OXYGEN
~History.~ The discovery of oxygen is generally attributed to the English
chemist Priestley, who in 1774 obtained the element by heating a
compound of mercury and oxygen, known as red oxide of mercury. It is
probable, however, that the Swedish chemist Scheele had previously
obtained it, although an account of his experiments was not published
until 1777. The name oxygen signifies acid former. It was given to the
element by the French chemist Lavoisier, since he believed that all
acids owe their characteristic properties to the presence of oxygen.
This view we now know to be incorrect.
~Occurrence.~ Oxygen is by far the most abundant of all the elements. It
occurs both in the free and in the combined state. In the free state it
occurs in the air, 100 volumes of dry air containing about 21 volumes of
oxygen. In the combined state it forms eight ninths of water and nearly
one half of the rocks composing the earth's crust. It is also an
important constituent of the compounds which compose plant and animal
tissues; for example, about 66% by weight of the human body is oxygen.
~Preparation.~ Although oxygen occurs in the free state in the atmosphere,
its separation from the nitrogen and other gases with which it is mixed
is such a difficult matter that in the laboratory it has been found more
convenient to prepare it from its compounds. The most important of the
laboratory methods are the following:
1. _Preparation from water._ Water is a compound, consisting of 11.18%
hydrogen and 88.82% oxygen. It is easily separated into these
constituents by passing an electric current through it under suitable
conditions. The process will be described in the chapter on water. While
this method of preparation is a simple one, it is not economical.
2. _Preparation from mercuric oxide._ This method is of interest, since
it is the one which led to the discovery of oxygen. The oxide, which
consists of 7.4% oxygen and 92.6% mercury, is placed in a small, glass
test tube and heated. The compound is in this way decomposed into
mercury which collects on the sides of the glass tube, forming a silvery
mirror, and oxygen which, being a gas, escapes from the tube. The
presence of the oxygen is shown by lighting the end of a splint,
extinguishing the flame and bringing the glowing coal into the mouth of
the tube. The oxygen causes the glowing coal to burst into a flame.
In a similar way oxygen may be obtained from its compounds with
some of the other elements. Thus manganese dioxide, a black
compound of manganese and oxygen, when heated to about 700°,
loses one third of its oxygen, while barium dioxide, when
heated, loses one half of its oxygen.
3. _Preparation from potassium chlorate (usual laboratory method)._
Potassium chlorate is a white solid which consists of 31.9% potassium,
28.9% chlorine, and 39.2% oxygen. When heated it undergoes a series of
changes in which all the oxygen is finally set free, leaving a compound
of potassium and chlorine called potassium chloride. The change may be
represented as follows:
/potassium\
| | (potassium / potassium \ (potassium
{ chlorine } = { } + oxygen
| | chlorate) \ chlorine / chloride)
\oxygen /
[Illustration: JOSEPH PRIESTLEY (English) (1733-1804)
School-teacher, theologian, philosopher, scientist; friend of Benjamin
Franklin; discoverer of oxygen; defender of the phlogiston theory; the
first to use mercury in a pneumatic trough, by which means he first
isolated in gaseous form hydrochloric acid, sulphur dioxide, and
ammonia]
The evolution of the oxygen begins at about 400°. It has been found,
however, that if the potassium chlorate is mixed with about one fourth
its weight of manganese dioxide, the oxygen is given off at a much lower
temperature. Just how the manganese dioxide brings about this result is
not definitely known. The amount of oxygen obtained from a given weight
of potassium chlorate is exactly the same whether the manganese dioxide
is present or not. So far as can be detected the manganese dioxide
undergoes no change.
[Illustration: Fig. 4]
~Directions for preparing oxygen.~ The manner of preparing oxygen from
potassium chlorate is illustrated in the accompanying diagram (Fig. 4).
A mixture consisting of one part of manganese dioxide and four parts of
potassium chlorate is placed in the flask A and gently heated. The
oxygen is evolved and escapes through the tube B. It is collected by
bringing over the end of the tube the mouth of a bottle completely
filled with water and inverted in a vessel of water, as shown in the
figure. The gas rises in the bottle and displaces the water. In the
preparation of large quantities of oxygen, a copper retort (Fig. 5) is
often substituted for the glass flask.
[Illustration: Fig. 5]
In the preparation of oxygen from potassium chlorate and manganese
dioxide, the materials used must be pure, otherwise a violent explosion
may occur. The purity of the materials is tested by heating a small
amount of the mixture in a test tube.
~The collection of gases.~ The method used for collecting oxygen
illustrates the general method used for collecting such gases as are
insoluble in water or nearly so. The vessel C (Fig. 4), containing the
water in which the bottles are inverted, is called a _pneumatic trough._
~Commercial methods of preparation.~ Oxygen can now be purchased stored
under great pressure in strong steel cylinders (Fig. 6). It is prepared
either by heating a mixture of potassium chlorate and manganese dioxide,
or by separating it from the nitrogen and other gases with which it is
mixed in the atmosphere. The methods employed for effecting this
separation will be described in subsequent chapters.
[Illustration: Fig. 6]
~Physical properties.~ Oxygen is a colorless, odorless, tasteless gas,
slightly heavier than air. One liter of it, measured at a temperature of
0° and under a pressure of one atmosphere, weighs 1.4285 g., while under
similar conditions one liter of air weighs 1.2923 g. It is but slightly
soluble in water. Oxygen, like other gases, may be liquefied by applying
very great pressure to the highly cooled gas. When the pressure is
removed the liquid oxygen passes again into the gaseous state, since its
boiling point under ordinary atmospheric pressure is -182.5°.
~Chemical properties.~ At ordinary temperatures oxygen is not very active
chemically. Most substances are either not at all affected by it, or the
action is so slow as to escape notice. At higher temperatures, however,
it is very active, and unites directly with most of the elements. This
activity may be shown by heating various substances until just ignited
and then bringing them into vessels of the gas, when they will burn with
great brilliancy. Thus a glowing splint introduced into a jar of oxygen
bursts into flame. Sulphur burns in the air with a very weak flame and
feeble light; in oxygen, however, the flame is increased in size and
brightness. Substances which readily burn in air, such as phosphorus,
burn in oxygen with dazzling brilliancy. Even substances which burn in
air with great difficulty, such as iron, readily burn in oxygen.
The burning of a substance in oxygen is due to the rapid combination of
the substance or of the elements composing it with the oxygen. Thus,
when sulphur burns both the oxygen and sulphur disappear as such and
there is formed a compound of the two, which is an invisible gas, having
the characteristic odor of burning sulphur. Similarly, phosphorus on
burning forms a white solid compound of phosphorus and oxygen, while
iron forms a reddish-black compound of iron and oxygen.
~Oxidation.~ The term _oxidation_ is applied to the chemical change which
takes place when a substance, or one of its constituent parts, combines
with oxygen. This process may take place rapidly, as in the burning of
phosphorus, or slowly, as in the oxidation (or rusting) of iron when
exposed to the air. It is always accompanied by the liberation of heat.
The amount of heat liberated by the oxidation of a definite weight of
any given substance is always the same, being entirely independent of
the rapidity of the process. If the oxidation takes place slowly, the
heat is generated so slowly that it is difficult to detect it. If the
oxidation takes place rapidly, however, the heat is generated in such a
short interval of time that the substance may become white hot or burst
into a flame.
~Combustion; kindling temperature.~ When oxidation takes place so rapidly
that the heat generated is sufficient to cause the substance to glow or
burst into a flame the process is called _combustion_. In order that any
substance may undergo combustion, it is necessary that it should be
heated to a certain temperature, known as the _kindling temperature._
This temperature varies widely for different bodies, but is always
definite for the same body. Thus the kindling temperature of phosphorus
is far lower than that of iron, but is definite for each. When any
portion of a substance is heated until it begins to burn the combustion
will continue without the further application of heat, provided the heat
generated by the process is sufficient to bring other parts of the
substance to the kindling temperature. On the other hand, if the heat
generated is not sufficient to maintain the kindling temperature,
combustion ceases.
~Oxides.~ The compounds formed by the oxidation of any element are called
_oxides_. Thus in the combustion of sulphur, phosphorus, and iron, the
compounds formed are called respectively oxide of sulphur, oxide of
phosphorus, and oxide of iron. In general, then, _an oxide is a compound
of oxygen with another element_. A great many substances of this class
are known; in fact, the oxides of all the common elements have been
prepared, with the exception of those of fluorine and bromine. Some of
these are familiar compounds. Water, for example, is an oxide of
hydrogen, and lime an oxide of the metal calcium.
~Products of combustion.~ The particular oxides formed by the combustion
of any substance are called _products of combustion_ of that substance.
Thus oxide of sulphur is the product of the combustion of sulphur; oxide
of iron is the product of the combustion of iron. It is evident that the
products of the combustion of any substance must weigh more than the
original substance, the increase in weight corresponding to the amount
of oxygen taken up in the act of combustion. For example, when iron
burns the oxide of iron formed weighs more than the original iron.
In some cases the products of combustion are invisible gases, so that
the substance undergoing combustion is apparently destroyed. Thus, when
a candle burns it is consumed, and so far as the eye can judge nothing
is formed during combustion. That invisible gases are formed, however,
and that the weight of these is greater than the weight of the candle
may be shown by the following experiment.
[Illustration: Fig. 7]
A lamp chimney is filled with sticks of the compound known as
sodium hydroxide (caustic soda), and suspended from the beam of
the balance, as shown in Fig. 7. A piece of candle is placed on
the balance pan so that the wick comes just below the chimney,
and the balance is brought to a level by adding weights to the
other pan. The candle is then lighted. The products formed pass
up through the chimney and are absorbed by the sodium
hydroxide. Although the candle burns away, the pan upon which
it rests slowly sinks, showing that the combustion is attended
by an increase in weight.
~Combustion in air and in oxygen.~ Combustion in air and in
oxygen differs only in rapidity, the products formed being
exactly the same. That the process should take place less
rapidly in the former is readily understood, for the air is
only about one fifth oxygen, the remaining four fifths being
inert gases. Not only is less oxygen available, but much of the
heat is absorbed in raising the temperature of the inert gases
surrounding the substance undergoing combustion, and the
temperature reached in the combustion is therefore less.
~Phlogiston theory of combustion.~ The French chemist Lavoisier
(1743-1794), who gave to oxygen its name was the first to show
that combustion is due to union with oxygen. Previous to his
time combustion was supposed to be due to the presence of a
substance or principle called _phlogiston_. One substance was
thought to be more combustible than another because it
contained more phlogiston. Coal, for example, was thought to be
very rich in phlogiston. The ashes left after combustion would
not burn because all the phlogiston had escaped. If the
phlogiston could be restored in any way, the substance would
then become combustible again. Although this view seems absurd
to us in the light of our present knowledge, it formerly had
general acceptance. The discovery of oxygen led Lavoisier to
investigate the subject, and through his experiments he arrived
at the true explanation of combustion. The discovery of oxygen
together with the part it plays in combustion is generally
regarded as the most important discovery in the history of
chemistry. It marked the dawn of a new period in the growth of
the science.
~Combustion in the broad sense.~ According to the definition given above,
the presence of oxygen is necessary for combustion. The term is
sometimes used, however, in a broader sense to designate any chemical
change attended by the evolution of heat and light. Thus iron and
sulphur, or hydrogen and chlorine under certain conditions, will combine
so rapidly that light is evolved, and the action is called a combustion.
Whenever combustion takes place in the air, however, the process is one
of oxidation.
~Spontaneous combustion.~ The temperature reached in a given
chemical action, such as oxidation, depends upon the rate at
which the reaction takes place. This rate is usually increased
by raising the temperature of the substances taking part in the
action.
When a slow oxidation takes place under such conditions that
the heat generated is not lost by being conducted away, the
temperature of the substance undergoing oxidation is raised,
and this in turn hastens the rate of oxidation. The rise in
temperature may continue in this way until the kindling
temperature of the substance is reached, when combustion
begins. Combustion occurring in this way is called _spontaneous
combustion_.
Certain oils, such as the linseed oil used in paints, slowly
undergo oxidation at ordinary temperatures, and not
infrequently the origin of fires has been traced to the
spontaneous combustion of oily rags. The spontaneous combustion
of hay has been known to set barns on fire. Heaps of coal have
been found to be on fire when spontaneous combustion offered
the only possible explanation.
~Importance of oxygen.~ 1. Oxygen is essential to life. Among living
organisms only certain minute forms of plant life can exist without it.
In the process of respiration the air is taken into the lungs where a
certain amount of oxygen is absorbed by the blood. It is then carried to
all parts of the body, oxidizing the worn-out tissues and changing them
into substances which may readily be eliminated from the body. The heat
generated by this oxidation is the source of the heat of the body. The
small amount of oxygen which water dissolves from the air supports all
the varied forms of aquatic animals.
2. Oxygen is also essential to decay. The process of decay is really a
kind of oxidation, but it will only take place in the presence of
certain minute forms of life known as bacteria. Just how these assist in
the oxidation is not known. By this process the dead products of animal
and vegetable life which collect on the surface of the earth are slowly
oxidized and so converted into harmless substances. In this way oxygen
acts as a great purifying agent.
3. Oxygen is also used in the treatment of certain diseases in which the
patient is unable to inhale sufficient air to supply the necessary
amount of oxygen.
OZONE
~Preparation.~ When electric sparks are passed through oxygen or air a
small percentage of the oxygen is converted into a substance called
_ozone_, which differs greatly from oxygen in its properties. The same
change can also be brought about by certain chemical processes. Thus, if
some pieces of phosphorus are placed in a bottle and partially covered
with water, the presence of ozone may soon be detected in the air
contained in the bottle. The conversion of oxygen into ozone is attended
by a change in volume, 3 volumes of oxygen forming 2 volumes of ozone.
If the resulting ozone is heated to about 300°, the reverse change
takes place, the 2 volumes of ozone being changed back into 3 volumes of
oxygen. It is possible that traces of ozone exist in the atmosphere,
although its presence there has not been definitely proved, the tests
formerly used for its detection having been shown to be unreliable.
~Properties.~ As commonly prepared, ozone is mixed with a large excess of
oxygen. It is possible, however, to separate the ozone and thus obtain
it in pure form. The gas so obtained has the characteristic odor noticed
about electrical machines when in operation. By subjecting it to great
pressure and a low temperature, the gas condenses to a bluish liquid,
boiling at -119°. When unmixed with other gases ozone is very explosive,
changing back into oxygen with the liberation of heat. Its chemical
properties are similar to those of oxygen except that it is far more
active. Air or oxygen containing a small amount of ozone is now used in
place of oxygen in certain manufacturing processes.
~The difference between oxygen and ozone.~ Experiments show that in
changing oxygen into ozone no other kind of matter is either added to
the oxygen or withdrawn from it. The question arises then, How can we
account for the difference in their properties? It must be remembered
that in all changes we have to take into account _energy_ as well as
_matter_. By changing the amount of energy in a substance we change its
properties. That oxygen and ozone contain different amounts of energy
may be shown in a number of ways; for example, by the fact that the
conversion of ozone into oxygen is attended by the liberation of heat.
The passage of the electric sparks through oxygen has in some way
changed the energy content of the element and thus it has acquired new
properties. _Oxygen and ozone must, therefore, be regarded as identical
so far as the kind of matter of which they are composed is concerned.
Their different properties are due to their different energy contents._
~Allotropic states or forms of matter.~ Other elements besides oxygen may
exist in more than one form. These different forms of the same element
are called _allotropic states_ or _forms_ of the element. These forms
differ not only in physical properties but also in their energy
contents. Elements often exist in a variety of forms which look quite
different. These differences may be due to accidental causes, such as
the size or shape of the particles or the way in which the element was
prepared. Only such forms, however, as have different energy contents
are properly called allotropic forms.
MEASUREMENT OF GAS VOLUMES
~Standard conditions.~ It is a well-known fact that the volume occupied by
a definite weight of any gas can be altered by changing the temperature
of the gas or the pressure to which it is subjected. In measuring the
volume of gases it is therefore necessary, for the sake of accuracy, to
adopt some standard conditions of temperature and pressure. The
conditions agreed upon are (1) a temperature of 0°, and (2) a pressure
equal to the average pressure exerted by the atmosphere at the sea
level, that is, 1033.3 g. per square centimeter. These conditions of
temperature and pressure are known as the _standard conditions_, and
when the volume of a gas is given it is understood that the measurement
was made under these conditions, unless it is expressly stated
otherwise. For example, the weight of a liter of oxygen has been given
as 1.4285 g. This means that one liter of oxygen, measured at a
temperature of 0° and under a pressure of 1033.3 g. per square
centimeter, weighs 1.4285 g.
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°, then the volume at 1° will be
V + 1/273 V; at 2° 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° is 750 cc.; find the volume it will occupy at
0°, 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 × 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 × 740 =
v × 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° 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 × 740)/(760 (1 + 0.00366 × 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° 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° (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 × 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°, the pressure remaining constant:
(a) 150 cc. at 10°; (b) 840 cc. at 273°.
8. A certain volume of gas is measured when the temperature is 20°. 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°?
11. Reduce the following volumes to standard conditions of temperature
and pressure: (a) 340 cc. at 12° and 753 mm; (b) 500 cc. at 15° and
740 mm.
12. What weight of potassium chlorate is necessary to prepare 250 l. of
oxygen at 20° 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?