphuric acid and zinc, if impure, are
likely to contain small amounts of arsenic. Such materials
should not be used in preparing hydrogen, since the arsenic
present combines with a portion of the hydrogen to form a very
poisonous gas known as arsine. On the other hand, chemically
pure sulphuric acid, i.e. sulphuric acid that is entirely free
from impurities, will not act upon chemically pure zinc. The
reaction may be started, however, by the addition of a few
drops of a solution of copper sulphate or platinum
tetrachloride.

~Physical properties.~ Hydrogen is similar to oxygen in that it is a
colorless, tasteless, odorless gas. It is characterized by its extreme
lightness, being the lightest of all known substances. One liter of the
gas weighs only 0.08984 g. On comparing this weight with that of an
equal volume of oxygen, viz., 1.4285 g., the latter is found to be 15.88
times as heavy as hydrogen. Similarly, air is found to be 14.38 times as
heavy as hydrogen. Soap bubbles blown with hydrogen rapidly rise in the
air. On account of its lightness it is possible to pour it upward from
one bottle into another. Thus, if the bottle A (Fig. 11) is filled
with hydrogen, placed mouth downward by the side of bottle _B_, filled
with air, and is then gradually inverted under B as indicated in the
figure, the hydrogen will flow upward into bottle _B_, displacing the
air. Its presence in bottle B may then be shown by bringing a lighted
splint to the mouth of the bottle, when the hydrogen will be ignited by
the flame. It is evident, from this experiment, that in order to retain
the gas in an open bottle the bottle must be placed mouth downward.

[Illustration Fig. 11]

Hydrogen is far more difficult to liquefy than any other gas, with the
exception of helium, a rare element recently found to exist in the
atmosphere. The English scientist Dewar, however, in 1898 succeeded not
only in obtaining hydrogen in liquid state but also as a solid. Liquid
hydrogen is colorless and has a density of only 0.07. Its boiling point
under atmospheric pressure is -252 deg.. Under diminished pressure the
temperature has been reduced to -262 deg.. The solubility of hydrogen in
water is very slight, being still less than that of oxygen.

Pure hydrogen produces no injurious results when inhaled. Of course one
could not live in an atmosphere of the gas, since oxygen is essential to
respiration.

~Chemical properties.~ At ordinary temperatures hydrogen is not an active
element. A mixture of hydrogen and chlorine, however, will combine with
explosive violence at ordinary temperature if exposed to the sunlight.
The union can be brought about also by heating. The product formed in
either case is hydrochloric acid. Under suitable conditions hydrogen
combines with nitrogen to form ammonia, and with sulphur to form the
foul-smelling gas, hydrogen sulphide. The affinity of hydrogen for
oxygen is so great that a mixture of hydrogen and oxygen or hydrogen
and air explodes with great violence when heated to the kindling
temperature (about 612 deg.). Nevertheless under proper conditions hydrogen
may be made to burn quietly in either oxygen or air. The resulting
hydrogen flame is almost colorless and is very hot. The combustion of
the hydrogen is, of course, due to its union with oxygen. The product of
the combustion is therefore a compound of hydrogen and oxygen. That this
compound is water may be shown easily by experiment.

[Illustration Fig. 12]

~Directions for burning hydrogen in air.~ The combustion of
hydrogen in air may be carried out safely as follows: The
hydrogen is generated in the bottle A (Fig. 12), is dried by
conducting it through the tube X, filled with some substance
(generally calcium chloride) which has a great attraction for
moisture, and escapes through the tube T, the end of which is
drawn out to a jet. The hydrogen first liberated mixes with the
air contained in the generator. If a flame is brought near the
jet before this mixture has all escaped, a violent and very
dangerous explosion results, since the entire apparatus is
filled with the explosive mixture. On the other hand, if the
flame is not applied until all the air has been expelled, the
hydrogen is ignited and burns quietly, since only the small
amount of it which escapes from the jet can come in contact
with the oxygen of the air at any one time. By holding a cold,
dry bell jar or bottle over the flame, in the manner shown in
the figure, the steam formed by the combustion of the hydrogen
is condensed, the water collecting in drops on the sides of the
jar.

~Precautions.~ In order to avoid danger it is absolutely necessary to
prove that the hydrogen is free from air before igniting it. This can be
done by testing small amounts of the escaping gas. A convenient and safe
method of doing this is to fill a test tube with the gas by inverting it
over the jet. The hydrogen, on account of its lightness, collects in the
tube, displacing the air. After holding it over the jet for a few
moments in order that it may be filled with the gas, the tube is gently
brought, mouth downward, to the flame of a burner placed not nearer than
an arm's length from the jet. If the hydrogen is mixed with air a slight
explosion occurs, but if pure it burns quietly in the tube. The
operation is repeated until the gas burns quietly, when the tube is
quickly brought back over the jet for an instant, whereby the escaping
hydrogen is ignited by the flame in the tube.

[Illustration. Fig. 13]

~A mixture of hydrogen and oxygen is explosive.~ That a mixture of
hydrogen and air is explosive may be shown safely as follows: A cork
through which passes a short glass tube about 1 cm. in diameter is
fitted air-tight into the tubule of a bell jar of 2 l. or 3 l. capacity.
(A thick glass bottle with bottom removed may be used.) The tube is
closed with a small rubber stopper and the bell jar filled with
hydrogen, the gas being collected over water. When entirely filled with
the gas the jar is removed from the water and supported by blocks of
wood in order to leave the bottom of the jar open, as shown in Fig. 13.
The stopper is now removed from the tube in the cork, and the hydrogen,
which on account of its lightness escapes from the tube, is at once
lighted. As the hydrogen escapes, the air flows in at the bottom of the
jar and mixes with the remaining portion of the hydrogen, so that a
mixture of the two soon forms, and a loud explosion results. The
explosion is not dangerous, since the bottom of the jar is open, thus
leaving room for the expansion of the hot gas.

Since air is only one fifth oxygen, the remainder being inert gases, it
may readily be inferred that a mixture of hydrogen with pure oxygen
would be far more explosive than a mixture of hydrogen with air. Such
mixtures should not be made except in small quantities and by
experienced workers.

~Hydrogen does not support combustion.~ While hydrogen is readily
combustible, it is not a supporter of combustion. In other words,
substances will not burn in it. This may be shown by bringing a lighted
candle supported by a stiff wire into a bottle or cylinder of the pure
gas, as shown in Fig. 14. The hydrogen is ignited by the flame of the
candle and burns at the mouth of the bottle, where it comes in contact
with the oxygen in the air. When the candle is thrust up into the gas,
its flame is extinguished on account of the absence of oxygen. If slowly
withdrawn, the candle is relighted as it passes through the layer of
burning hydrogen.

[Illustration: Fig. 14]

[Illustration: Fig. 15]

~Reduction.~ On account of its great affinity for oxygen, hydrogen has the
power of abstracting it from many of its compounds. Thus, if a stream of
hydrogen, dried by passing through the tube B (Fig. 15), filled with
calcium chloride, is conducted through the tube C containing some
copper oxide, heated to a moderate temperature, the hydrogen abstracts
the oxygen from the copper oxide. The change may be represented as
follows:

hydrogen + {copper} {hydrogen}
{oxygen}(copper oxide) = {oxygen }(water) + copper

The water formed collects in the cold portions of the tube C near its
end. In this experiment the copper oxide is said to undergo reduction.
_Reduction may therefore be defined as the process of withdrawing oxygen
from a compound._

~Relation of reduction to oxidation.~ At the same time that the copper
oxide is reduced it is clear that the hydrogen is oxidized, for it
combines with the oxygen given up by the copper oxide. The two processes
are therefore very closely related, and it usually happens that when one
substance is oxidized some other substance is reduced. That substance
which gives up its oxygen is called an _oxidizing agent_, while the
substance which unites with the oxygen is called a _reducing agent_.

~The oxyhydrogen blowpipe.~ This is a form of apparatus used for burning
hydrogen in pure oxygen. As has been previously stated, the flame
produced by the combustion of hydrogen in the air is very hot. It is
evident that if pure oxygen is substituted for air, the temperature
reached will be much higher, since there are no inert gases to absorb
the heat. The oxyhydrogen blowpipe, used to effect this combination,
consists of a small tube placed within a larger one, as shown in Fig.
16.

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