CARBON AND SOME OF ITS SIMPLER COMPOUNDS
~The family.~ Carbon stands at the head of a family of elements in the
fourth group in the periodic table. The resemblances between the
elements of this family, while quite marked, are not so striking as in
the case of the elements of the chlorine family. With the exception of
carbon, these elements are comparatively rare, and need not be taken up
in detail in this chapter. Titanium will be referred to again in
connection with silicon which it very closely resembles.
~Occurrence.~ Carbon is found in nature in the uncombined state in several
forms. The diamond is practically pure carbon, while graphite and coal
are largely carbon, but contain small amounts of other substances. Its
natural compounds are exceedingly numerous and occur as gases, liquids,
and solids. Carbon dioxide is its most familiar gaseous compound.
Natural gas and petroleum are largely compounds of carbon with hydrogen.
The carbonates, especially calcium carbonate, constitute great strata of
rocks, and are found in almost every locality. All living organisms,
both plant and animal, contain a large percentage of this element, and
the number of its compounds which go to make up all the vast variety of
animate nature is almost limitless. Over one hundred thousand definite
compounds containing carbon have been prepared. In the free state carbon
occurs in three allotropic forms, two of which are crystalline and one
amorphous.
~Crystalline carbon.~ Crystalline carbon occurs in two forms,--diamond and
graphite.
1. _Diamond._ Diamonds are found in considerable quantities in several
localities, especially in South Africa, the East Indies, and Brazil. The
crystals belong to the regular system, but the natural stones do not
show this very clearly. When found they are usually covered with a rough
coating which is removed in the process of cutting. Diamond cutting is
carried on most extensively in Holland.
The density of the diamond is 3.5, and, though brittle, it is one of the
hardest of substances. Black diamonds, as well as broken and imperfect
stones which are valueless as gems, are used for grinding hard
substances. Few chemical reagents have any action on the diamond, but
when heated in oxygen or the air it blackens and burns, forming carbon
dioxide.
Lavoisier first showed that carbon dioxide is formed by the combustion
of the diamond; and Sir Humphry Davy in 1814 showed that this is the
only product of combustion, and that the diamond is pure carbon.
~The diamond as a gem.~ The pure diamond is perfectly transparent
and colorless, but many are tinted a variety of colors by
traces of foreign substances. Usually the colorless ones are
the most highly prized, although in some instances the color
adds to the value; thus the famous Hope diamond is a beautiful
blue. Light passing through a diamond is very much refracted,
and to this fact the stone owes its brilliancy and sparkle.
~Artificial preparation of diamonds.~ Many attempts have been
made to produce diamonds artificially, but for a long time
these always ended in failure, graphite and not diamonds being
the product obtained. The French chemist Moissan, in his
extended study of chemistry at high temperatures, finally
succeeded (1893) in making some small ones. He accomplished
this by dissolving carbon in boiling iron and plunging the
crucible containing the mixture into water, as shown in Fig.
58. Under these conditions the carbon crystallized in the iron
in the form of the diamond. The diamonds were then obtained by
dissolving away the iron in hydrochloric acid.
[Illustration: Fig. 58]
2. _Graphite._ This form of carbon is found in large quantities,
especially in Ceylon, Siberia, and in some localities of the United
States and Canada. It is a shining black substance, very soft and greasy
to the touch. Its density is about 2.15. It varies somewhat in
properties according to the locality in which it is found, and is more
easily attacked by reagents than is the diamond. It is also manufactured
by heating carbon with a small amount of iron (3%) in an electric
furnace. It is used in the manufacture of lead pencils and crucibles, as
a lubricant, and as a protective covering for iron in the form of a
polish or a paint.
~Amorphous carbon.~ Although there are many varieties of amorphous carbon
known, they are not true allotropic modifications. They differ merely in
their degree of purity, their fineness of division, and in their mode of
preparation. These substances are of the greatest importance, owing to
their many uses in the arts and industries. As they occur in nature, or
are made artificially, they are nearly all impure carbon, the impurity
depending on the particular substance in question.
1. _Pure carbon._ Pure amorphous carbon is best prepared by charring
sugar. This is a substance consisting of carbon, hydrogen, and oxygen,
the latter two elements being present in the ratio of one oxygen atom to
two of hydrogen. When sugar is strongly heated the oxygen and hydrogen
are driven off in the form of water and pure carbon is left behind.
Prepared in this way it is a soft, lustrous, very bulky, black powder.
2. _Coal and coke._ Coals of various kinds were probably formed from
vast accumulations of vegetable matter in former ages, which became
covered over with earthy material and were thus protected from rapid
decay. Under various natural agencies the organic matter was slowly
changed into coal. In anthracite these changes have gone the farthest,
and this variety of coal is nearly pure carbon. Soft or bituminous coals
contain considerable organic matter besides carbon and mineral
substances. When heated strongly out of contact with air the organic
matter is decomposed and the resulting volatile matter is driven off in
the form of gases and vapors, and only the mineral matter and carbon
remain behind. The gaseous product is chiefly illuminating gas and the
solid residue is _coke_. Some of the coke is found as a dense cake on
the sides and roof of the retort. This is called retort carbon and is
quite pure.
3. _Charcoal._ This is prepared from wood in the same way that coke is
made from coal. When the process is carried on in retorts the products
expelled by the heat are saved. Among these are many valuable substances
such as wood alcohol and acetic acid. Where timber is abundant the
process is carried out in a wasteful way, by merely covering piles of
wood with sod and setting the wood on fire. Some wood burns and the heat
from this decomposes the wood not burned, forming charcoal from it. The
charcoal, of course, contains the mineral part of the wood from which it
is formed.
4. _Bone black._ This is sometimes called animal charcoal, and is made
by charring bones and animal refuse. The organic part of the materials
is thus decomposed and carbon is left in a very finely divided state,
scattered through the mineral part which consists largely of calcium
phosphate. For some uses this mineral part is removed by treatment with
hydrochloric acid and prolonged washing.
5. _Lampblack._ Lampblack and soot are products of imperfect combustion
of oil and coal, and are deposited from a smoky flame on a cold surface.
The carbon in this form is very finely divided and usually contains
various oily materials.
~Properties.~ While the various forms of carbon differ in many properties,
especially in color and hardness, yet they are all odorless, tasteless
solids, insoluble in water and characterized by their stability towards
heat. Only in the intense heat of the electric arc does carbon
volatilize, passing directly from the solid state into a vapor. Owing to
this fact the inside surface of an incandescent light bulb after being
used for some time becomes coated with a dark film of carbon. It is not
acted on at ordinary temperatures by most reagents, but at a higher
temperature it combines directly with many of the elements, forming
compounds called _carbides_. When heated in the presence of sufficient
oxygen it burns, forming carbon dioxide.
~Uses of carbon.~ The chief use of amorphous carbon is for fuel to furnish
heat and power for all the uses of civilization. An enormous quantity of
carbon in the form of the purer coals, coke, and charcoal is used as a
reducing agent in the manufacture of the various metals, especially in
the metallurgy of iron. Most of the metals are found in nature as
oxides, or in forms which can readily be converted into oxides. When
these oxides are heated with carbon the oxygen is abstracted, leaving
the metal. Retort carbon and coke are used to make electric light
carbons and battery plates, while lampblack is used for indelible inks,
printer's ink, and black varnishes. Bone black and charcoal have the
property of absorbing large volumes of certain gases, as well as smaller
amounts of organic matter; hence they are used in filters to remove
noxious gases and objectionable colors and odors from water. Bone black
is used extensively in the sugar refineries to remove coloring matter
from the impure sugars.
~Chemistry of carbon compounds.~ Carbon is remarkable for the very large
number of compounds which it forms with the other elements, especially
with oxygen and hydrogen. Compounds containing carbon are more numerous
than all others put together, and the chemistry of these substances
presents peculiarities not met with in the study of other substances.
For these reasons the systematic study of carbon compounds, or of
_organic chemistry_ as it is usually called, must be deferred until the
student has gained some knowledge of the chemistry of other elements. An
acquaintance with a few of the most familiar carbon compounds is,
however, essential for the understanding of the general principles of
chemistry.
~Compounds of carbon with hydrogen,--the hydrocarbons.~ Carbon unites with
hydrogen to form a very large number of compounds called _hydrocarbons_.
Petroleum and natural gas are essentially mixtures of a great variety of
these hydrocarbons. Many others are found in living plants, and still
others are produced by the decay of organic matter in the absence of
air. Only two of them, methane and acetylene, will be discussed here.
~Methane~ (_marsh gas_) (CH_{4}). This is one of the most important of
these hydrocarbons, and constitutes about nine tenths of natural gas. As
its name suggests, it is formed in marshes by the decay of vegetable
matter under water, and bubbles of the gas are often seen to rise when
the dead leaves on the bottom of pools are stirred. It also collects in
mines, and, when mixed with air, is called _fire damp_ by the miners
because of its great inflammability, damp being an old name for a gas.
It is formed when organic matter, such as coal or wood, is heated in
closed vessels, and is therefore a principal constituent of coal gas.
~Preparation.~ Methane is prepared in the laboratory by heating sodium or
calcium acetate with soda-lime. Equal weights of fused sodium acetate
and soda-lime are thoroughly dried, then mixed and placed in a
good-sized, hard-glass test tube fitted with a one-holed stopper and
delivery tube. The mixture is gradually heated, and when the air has
been displaced from the tube the gas is collected in bottles by
displacement of water. Soda-lime is a mixture of sodium and calcium
hydroxides. Regarding it as sodium hydroxide alone, the equation is
NaC_{2}H_{3}O_{2} + NaOH = Na_{2}CO_{3} + CH_{4}.
~Properties.~ Methane is a colorless, odorless gas whose density is 0.55.
It is difficult to liquefy, boiling at -155° under standard pressure,
and is almost insoluble in water. It burns with a pale blue flame,
liberating much heat, and when mixed with oxygen is very explosive.
~Davy's safety lamp.~ In 1815 Sir Humphry Davy invented a lamp for the use
of miners, to prevent the dreadful mine explosions then common, due to
methane mixed with air. The invention consisted in surrounding the upper
part of the common miner's lamp with a mantle of wire gauze and the
lower part with glass (Fig. 59). It has been seen that two gases will
not combine until raised to their kindling temperature, and if while
combining they are cooled below this point, the combination ceases. A
flame will not pass through a wire gauze because the metal, being a good
conductor of heat, takes away so much heat from the flame that the gases
are cooled below the kindling temperature. When a lamp so protected is
brought into an explosive mixture the gases inside the wire mantle burn
in a series of little explosions, giving warning to the miner that the
air is unsafe.
[Illustration: Fig. 59]
~Acetylene~ (C_{2}H_{2}). This is a colorless gas usually having a
disagreeable odor due to impurities. It is now made in large quantities
from calcium carbide (CaC_{2}). This substance is formed when coal and
lime are heated together in an electric furnace. When treated with water
the carbide is decomposed, yielding acetylene:
CaC_{2} + 2H_{2}O = C_{2}H_{2} + Ca(OH)_{2}.
Under ordinary conditions the gas burns with a very smoky flame; in
burners constructed so as to secure a large amount of oxygen it burns
with a very brilliant white light, and hence is used as an illuminant.
~Laboratory preparation.~ The gas can be prepared readily in a generator
such as is shown in Fig. 60. The inner tube contains fragments of
calcium carbide, while the outer one is filled with water. As long as
the stopcock is closed the water cannot rise in the inner tube. When the
stopcock is open the water rises, and, coming into contact with the
carbide in the inner tube, generates acetylene. This escapes through the
stopcock, and after the air has been expelled may be lighted as it
issues from the burner.
[Illustration: Fig. 60]
Carbon forms two oxides, namely, carbon dioxide (CO_{2}) and carbon
monoxide (CO).
~Carbon dioxide~ (CO_{2}). Carbon dioxide is present in the air to the
extent of about 3 parts in 10,000, and this apparently small amount is
of fundamental importance in nature. In some localities it escapes from
the earth in great quantities, and many spring waters carry large
amounts of it in solution. When these highly charged spring waters reach
the surface of the earth, and the pressure on them is removed, the
carbon dioxide escapes with effervescence. It is a product of the
oxidation of all organic matter, and is therefore formed in fires as
well as in the process of decay. It is thrown off from the lungs of all
animals in respiration, and is a product of many fermentation processes
such as vinegar making and brewing. Combined with metallic oxides it
forms vast deposits of carbonates in nature.
~Preparation.~ In the laboratory carbon dioxide is always prepared by the
action of an acid upon a carbonate, usually calcium carbonate, the
apparatus shown in Fig. 39 serving the purpose very well. This reaction
might be expected to produce carbonic acid, thus:
CaCO_{3} + 2HCl = CaCl_{2} + H_{2}CO_{3}.
Carbonic acid is very unstable, however, and decomposes into its
anhydride, CO_{2}, and water, thus:
H_{2}CO_{3} = H_{2}O + CO_{2}.
The complete reaction is represented by the equation
CaCO_{3} + 2HCl = CaCl_{2} + CO_{2} + H_{2}O.
~Physical properties.~ Carbon dioxide is a colorless, practically odorless
gas whose density is 1.5. Its weight may be inferred from the fact that
it can be siphoned, or poured like water, from one vessel downward into
another. At 15° and under ordinary pressure it dissolves in its own
volume of water and imparts a somewhat biting, pungent taste to it. It
is easily condensed, and is now prepared commercially in this form by
pumping the gas into steel cylinders (see Fig. 6) which are kept cold
during the process. When the liquid is permitted to escape into the air
part of it instantly evaporates, and in so doing absorbs so much heat
that another portion is solidified, the solid form strikingly resembling
snow in appearance. This snow is very cold and mercury can easily be
frozen with it.
~Solid carbon dioxide.~ Cylinders of liquid carbon dioxide are
inexpensive, and should be available in every school. To demonstrate the
properties of solid carbon dioxide, the cylinder should be placed across
the table and supported in such a way that the stopcock end is several
inches lower than the other end. A loose bag is made by holding the
corners of a handkerchief around the neck of the stopcock, and the cock
is then turned on so that the gas rushes out in large quantities. Very
quickly a considerable quantity of the snow collects in the
handkerchief. To freeze mercury, press a piece of filter paper into a
small evaporating dish and pour the mercury upon it. Coil a flat spiral
upon the end of a wire, and dip the spiral into the mercury. Place a
quantity of solid carbon dioxide upon the mercury and pour 10 cc.-15 cc.
of ether over it. In a minute or two the mercury will solidify and may
be removed from the dish by the wire serving as a handle. The filter
paper is to prevent the mercury from sticking to the dish; the ether
dissolves the solid carbon dioxide and promotes its rapid conversion
into gas.
~Chemical properties.~ Carbon dioxide is incombustible, since it is, like
water, a product of combustion. It does not support combustion, as does
nitrogen peroxide, because the oxygen in it is held in very firm
chemical union with the carbon. Very strong reducing agents, such as
highly heated carbon, can take away half of its oxygen:
CO_{2} + C = 2CO.
~Uses.~ The relation of carbon dioxide to plant life has been discussed in
a previous chapter. Water highly charged with carbon dioxide is used for
making soda water and similar beverages. Since it is a non-supporter of
combustion and can be generated readily, carbon dioxide is also used as
a fire extinguisher. Some of the portable fire extinguishers are simply
devices for generating large amounts of the gas. It is not necessary
that all the oxygen should be kept away from the fire in order to
smother it. A burning candle is extinguished in air which contains only
2.5% of carbon dioxide.
~Carbonic acid~ (H_{2}CO_{3}). Like most of the oxides of the non-metallic
elements, carbon dioxide is an acid anhydride. It combines with water to
form an acid of the formula H_{2}CO_{3}, called carbonic acid:
H_{2}O + CO_{2} = H_{2}CO_{3}.
The acid is, however, very unstable and cannot be isolated. Only a very
small amount of it is actually formed when carbon dioxide is passed into
water, as is evident from the small solubility of the gas. If, however,
a base is present in the water, salts of carbonic acid are formed, and
these are quite stable:
2NaOH + H_{2}O + CO_{2} = Na_{2}CO_{3} + 2H_{2}O.
~Action of carbon dioxide on bases.~ This conduct is explained by the
principles of reversible reactions. The equation
H_{2}O +CO_{2} <--> H_{2}CO_{3}
is a reversible equation, and the extent to which the reaction
progresses depends upon the relative concentrations of each of the three
factors in it. Equilibrium is ordinarily reached when very little
H_{2}CO_{3} is formed. If a base is present in the water to combine with
the H_{2}CO_{3} as fast as it is formed, all of the CO_{2} is converted
into H_{2}CO_{3}, and thence into a carbonate.
~Salts of carbonic acid,--carbonates.~ The carbonates form a very
important class of salts. They are found in large quantities in nature,
and are often used in chemical processes. Only the carbonates of sodium,
potassium, and ammonium are soluble, and these can be made by the action
of carbon dioxide on solutions of the bases, as has just been explained.
The insoluble carbonates are formed as precipitates when soluble salts
are treated with a solution of a soluble carbonate. Thus the insoluble
calcium carbonate can be made by bringing together solutions of calcium
chloride and sodium carbonate:
CaCl_{2} + Na_{2}CO_{3} = CaCO_{3} + 2NaCl.
Most of the carbonates are decomposed by heat, yielding an oxide of the
metal and carbon dioxide. Thus lime (calcium oxide) is made by strongly
heating calcium carbonate:
CaCO_{3} = CaO + CO_{2}.
~Acid carbonates.~ Like all acids containing two acid hydrogen atoms,
carbonic acid can form both normal and acid salts. The acid carbonates
are made by treating a normal carbonate with an excess of carbonic acid.
With few exceptions they are very unstable, heat decomposing them even
when in solution.
~Action of carbon dioxide on calcium hydroxide.~ If carbon dioxide is
passed into clear lime water, calcium carbonate is at first
precipitated:
H_{2}O + CO_{2} = H_{2}CO_{3},
Ca(OH)_{2} + H_{2}CO_{3} = CaCO_{3} + 2H_{2}O.
Advantage is taken of this reaction in testing for the presence of
carbon dioxide, as already explained in the chapter on the atmosphere.
If the current of carbon dioxide is continued, the precipitate soon
dissolves, because the excess of carbonic acid forms calcium acid
carbonate which is soluble:
CaCO_{3} + H_{2}CO_{3} = Ca(HCO_{3})_{2}.
If now the solution is heated, the acid carbonate is decomposed and
calcium carbonate once more precipitated:
Ca(HCO_{3})_{2} = CaCO_{3} + H_{2}CO_{3}.
~Carbon monoxide (CO).~ Carbon monoxide can be made in a number of ways,
the most important of which are the three following:
1. _By the partial oxidation of carbon._ If a slow current of air is
conducted over highly heated carbon, the monoxide is formed, thus:
C + O = CO
It is therefore often formed in stoves when the air draught is
insufficient. Water gas, which contains large amounts of carbon
monoxide, is made by partially oxidizing carbon with steam:
C + H_{2}O = CO + 2H.
2. _By the partial reduction of carbon dioxide._ When carbon dioxide is
conducted over highly heated carbon it is reduced to carbon monoxide by
the excess of carbon:
CO_{2} + C = 2CO.
When coal is burning in a stove or grate carbon dioxide is at first
formed in the free supply of air, but as the hot gas rises through the
glowing coal it is reduced to carbon monoxide. When the carbon monoxide
reaches the free air above the coal it takes up oxygen to form carbon
dioxide, burning with the blue flame so familiar above a bed of coals,
especially in the case of hard coals.
3. _By the decomposition of oxalic acid._ In the laboratory carbon
monoxide is usually prepared by the action of concentrated sulphuric
acid upon oxalic acid. The latter substance has the formula
C_{2}H_{2}O_{4}. The sulphuric acid, owing to its affinity for water,
decomposes the oxalic acid, as represented in the equation
C_{2}H_{2}O_{4} + (H_{2}SO_{4}) = (H_{2}SO_{4}) + H_{2}O + CO_{2} + CO.
~Properties.~ Carbon monoxide is a light, colorless, almost odorless gas,
very difficult to liquefy. Chemically it is very active, combining
directly with a great many substances. It has a great affinity for
oxygen and is therefore combustible and a good reducing agent. Thus, if
carbon monoxide is passed over hot copper oxide, the copper is reduced
to the metallic state:
CuO + CO = Cu + CO_{2}.
When inhaled it combines with the red coloring matter of the blood and
in this way prevents the absorption of oxygen, so that even a small
quantity of the gas may prove fatal.
[Illustration: Fig. 61]
~The reducing power of carbon monoxide.~ Fig. 61 illustrates a
method of showing the reducing power of carbon monoxide. The
gas is generated by gently heating 7 or 8 g. of oxalic acid
with 25 cc. of concentrated sulphuric acid in a 200 cc. flask
A. The bottle B contains a solution of sodium hydroxide,
which removes the carbon dioxide formed along with the
monoxide. C contains a solution of calcium hydroxide to show
that the carbon dioxide is completely removed. E is a
hard-glass tube containing 1 or 2 g. of copper oxide, which is
heated by a burner. The black copper oxide is reduced to
reddish metallic copper by the carbon monoxide, which is
thereby changed to carbon dioxide. The presence of the carbon
dioxide is shown by the precipitate in the calcium hydroxide
solution in D. Any unchanged carbon monoxide is collected
over water in F.
~Carbon disulphide~ (CS_{2}). Just as carbon combines with oxygen to form
carbon dioxide, so it combines with sulphur to form carbon disulphide
(CS_{2}). This compound has been described in the chapter on sulphur.
~Hydrocyanic acid~ (_prussic acid_)(HCN). Under the proper conditions
carbon unites with nitrogen and hydrogen to form the acid HCN, called
hydrocyanic acid. It is a weak, volatile acid, and is therefore easily
prepared by treating its salts with sulphuric acid:
KCN + H_{2}SO_{4} = KHSO_{4} + HCN.
It is most familiar as a gas, though it condenses to a colorless liquid
boiling at 26°. It has a peculiar odor, suggesting bitter almonds, and
is extremely poisonous either when inhaled or when taken into the
stomach. A single drop may cause death. It dissolves readily in water,
its solution being commonly called prussic acid.
The salts of hydrocyanic acid are called _cyanides_, the cyanides of
sodium and potassium being the best known. These are white solids and
are extremely poisonous.
~Solutions of potassium cyanide are alkaline.~ A solution of potassium
cyanide turns red litmus blue, and must therefore contain hydroxyl ions.
The presence of these ions is accounted for in the following way.
Although water is so little dissociated into its ions H^{+} and OH^{-}
that for most purposes we may neglect the dissociation, it is
nevertheless measurably dissociated. Hydrocyanic acid is one of the
weakest of acids, and dissociates to an extremely slight extent. When a
cyanide such as potassium cyanide dissolves it freely dissociates, and
the CN^{-} ions must come to an equilibrium with the H^{+} ions derived
from the water:
H^{+} + CN^{-} <--> HCN.
The result of this equilibrium is that quite a number of H^{+} ions from
the water are converted into undissociated HCN molecules. But for every
H^{+} ion so removed an OH^{-} ion remains free, and this will give the
solution alkaline properties.
EXERCISES
1. How can you prove that the composition of the different allotropic
forms of carbon is the same?
2. Are lampblack and bone black allotropic forms of carbon? Will equal
amounts of heat be liberated in the combustion of 1 g. of each?
3. How could you judge of the relative purity of different forms of
carbon?
4. Apart from its color, why should carbon be useful in the preparation
of inks and paints?
5. Could asbestos fibers be used to replace the wire in a safety lamp?
6. Why do most acids decompose carbonates?
7. What effect would doubling the pressure have upon the solubility of
carbon dioxide in water?
8. What compound would be formed by passing carbon dioxide into a
solution of ammonium hydroxide? Write the equation.
9. Write equations for the preparation of K_{2}CO_{3}; of BaCO_{3}; of
MgCO_{3}.
10. In what respects are carbonic and sulphurous acids similar?
11. Give three reasons why the reaction which takes place when a
solution of calcium acid carbonate is heated, completes itself.
12. How could you distinguish between carbonates and sulphites?
13. How could you distinguish between oxygen, hydrogen, nitrogen,
nitrous oxide, and carbon dioxide?
14. Could a solution of sodium hydroxide be substituted for the solution
of calcium hydroxide in testing for carbon dioxide?
15. What weight of sodium hydroxide is necessary to neutralize the
carbonic acid formed by the action of hydrochloric acid on 100 g. of
calcium carbonate?
16. What weight of calcium carbonate would be necessary to prepare
sufficient carbon dioxide to saturate 10 l. of water at 15° and under
ordinary pressure?
17. On the supposition that calcium carbide costs 12 cents a kilogram,
what would be the cost of an amount sufficient to generate 100 l. of
acetylene measured at 20° and 740 mm.?
18. How would the volume of a definite amount of carbon monoxide compare
with the volume of carbon dioxide formed by its combustion, the
measurements being made under the same conditions?