COMPOUNDS OF NITROGEN
~Occurrence.~ As has been stated in a former chapter, nitrogen constitutes
a large fraction of the atmosphere. The compounds of nitrogen, however,
cannot readily be obtained from this source, since at any ordinary
temperature nitrogen is able to combine directly with very few of the
elements.
In certain forms of combination nitrogen occurs in the soil from which
it is taken up by plants and built into complex substances composed
chiefly of carbon, hydrogen, oxygen, and nitrogen. Animals feeding on
these plants assimilate the nitrogenous matter, so that this element is
an essential constituent of both plants and animals.
~Decomposition of organic matter by bacteria.~ When living matter dies and
undergoes decay complicated chemical reactions take place, one result of
which is that the nitrogen of the organic matter is set free either as
the element nitrogen, or in the form of simple compounds, such as
ammonia (NH_{3}) or oxides of nitrogen. Experiment has shown that all
such processes of decay are due to the action of different kinds of
bacteria, each particular kind effecting a different change.
~Decomposition of organic matter by heat.~ When organic matter is strongly
heated decomposition into simpler substances takes place in much the
same way as in the case of bacterial decomposition. Coal is a complex
substance of vegetable origin, consisting largely of carbon, but also
containing hydrogen, oxygen, and nitrogen. When this is heated in a
closed vessel so that air is excluded, about one seventh of the nitrogen
is converted into ammonia, and this is the chief source from which
ammonia and its compounds are obtained.
COMPOUNDS OF NITROGEN WITH HYDROGEN
~Ammonia~ (NH_{3}). Several compounds consisting exclusively of nitrogen
and hydrogen are known, but only one, ammonia, need be considered here.
~Preparation of ammonia.~ Ammonia is prepared in the laboratory by a
different method from the one which is used commercially.
1. _Laboratory method._ In the laboratory ammonia is prepared from
ammonium chloride, a compound having the formula NH_{4}Cl, and obtained
in the manufacture of coal gas. As will be shown later in the chapter,
the group NH_{4} in this compound acts as a univalent radical and is
known as _ammonium_. When ammonium chloride is warmed with sodium
hydroxide, the ammonium and sodium change places, the reaction being
expressed in the following equation.
NH_{4}Cl + NaOH = NaCl + NH_{4}OH.
The ammonium hydroxide (NH_{4}OH) so formed is unstable and breaks down
into water and ammonia.
NH_{4}OH = NH_{3} + H_{2}O.
Calcium hydroxide (Ca(OH)_{2}) is frequently used in place of the more
expensive sodium hydroxide, the equations being
2NH_{4}Cl + Ca(OH)_{2} = CaCl_{2} + 2NH_{4}OH,
2NH_{4}OH = 2H_{2}O + 2NH_{3}.
In the preparation, the ammonium chloride and calcium hydroxide
are mixed together and placed in a flask arranged as shown in
Fig. 35. The mixture is gently warmed, when ammonia is evolved
as a gas and is collected by displacement of air.
[Illustration: Fig. 35]
2. _Commercial method._ Nearly all the ammonia of commerce comes from
the gasworks. Ordinary illuminating gas is made by distilling coal, as
will be explained later, and among the products of this distillation a
solution of ammonia in water is obtained. This solution, known as _gas
liquor_, contains not only ammonia but other soluble substances. Most of
these combine chemically with lime, while ammonia does not; if then lime
is added to the gas liquor and the liquor is heated, the ammonia is
driven out from the mixture. It may be dissolved again in pure, cold
water, forming _aqua ammonia_, or the ammonia water of commerce.
~Preparation from hydrogen and nitrogen.~ When electric sparks
are passed for some time through a mixture of hydrogen and
nitrogen, a small percentage of the two elements in the mixture
is changed into ammonia. The action soon ceases, however, for
the reason that ammonia is decomposed by the electric
discharge. The reaction expressed in the equation
N + 3H = NH_{3}
can therefore go in either direction depending upon the
relative quantities of the substances present. This recalls the
similar change from oxygen into ozone, which soon ceases
because the ozone is in turn decomposed into oxygen.
~Physical properties.~ Under ordinary conditions ammonia is a gas whose
density is 0.59. It is therefore little more than half as heavy as air.
It is easily condensed into a colorless liquid, and can now be purchased
in liquid form in steel cylinders. The gas is colorless and has a
strong, suffocating odor. It is extremely soluble in water, 1 l. of
water at 0° and 760 mm. pressure dissolving 1148 l. of the gas. In
dissolving this large volume of gas the water expands considerably, so
that the density of the solution is less than that of water, the
strongest solutions having a density of 0.88.
~Chemical properties.~ Ammonia will not support combustion, nor will it
burn under ordinary conditions. In an atmosphere of oxygen it burns with
a feeble, yellowish flame. When quite dry it is not a very active
substance, but when moist it combines with a great many substances,
particularly with acids.
~Uses.~ It has been stated that ammonia can be condensed to a liquid by
the application of pressure. If the pressure is removed from the liquid
so obtained, it rapidly passes again into the gaseous state and in so
doing absorbs a large amount of heat. Advantage is taken of this fact in
the preparation of artificial ice. Large quantities of ammonia are also
used in the preparation of ammonium compounds.
~The manufacture of artificial ice.~ Fig. 36 illustrates the
method of preparing artificial ice. The ammonia gas is
liquefied in the pipes X by means of the pump Y. The heat
generated is absorbed by water flowing over the pipes. The
pipes lead into a large brine tank, a cross section of which is
shown in the figure. Into the brine (concentrated solution of
common salt) contained in this tank are dipped the vessels A,
B, C, filled with pure water. The pressure is removed from
the liquid ammonia as it passes into the pipes immersed in the
brine, and the heat absorbed by the rapid evaporation of the
liquid lowers the temperature of the brine below zero. The
water in A, B, C is thereby frozen into cakes of ice. The
gaseous ammonia resulting from the evaporation of the liquid
ammonia is again condensed, so that the process is continuous.
[Illustration Fig. 36]
~Ammonium hydroxide~ (NH_{4}OH). The solution of ammonia in water is found
to have strong basic properties and therefore contains hydroxyl ions. It
turns red litmus blue; it has a soapy feel; it neutralizes acids,
forming salts with them. It seems probable, therefore, that when ammonia
dissolves in water it combines chemically with it according to the
equation
NH_{3} + H_{2}O = NH_{4}OH,
and that it is the substance NH_{4}OH, called ammonium hydroxide, which
has the basic properties, dissociating into the ions NH_{4} and OH.
Ammonium hydroxide has never been obtained in a pure state. At every
attempt to isolate it the substance breaks up into water and ammonia,--
NH_{4}OH = NH_{3} + H_{2}O.
~The ammonium radical.~ The radical NH_{4} plays the part of a metal in
many chemical reactions and is called ammonium. The ending _-ium_ is
given to the name to indicate the metallic properties of the substance,
since the names of the metals in general have that ending. The salts
formed by the action of the base ammonium hydroxide on acids are called
ammonium salts. Thus, with hydrochloric acid, ammonium chloride is
formed in accordance with the equation
NH_{4}OH + HCl = NH_{4}Cl + H_{2}O.
Similarly, with nitric acid, ammonium nitrate (NH_{4}NO_{3}) is formed,
and with sulphuric acid, ammonium sulphate ((NH_{4})_{2}S0_{4}).
It will be noticed that in the neutralization of ammonium hydroxide by
acids the group NH_{4} replaces one hydrogen atom of the acid, just as
sodium does. The group therefore acts as a univalent metal.
~Combination of nitrogen with hydrogen by volume.~ Under suitable
conditions ammonia can be decomposed into nitrogen and hydrogen by
passing electric sparks through the gas. Accurate measurement has shown
that when ammonia is decomposed, two volumes of the gas yield one volume
of nitrogen and three volumes of hydrogen. Consequently, if the two
elements were to combine directly, one volume of nitrogen would combine
with three volumes of hydrogen to form two volumes of ammonia. Here, as
in the formation of steam from hydrogen and oxygen, small whole numbers
serve to indicate the relation between the volumes of combining gases
and that of the gaseous product.
COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROGEN
In addition to ammonium hydroxide, nitrogen forms several compounds with
hydrogen and oxygen, of which nitric acid (HNO_{3}) and nitrous acid
(HNO_{2}) are the most familiar.
~Nitric acid~ (HNO_{3}). Nitric acid is not found to any extent in nature,
but some of its salts, especially sodium nitrate (NaNO_{3}) and
potassium nitrate (KNO_{3}) are found in large quantities. From these
salts nitric acid can be obtained.
[Illustration Fig. 37]
~Preparation of nitric acid.~ When sodium nitrate is treated with
concentrated cold sulphuric acid, no chemical action seems to take
place. If, however, the mixture is heated in a retort, nitric acid is
given off as a vapor and may be easily condensed to a liquid by passing
the vapor into a tube surrounded by cold water, as shown in Fig. 37. An
examination of the liquid left in the retort shows that it contains
sodium acid sulphate (NaHSO_{4}), so that the reaction may be
represented by the equation
NaNO_{3} + H_{2}SO_{4} = NaHSO_{4} + HNO_{3}.
If a smaller quantity of sulphuric acid is taken and the
mixture is heated to a high temperature, normal sodium sulphate
is formed:
2NaNO_{3} + H_{2}SO_{4} = Na_{2}SO_{4} + 2HNO_{3}.
In this case, however, the higher temperature required
decomposes a part of the nitric acid.
~The commercial preparation of nitric acid.~ Fig. 38 illustrates
a form of apparatus used in the preparation of nitric acid on a
large scale. Sodium nitrate and sulphuric acid are heated in
the iron retort A. The resulting acid vapors pass in the
direction indicated by the arrows, and are condensed in the
glass tubes B, which are covered with cloth kept cool by
streams of water. These tubes are inclined so that the liquid
resulting from the condensation of the vapors runs back into
C and is drawn off into large vessels (D).
[Illustration Fig. 38]
~Physical properties of nitric acid.~ Pure nitric acid is a colorless
liquid, which boils at about 86° and has a density of 1.56. The
concentrated acid of commerce contains about 68% of the acid, the
remainder being water. Such a mixture has a density of 1.4. The
concentrated acid fumes somewhat in moist air, and has a sharp choking
odor.
~Chemical properties.~ The most important chemical properties of nitric
acid are the following.
1. _Acid properties._ As the name indicates, this substance is an acid,
and has all the properties of that class of substances. It changes blue
litmus red and has a sour taste in dilute solutions. It forms hydrogen
ions in solution and neutralizes bases forming salts. It also acts upon
the oxides of most metals, forming a salt and water. It is one of the
strongest acids.
2. _Decomposition on heating._ When boiled, or exposed for some time to
sunlight, it suffers a partial decomposition according to the equation
2HNO_{3} = H_{2}O + 2NO_{2} + O.
The substance NO_{2}, called nitrogen peroxide, is a brownish gas, which
is readily soluble in water and in nitric acid. It therefore dissolves
in the undecomposed acid, and imparts a yellowish or reddish color to
it. Concentrated nitric acid highly charged with this substance is
called _fuming nitric acid_.
3. _Oxidizing action._ According to its formula, nitric acid contains a
large percentage of oxygen, and the reaction just mentioned shows that
the compound is not a very stable one, easily undergoing decomposition.
These properties should make it a good oxidizing agent, and we find that
this is the case. Under ordinary circumstances, when acting as an
oxidizing agent, it is decomposed according to the equation
2HNO_{3} = H_{2}O + 2NO + 3O.
The oxygen is taken up by the substance oxidized, and not set free, as
is indicated in the equation. Thus, if carbon is oxidized by nitric
acid, the oxygen combines with carbon, forming carbon dioxide (CO_{2}):
C + 2O = CO_{2}.
4. _Action on metals._ We have seen that when an acid acts upon a metal
hydrogen is set free. Accordingly, when nitric acid acts upon a metal,
such as copper, we should expect the reaction to take place which is
expressed in the equation
Cu + 2HNO_{3} = Cu(NO_{3})_{2} + 2H.
This reaction does take place, but the hydrogen set free is immediately
oxidized to water by another portion of the nitric acid according to the
equation
HNO_{3} + 3H = 2H_{2}O + NO.
As these two equations are written, two atoms of hydrogen are given off
in the first equation, while three are used up in the second. In order
that the hydrogen may be equal in the two equations, we must multiply
the first by 3 and the second by 2. We shall then have
3Cu + 6HNO_{3} = 3Cu(NO_{3})_{2} + 6H,
2HNO_{3} + 6H = 4H_{2}O + 2NO.
The two equations may now be combined into one by adding the quantities
on each side of the equality sign, canceling the hydrogen which is given
off in the one reaction and used up in the other. We shall then have the
equation
3Cu + 8HNO_{3} = 3Cu(NO_{3})_{2} + 2NO + 4H_{2}O.
A number of other reactions may take place when nitric acid acts upon
metals, resulting in the formation of other oxides of nitrogen, free
nitrogen, or even ammonia. The reaction just given is, however, the
usual one.
~Importance of steps in a reaction.~ This complete equation has
the advantage of making it possible to calculate very easily
the proportions in which the various substances enter into the
reaction or are formed in it. It is unsatisfactory in that it
does not give full information about the way in which the
reaction takes place. For example, it does not suggest that
hydrogen is at first formed, and subsequently transformed into
water. It is always much more important to remember the steps
in a chemical reaction than to remember the equation expressing
the complete action; for if these steps in the reaction are
understood, the complete equation is easily obtained in the
manner just described.
~Salts of nitric acid,--nitrates.~ The salts of nitric acid are called
nitrates. Many of these salts will be described in the study of the
metals. They are all soluble in water, and when heated to a high
temperature undergo decomposition. In a few cases a nitrate on being
heated evolves oxygen, forming a nitrite:
NaNO_{3} = NaNO_{2} + O.
In other cases the decomposition goes further, and the metal is left as
oxide:
Cu(NO_{3})_{2} = CuO + 2NO_{2} + O.
~Nitrous acid~ (HNO_{2}). It is an easy matter to obtain sodium nitrite
(NaNO_{2}), as the reaction given on the previous page indicates.
Instead of merely heating the nitrate, it is better to heat it together
with a mild reducing agent, such as lead, when the reaction takes place
which is expressed by the equation
NaNO_{3} + Pb = PbO + NaNO_{2}.
When sodium nitrite is treated with an acid, such as sulphuric acid, it
is decomposed and nitrous acid is set free:
NaNO_{2} + H_{2}SO_{4} = NaHSO_{4} + HNO_{2}.
The acid is very unstable, however, and decomposes readily into water
and nitrogen trioxide (N_{2}O_{3}):
2HNO_{2} = H_{2}O + N_{2}O_{3}.
Dilute solutions of the acid, however, can be obtained.
COMPOUNDS OF NITROGEN WITH OXYGEN
Nitrogen combines with oxygen to form five different oxides. The
formulas and names of these are as follows:
N_{2}O nitrous oxide.
NO nitric oxide.
NO_{2} nitrogen peroxide.
N_{2}O_{3} nitrogen trioxide, or nitrous anhydride.
N_{2}O_{5} nitrogen pentoxide, or nitric anhydride.
These will now be briefly discussed.
~Nitrous oxide~ (_laughing gas_) (N_{2}O). Ammonium nitrate, like all
nitrates, undergoes decomposition when heated; and owing to the fact
that it contains no metal, but does contain both oxygen and hydrogen,
the reaction is a peculiar one. It is represented by the equation
NH_{4}NO_{3} = 2H_{2}O + N_{2}O.
The oxide of nitrogen so formed is called nitrous oxide or laughing gas.
It is a colorless gas having a slight odor. It is somewhat soluble in
water, and in solution has a slightly sweetish taste. It is easily
converted into a liquid and can be purchased in this form. When inhaled
it produces a kind of hysteria (hence the name "laughing gas"), and even
unconsciousness and insensibility to pain if taken in large amounts. It
has long been used as an anæsthetic for minor surgical operations, such
as those of dentistry, but owing to its unpleasant after effects it is
not so much in use now as formerly.
Chemically, nitrous oxide is remarkable for the fact that it is a very
energetic oxidizing agent. Substances such as carbon, sulphur, iron, and
phosphorus burn in it almost as brilliantly as in oxygen, forming oxides
and setting free nitrogen. Evidently the oxygen in nitrous oxide cannot
be held in very firm combination by the nitrogen.
[Illustration Fig. 39]
~Nitric oxide~ (NO). We have seen that when nitric acid acts upon metals,
such as copper, the reaction represented by the following equation takes
place:
3Cu + 8HNO_{3} = 3Cu(NO_{3})_{3} + 2NO + 4H_{2}O.
Nitric oxide is most conveniently prepared in this way. The metal is
placed in the flask A (Fig. 39) and the acid added slowly through the
funnel tube B. The gas escapes through C and is collected over
water.
Pure nitric oxide is a colorless gas, slightly heavier than air, and is
practically insoluble in water. It is a difficult gas to liquefy. Unlike
nitrous oxide, nitric oxide does not part with its oxygen easily, and
burning substances introduced into this gas are usually extinguished. A
few substances like phosphorus, which have a very strong affinity for
oxygen and which are burning energetically in the air, will continue to
burn in an atmosphere of nitric oxide. In this case the nitric oxide
loses all of its oxygen and the nitrogen is set free as gas.
~Action of nitric oxide with oxygen.~ When nitric oxide comes into contact
with oxygen or with the air, it at once combines with the oxygen even at
ordinary temperatures, forming a reddish-yellow gas of the formula
NO_{2}, which is called nitrogen peroxide. This action is not energetic
enough to produce a flame, though considerable heat is set free.
~Nitrogen peroxide~ (NO_{2}). This gas, as we have just seen, is formed by
allowing nitric oxide to come into contact with oxygen. It can also be
made by heating certain nitrates, such as lead nitrate:
Pb(NO_{3})_{2} = PbO + 2NO_{2} + O.
It is a reddish-yellow gas of unpleasant odor, which is quite poisonous
when inhaled. It is heavier than air and is easily condensed to a
liquid. It dissolves in water, but this solution is not a mere physical
solution; the nitrogen peroxide is decomposed, forming a mixture of
nitric and nitrous acids:
2NO_{2} + H_{2}O = HNO_{2} + HNO_{3}.
Nitrogen peroxide will not combine with more oxygen; it will, however,
give up a part of its oxygen to burning substances, acting as an
oxidizing agent:
NO_{2} = NO + O.
~Acid anhydrides.~ The oxides N_{2}O_{3} (nitrogen trioxide) and
N_{2}O_{5} (nitrogen pentoxide) are rarely prepared and need not be
separately described. They bear a very interesting relation to the acids
of nitrogen. When dissolved in water they combine with the water,
forming acids:
N_{2}O_{3} + H_{2}O = 2HNO_{2},
N_{2}O_{5} + H_{2}O = 2HNO_{3}.
On the other hand, nitrous acid very easily decomposes, yielding water
and nitrogen trioxide, and by suitable means nitric acid likewise may be
decomposed into water and nitrogen pentoxide:
2HNO_{2} = H_{2}O + N_{2}O_{3},
2HNO_{3} = H_{2}O + N_{2}O_{5}.
In view of the close relation between these oxides and the corresponding
acids, they are called _anhydrides_ of the acids, N_{2}O_{3} being
nitrous anhydride and N_{2}O_{5} nitric anhydride.
DEFINITION: _Any oxide which will combine with water to form an acid, or
which together with water is formed by the decomposition of an acid, is
called an anhydride of that acid._
EXERCISES
1. Perfectly dry ammonia does not affect litmus paper. Explain.
2. Can ammonia be dried by passing the gas through concentrated
sulphuric acid? Explain.
3. Ammonium hydroxide is a weak base, i.e. it is not highly dissociated.
When it is neutralized by strong acids the heat of reaction is less than
when strong bases are so neutralized. Suggest some possible cause for
this.
4. Why is brine used in the manufacture of artificial ice?
5. Discuss the energy changes which take place in the manufacture of
artificial ice.
6. What weight of ammonium chloride is necessary to furnish enough
ammonia to saturate 1 l. of water at 0° and 760 mm.?
7. What weight of sodium nitrate is necessary to prepare 100 cc. of
commercial nitric acid? What weight of potassium nitrate is necessary to
furnish the same weight of acid?
8. 100 l. of nitrogen peroxide were dissolved in water and neutralized
with sodium hydroxide. What substances were formed and how much of
each?(1 l. nitrogen peroxide weighs 2.05 grams.)
9. How many liters of nitrous oxide, measured under standard conditions,
can be prepared from 10 g. of ammonium nitrate?
10. What weight of copper is necessary to prepare 50 l. of nitric oxide
under standard conditions?
11. (a) Calculate the percentage composition of the oxides of
nitrogen. (b) What important law does this series of substances
illustrate?
12. Write the equations representing the reactions between ammonium
hydroxide, and sulphuric acid and nitric acid respectively, in
accordance with the theory of electrolytic dissociation.
13. In the same way, write the equations representing the reactions
between nitric acid and each of the following bases: NaOH, KOH,
NH_{4}OH, Ca(OH)_{2}.