THE ALKALI METALS
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| | | | |
| SYMBOL | ATOMIC | DENSITY | MELTING | FIRST PREPARED
| | WEIGHT | | POINT |
__________|________|________|_________|_________|________________
| | | | |
Lithium | Li | 7.03 | 0.59 | 186.° | Davy 1820
Sodium | Na | 23.05 | 0.97 | 97.6° | " 1807
Potassium | K | 39.15 | 0.87 | 62.5° | " 1807
Rubidium | Rb | 85.5 | 1.52 | 38.5° | Bunsen 1861
Cæsium | Cs | 132.9 | 1.88 | 26.5° | " 1860
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~The family.~ The metals listed in the above table constitute the even
family in Group I in the periodic arrangement of the elements, and
therefore form a natural family. The name alkali metals is commonly
applied to the family for the reason that the hydroxides of the most
familiar members of the family, namely sodium and potassium, have long
been called alkalis.
1. _Occurrence._ While none of these metals occur free in nature, their
compounds are very widely distributed, being especially abundant in sea
and mineral waters, in salt beds, and in many rocks. Only sodium and
potassium occur in abundance, the others being rarely found in any
considerable quantity.
2. _Preparation._ The metals are most conveniently prepared by the
electrolysis of their fused hydroxides or chlorides, though it is
possible to prepare them by reducing their oxides or carbonates with
carbon.
3. _Properties._ They are soft, light metals, having low melting points
and small densities, as is indicated in the table. Their melting points
vary inversely with their atomic weights, while their densities (sodium
excepted) vary directly with these. The pure metals have a silvery
luster but tarnish at once when exposed to the air, owing to the
formation of a film of oxide upon the surface of the metal. They are
therefore preserved in some liquid, such as coal oil, which contains no
oxygen. Because of their strong affinity for oxygen they decompose water
with great ease, forming hydroxides and liberating hydrogen in
accordance with the equation
M + H_{2}O = MOH + H,
where M stands for any one of these metals. These hydroxides are white
solids; they are readily soluble in water and possess very strong basic
properties. These bases are nearly equal in strength, that is, they all
dissociate in water to about the same extent.
4. _Compounds._ The alkali metals almost always act as univalent
elements in the formation of compounds, the composition of which can be
represented by such formulas as MH, MCl, MNO_{3}, M_{2}SO_{4},
M_{3}PO_{4}. These compounds, when dissolved in water, dissociate in
such a way as to form simple, univalent metallic ions which are
colorless. With the exception of lithium these metals form very few
insoluble compounds, so that it is not often that precipitates
containing them are obtained. Only sodium and potassium will be studied
in detail, since the other metals of the family are of relatively small
importance.
The compounds of sodium and potassium are so similar in properties that
they can be used interchangeably for most purposes. Other things being
equal, the sodium compounds are prepared in preference to those of
potassium, since they are cheaper. When a given sodium compound is
deliquescent, or is so soluble that it is difficult to purify, the
corresponding potassium compound is prepared in its stead, provided its
properties are more desirable in these respects.
SODIUM
~Occurrence in nature.~ Large deposits of sodium chloride have been found
in various parts of the world, and the water of the ocean and of many
lakes and springs contains notable quantities of it. The element also
occurs as a constituent of many rocks and is therefore present in the
soil formed by their disintegration. The mineral cryolite
(Na_{3}AlF_{6}) is an important substance, and the nitrate, carbonate,
and borate also occur in nature.
~Preparation.~ In 1807 Sir Humphry Davy succeeded in preparing very small
quantities of metallic sodium by the electrolysis of the fused
hydroxide. On account of the cost of electrical energy it was for many
years found more economical to prepare it by reducing the carbonate with
carbon in accordance with the following equation:
Na_{2}CO_{3} + 2C = 2Na + 3CO.
The cost of generating the electric current has been diminished to such
an extent, however, that it is now more economical to prepare sodium by
Davy's original method, namely, by the electrolysis of the fused
hydroxide or chloride. When the chloride is used the process is
difficult to manage, owing to the higher temperature required to keep
the electrolyte fused, and because of the corroding action of the fused
chloride upon the containing vessel.
[Illustration: SIR HUMPHRY DAVY (English) (1778-1829)
Isolated sodium, lithium, potassium, barium, strontium, and calcium by
means of electrolysis; demonstrated the elementary nature of chlorine;
invented the safety lamp; discovered the stupefying effects of nitrous
oxide]
~Technical preparation.~ The sodium hydroxide is melted in a
cylindrical iron vessel (Fig. 76) through the bottom of which
rises the cathode K. The anodes A, several in number, are
suspended around the cathode from above. A cylindrical vessel
C floats in the fused alkali directly over the cathode, and
under this cap the sodium and hydrogen liberated at the cathode
collect. The hydrogen escapes by lifting the cover, and the
sodium, protected from the air by the hydrogen, is skimmed or
drained off from time to time. Oxygen is set free upon the
anode and escapes into the air through the openings O without
coming into contact with the sodium or hydrogen. This process
is carried on extensively at Niagara Falls.
[Illustration: Fig. 76]
~Properties.~ Sodium is a silver-white metal about as heavy as water, and
so soft that it can be molded easily by the fingers or pressed into
wire. It is very active chemically, combining with most of the
non-metallic elements, such as oxygen and chlorine, with great energy.
It will often withdraw these elements from combination with other
elements, and is thus able to decompose water and the oxides and
chlorides of many metals.
~Sodium peroxide~ (NaO). Since sodium is a univalent element we should
expect it to form an oxide of the formula Na_{2}O. While such an oxide
can be prepared, the peroxide (NaO) is much better known. It is a
yellowish-white powder made by burning sodium in air. Its chief use is
as an oxidizing agent. When heated with oxidizable substances it gives
up a part of its oxygen, as shown in the equation
2NaO = Na_{2}O + O.
Water decomposes it in accordance with the equation
2NaO + 2H_{2}O = 2NaOH + H_{2}O_{2}.
Acids act readily upon it, forming a sodium salt and hydrogen peroxide:
2NaO + 2HCl = 2NaCl + H_{2}O_{2}.
In these last two reactions the hydrogen dioxide formed may decompose
into water and oxygen if the temperature is allowed to rise:
H_{2}O_{2} = H_{2}O + O.
~Peroxides.~ It will be remembered that barium dioxide (BaO_{2})
yields hydrogen dioxide when treated with acids, and that
manganese dioxide gives up oxygen when heated with sulphuric
acid. Oxides which yield either hydrogen dioxide or oxygen when
treated with water or an acid are called peroxides.
~Sodium hydroxide~ (_caustic soda_) (NaOH). 1. _Preparation._ Sodium
hydroxide is prepared commercially by several processes.
(a) In the older process, still in extensive use, sodium carbonate is
treated with calcium hydroxide suspended in water. Calcium carbonate is
precipitated according to the equation
Na_{2}CO_{3} + Ca(OH)_{2} = CaCO_{3} + 2NaOH.
The dilute solution of sodium hydroxide, filtered from the calcium
carbonate, is evaporated to a paste and is then poured into molds to
solidify. It is sold in the form of slender sticks.
(b) The newer methods depend upon the electrolysis of sodium chloride.
In the Castner process a solution of salt is electrolyzed, the reaction
being expressed as follows:
NaCl + H_{2}O = NaOH + H + Cl.
The chlorine escapes as a gas, and by an ingenious mechanical device the
sodium hydroxide is prevented from mixing with the salt in the solution.
In the Acker process the electrolyte is _fused_ sodium chloride. The
chlorine is evolved as a gas at the anode, while the sodium alloys with
the melted lead which forms the cathode. When this alloy is treated with
water the following reaction takes place:
Na + H_{2}O = NaOH + H.
[Illustration: Fig. 77]
~Technical process.~ A sketch of an Acker furnace is represented in Fig.
77. The furnace is an irregularly shaped cast-iron box, divided into
three compartments, A, B, and C. Compartment A is lined with
magnesia brick. Compartments B and C are filled with melted lead,
which also covers the bottom of A to a depth of about an inch. Above
this layer in A is fused salt, into which dip carbon anodes D. The
metallic box and melted lead is the cathode.
When the furnace is in operation chlorine is evolved at the
anodes, and is drawn away through a pipe (not represented) to
the bleaching-powder chambers. Sodium is set free at the
surface of the melted lead in A, and at once alloys with it.
Through the pipe E a powerful jet of steam is driven through
the lead in B upwards into the narrow tube F. This forces
the lead alloy up through the tube and over into the chamber
G.
In this process the steam is decomposed by the sodium in the
alloy, forming melted sodium hydroxide and hydrogen. The melted
lead and sodium hydroxide separate into two layers in G, and
the sodium hydroxide, being on top, overflows into tanks from
which it is drawn off and packed in metallic drums. The lead is
returned to the other compartments of the furnace by a pipe
leading from H to I. Compartment C serves merely as a
reservoir for excess of melted lead.
2. _Properties._ Sodium hydroxide is a white, crystalline, brittle
substance which rapidly absorbs water and carbon dioxide from the air.
As the name (caustic soda) indicates, it is a very corrosive substance,
having a disintegrating action on most animal and vegetable tissues. It
is a strong base. It is used in a great many chemical industries, and
under the name of lye is employed to a small extent as a cleansing agent
for household purposes.
~Sodium chloride~ (_common salt_) (NaCl). 1. _Preparation._ Sodium
chloride, or common salt, is very widely distributed in nature. Thick
strata, evidently deposited at one time by the evaporation of salt
water, are found in many places. In the United States the most important
localities for salt are New York, Michigan, Ohio, and Kansas. Sometimes
the salt is mined, especially if it is in the pure form called rock
salt. More frequently a strong brine is pumped from deep wells sunk into
the salt deposit, and is then evaporated in large pans until the salt
crystallizes out. The crystals are in the form of small cubes and
contain no water of crystallization; some water is, however, held in
cavities in the crystals and causes the salt to decrepitate when heated.
2. _Uses._ Since salt is so abundant in nature it forms the starting
point in the preparation of all compounds containing either sodium or
chlorine. This includes many substances of the highest importance to
civilization, such as soap, glass, hydrochloric acid, soda, and
bleaching powder. Enormous quantities of salt are therefore produced
each year. Small quantities are essential to the life of man and
animals. Pure salt does not absorb moisture; the fact that ordinary salt
becomes moist in air is not due to a property of the salt, but to
impurities commonly occurring in it, especially calcium and magnesium
chlorides.
~Sodium sulphate~ (_Glauber's salt_) (Na_{2}SO_{4}·10H_{2}O). This salt is
prepared by the action of sulphuric acid upon sodium chloride,
hydrochloric acid being formed at the same time:
2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl.
Some sodium sulphate is prepared by the reaction represented in the
equation
MgSO_{4} + 2NaCl = Na_{2}SO_{4} + MgCl_{2}.
The magnesium sulphate required for this reaction is obtained in large
quantities in the manufacture of potassium chloride, and being of little
value for any other purpose is used in this way. The reaction depends
upon the fact that sodium sulphate is the least soluble of any of the
four factors in the equation, and therefore crystallizes out when hot,
saturated solutions of magnesium sulphate and sodium chloride are mixed
together and the resulting mixture cooled.
Sodium sulphate forms large efflorescent crystals. The salt is
extensively used in the manufacture of sodium carbonate and glass. Small
quantities are used in medicine.
~Sodium sulphite~ (Na_{2}SO_{3}·7H_{2}O). Sodium sulphite is prepared by
the action of sulphur dioxide upon solutions of sodium hydroxide, the
reaction being analogous to the action of carbon dioxide upon sodium
hydroxide. Like the carbonate, the sulphite is readily decomposed by
acids:
Na_{2}SO_{3} + 2HCl = 2NaCl + H_{2}O + SO_{2}.
Because of this reaction sodium sulphite is used as a convenient source
of sulphur dioxide. It is also used as a disinfectant and a
preservative.
~Sodium thiosulphate~ (_hyposulphite of soda or "hypo"_)
(Na_{2}S_{2}O_{3}·5H_{2}O). This salt, commonly called sodium
hyposulphite, or merely hypo, is made by boiling a solution of sodium
sulphite with sulphur:
Na_{2}SO_{3} + S = Na_{2}S_{2}O_{3}.
It is used in photography and in the bleaching industry, to absorb the
excess of chlorine which is left upon the bleached fabrics.
~Thio compounds.~ The prefix "thio" means sulphur. It is used to
designate substances which may be regarded as derived from
oxygen compounds by replacing the whole or a part of their
oxygen with sulphur. The thiosulphates may be regarded as
sulphates in which one atom of oxygen has been replaced by an
atom of sulphur. This may be seen by comparing the formula
Na_{2}SO_{4} (sodium sulphate) with the formula
Na_{2}S_{2}O_{3} (sodium thiosulphate).
~Sodium carbonate~ (_sal soda_)(Na_{2}CO_{3}·10H_{2}O). There are two
different methods now employed in the manufacture of this important
substance.
1. _Le Blanc process._ This older process involves several distinct
reactions, as shown in the following equations.
(a) Sodium chloride is first converted into sodium sulphate:
2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + 2HCl.
(b) The sodium sulphate is next reduced to sulphide by heating it with
carbon:
Na_{2}SO_{4} + 2C = Na_{2}S + 2CO_{2}.
(c) The sodium sulphide is then heated with calcium carbonate, when
double decomposition takes place:
Na_{2}S + CaCO_{3} = CaS + Na_{2}CO_{3}.
~Technical preparation of sodium carbonate.~ In a manufacturing
plant the last two reactions take place in one process. Sodium
sulphate, coal, and powdered limestone are heated together to a
rather high temperature. The coal reduces the sulphate to
sulphide, which in turn reacts upon the calcium carbonate. Some
limestone is decomposed by the heat, forming calcium oxide.
When treated with water the calcium oxide is changed into
hydroxide, and this prevents the water from decomposing the
insoluble calcium sulphide.
The crude product of the process is a hard black cake called
black ash. On digesting this mass with water the sodium
carbonate passes into solution. The pure carbonate is obtained
by evaporation of this solution, crystallizing from it in
crystals of the formula Na_{2}CO_{3}·10H_{2}O. Since over 60%
of this salt is water, the crystals are sometimes heated until
it is driven off. The product is called calcined soda, and is,
of course, more valuable than the crystallized salt.
2. _Solvay process._ This more modern process depends upon the reactions
represented in the equations
NaCl + NH_{4}HCO_{3} = NaHCO_{3} + NH_{4}Cl,
2NaHCO_{3} = Na_{2}CO_{3} + H_{2}O + CO_{2}.
The reason the first reaction takes place is that sodium hydrogen
carbonate is sparingly soluble in water, while the other compounds are
freely soluble. When strong solutions of sodium chloride and of ammonium
hydrogen carbonate are brought together the sparingly soluble sodium
hydrogen carbonate is precipitated. This is converted into the normal
carbonate by heating, the reaction being represented in the second
equation.
~Technical preparation.~ In the Solvay process a very
concentrated solution of salt is first saturated with ammonia
gas, and a current of carbon dioxide is then conducted into the
solution. In this way ammonium hydrogen carbonate is formed:
NH_{3} + H_{2}O + CO_{2} = NH_{4}HCO_{3}.
This enters into double decomposition with the salt, as shown
in the first equation under the Solvay process. After the
sodium hydrogen carbonate has been precipitated the mother
liquors containing ammonium chloride are treated with lime:
2NH_{4}Cl + CaO = CaCl_{2} + 2 NH_{3} + H_{2}O.
The lime is obtained by burning limestone:
CaCO_{3} = CaO + CO_{2}.
The ammonia and carbon dioxide evolved in the latter two
reactions are used in the preparation of an additional quantity
of ammonium hydrogen carbonate. It will thus be seen that there
is no loss of ammonia. The only materials permanently used up
are calcium carbonate and salt, while the only waste product is
calcium chloride.
~Historical.~ In former times sodium carbonate was made by
burning seaweeds and extracting the carbonate from their ash.
On this account the salt was called _soda ash_, and the name is
still in common use. During the French Revolution this supply
was cut off, and in behalf of the French government Le Blanc
made a study of methods of preparing the carbonate directly
from salt. As a result he devised the method which bears his
name, and which was used exclusively for many years. It has
been replaced to a large extent by the Solvay process, which
has the advantage that the materials used are inexpensive, and
that the ammonium hydrogen carbonate used can be regenerated
from the products formed in the process. Much expense is also
saved in fuel, and the sodium hydrogen carbonate, which is the
first product of the process, has itself many commercial uses.
The Le Blanc process is still used, however, since the
hydrochloric acid generated is of value.
~By-products.~ The substances obtained in a given process, aside
from the main product, are called the by-products. The success
of many processes depends upon the value of the by-products
formed.
Thus hydrochloric acid, a by-product in the Le Blanc process,
is valuable enough to make the process pay, even though sodium
carbonate can be made cheaper in other ways.
~Properties of sodium carbonate.~ Sodium carbonate forms large crystals of
the formula Na_{2}CO_{3} · 10 H_{2}O. It has a mild alkaline reaction
and is used for laundry purposes under the name of washing soda. Mere
mention of the fact that it is used in the manufacture of glass, soap,
and many chemical reagents will indicate its importance in the
industries. It is one of the few soluble carbonates.
~Sodium hydrogen carbonate~ (_bicarbonate of soda_) (NaHCO_{3}). This
salt, commonly called bicarbonate of soda, or baking soda, is made by
the Solvay process, as explained above, or by passing carbon dioxide
into strong solutions of sodium carbonate:
Na_{2}CO_{3} + H_{2}O + CO_{2} = 2NaHCO_{3}.
The bicarbonate, being sparingly soluble, crystallizes out. A mixture of
the bicarbonate with some substance (the compound known as cream of
tartar is generally used) which slowly reacts with it, liberating carbon
dioxide, is used largely in baking. The carbon dioxide generated forces
its way through the dough, thus making it porous and light.
~Sodium nitrate~ (_Chili saltpeter_) (NaNO_{3}). This substance is found
in nature in arid regions in a number of places, where it has been
formed apparently by the decay of organic substances in the presence of
air and sodium salts. The largest deposits are in Chili, and most of the
nitrate of commerce comes from that country. Smaller deposits occur in
California and Nevada. The commercial salt is prepared by dissolving the
crude nitrate in water, allowing the insoluble earthy materials to
settle, and evaporating the clear solution so obtained to
crystallization. The soluble impurities remain for the most part in the
mother liquors.
Since this salt is the only nitrate found extensively in nature, it is
the material from which other nitrates as well as nitric acid are
prepared. It is used in enormous quantities in the manufacture of
sulphuric acid and potassium nitrate, and as a fertilizer.
~Sodium phosphate~ (Na_{2}HPO_{4}·12H_{2}O). Since phosphoric acid has
three replaceable hydrogen atoms, three sodium phosphates are
possible,--two acid salts and one normal. All three can be made without
difficulty, but disodium phosphate is the only one which is largely
used, and is the salt which is commonly called sodium phosphate. It is
made by the action of phosphoric acid on sodium carbonate:
Na_{2}CO_{3} + H_{3}PO_{4} = Na_{2}HPO_{4} + CO_{2} + H_{2}O.
It is interesting as being one of the few phosphates which are soluble
in water, and is the salt commonly used when a soluble phosphate is
needed.
~Normal sodium phosphate~ (Na_{3}PO_{4}). Although this is a normal salt
its solution has a strongly alkaline reaction. This is due to the fact
that the salt hydrolyzes in solution into sodium hydroxide and disodium
phosphate, as represented in the equation
Na_{3}PO_{4} + H_{2}O = Na_{2}HPO_{4} + NaOH.
Sodium hydroxide is strongly alkaline, while disodium phosphate is
nearly neutral in reaction. The solution as a whole is therefore
alkaline. The salt is prepared by adding a large excess of sodium
hydroxide to a solution of disodium phosphate and evaporating to
crystallization. The excess of the sodium hydroxide reverses the
reaction of hydrolysis and the normal salt crystallizes out.
~Sodium tetraborate ~(_borax_) (Na_{2}B_{4}O_{7}·10H_{2}O). The properties
of this important compound have been discussed under the head of boron.
POTASSIUM
~Occurrence in nature.~ Potassium is a constituent of many common rocks
and minerals, and is therefore a rather abundant element, though not so
abundant as sodium. Feldspar, which occurs both by itself and as a
constituent of granite, contains considerable potassium. The element is
a constituent of all clay and of mica and also occurs in very large
deposits at Stassfurt, Germany, in the form of the chloride and
sulphate, associated with compounds of sodium and magnesium. In small
quantities it is found as nitrate and in many other forms.
The natural decomposition of rocks containing potassium gives rise to
various compounds of the element in all fertile soils. Its soluble
compounds are absorbed by growing plants and built up into complex
vegetable substances; when these are burned the potassium remains in the
ash in the form of the carbonate. Crude carbonate obtained from wood
ashes was formerly the chief source of potassium compounds; they are now
mostly prepared from the salts of the Stassfurt deposits.
~Stassfurt salts.~ These salts form very extensive deposits in
middle and north Germany, the most noted locality for working
them being at Stassfurt. The deposits are very thick and rest
upon an enormous layer of common salt. They are in the form of
a series of strata, each consisting largely of a single mineral
salt. A cross section of these deposits is shown in Fig. 78.
While these strata are salts from a chemical standpoint, they
are as solid and hard as many kinds of stone, and are mined as
stone or coal would be. Since the strata differ in general
appearance, each can be mined separately, and the various
minerals can be worked up by methods adapted to each particular
case. The chief minerals of commercial importance in these
deposits are the following:
Sylvine KCl.
Anhydrite CaSO_{4}.
Carnallite KCl·MgCl_{2}·6H_{2}O.
Kainite K_{2}SO_{4}·MgSO_{4}·MgCl_{2}·6H_{2}O.
Polyhalite K_{2}SO_{4}·MgSO_{4}·2CaSO_{4}·2H_{2}O.
Kieserite MgSO_{4}·H_{2}O.
Schönite K_{2}SO_{4}·MgSO_{4}·6H_{2}O.
~Preparation and properties.~ The metal is prepared by the same method
used in the preparation of sodium. In most respects it is very similar
to sodium, the chief difference being that it is even more energetic in
its action upon other substances. The freshly cut, bright surface
instantly becomes dim through oxidation by the air. It decomposes water
very vigorously, the heat of reaction being sufficient to ignite the
hydrogen evolved. It is somewhat lighter than sodium and is preserved
under gasoline.
[Illustration: Fig. 78]
~Potassium hydroxide~ (_caustic potash_) (KOH). Potassium hydroxide is
prepared by methods exactly similar to those used in the preparation of
sodium hydroxide, which compound it closely resembles in both physical
and chemical properties. It is not used to any very great extent, being
replaced by the cheaper sodium hydroxide.
~Action of the halogen elements on potassium hydroxide.~ When any one of
the three halogen elements--chlorine, bromine, and iodine--is added to a
solution of potassium hydroxide a reaction takes place, the nature of
which depends upon the conditions of the experiment. Thus, when chlorine
is passed into a cold dilute solution of potassium hydroxide the
reaction expressed by the following equation takes place:
(1) 2KOH + 2Cl = KCl + KClO + H_{2}O.
If the solution of hydroxide is concentrated and hot, on the other hand,
the potassium hypochlorite formed according to equation (1) breaks down
as fast as formed:
(2) 3KClO = KClO_{3} + 2KCl.
Equation (1), after being multiplied by 3, may be combined with equation
(2), giving the following:
(3) 6KOH + 6Cl = 5KCl + KClO_{3} + 3H_{2}O.
This represents in a single equation the action of chlorine on hot,
concentrated solutions of potassium hydroxide. By means of these
reactions one can prepare potassium chloride, potassium hypochlorite,
and potassium chlorate. By substituting bromine or iodine for chlorine
the corresponding compounds of these elements are obtained. Some of
these compounds can be obtained in cheaper ways.
If the halogen element is added to a solution of sodium hydroxide or
calcium hydroxide, the reaction which takes place is exactly similar to
that which takes place with potassium hydroxide. It is possible,
therefore, to prepare in this way the sodium and calcium compounds
corresponding to the potassium compounds given above.
~Potassium chloride~ (KCl). This salt occurs in nature in sea water, in
the mineral sylvine, and, combined with magnesium chloride, as
carnallite (KCl·MgCl_{2}·6H_{2}O). It is prepared from carnallite by
saturating boiling water with the mineral and allowing the solution to
cool. The mineral decomposes while in solution, and the potassium
chloride crystallizes out on cooling, while the very soluble magnesium
chloride remains in solution. The salt is very similar to sodium
chloride both in physical and chemical properties. It is used in the
preparation of nearly all other potassium salts, and, together with
potassium sulphate, is used as a fertilizer.
~Potassium bromide~ (KBr). When bromine is added to a hot concentrated
solution of potassium hydroxide there is formed a mixture of potassium
bromide and potassium bromate in accordance with the reactions already
discussed. There is no special use for the bromate, so the solution is
evaporated to dryness, and the residue, consisting of a mixture of the
bromate and bromide, is strongly heated. This changes the bromate to
bromide, as follows:
KBrO_{3} = KBr +3O.
The bromide is then crystallized from water, forming large colorless
crystals. It is used in medicine and in photography.
~Potassium iodide~ (KI). Potassium iodide may be made by exactly the same
method as has just been described for the bromide, substituting iodine
for bromine. It is more frequently made as follows. Iron filings are
treated with iodine, forming the compound Fe_{3}I_{8}; on boiling this
substance with potassium carbonate the reaction represented in the
following equation occurs:
Fe_{3}I_{8} + 4K_{2}CO_{3} = Fe_{3}O_{4} + 8KI + 4CO_{2}.
Potassium iodide finds its chief use in medicine.
~Potassium chlorate~ (KClO_{3}). This salt, as has just been explained,
can be made by the action of chlorine on strong potassium hydroxide
solutions. The chief use of potassium chlorate is as an oxidizing agent
in the manufacture of matches, fireworks, and explosives; it is also
used in the preparation of oxygen and in medicine.
~Commercial preparation.~ By referring to the reaction between
chlorine and hot concentrated solutions of potassium hydroxide,
it will be seen that only one molecule of potassium chlorate is
formed from six molecules of potassium hydroxide. Partly
because of this poor yield and partly because the potassium
hydroxide is rather expensive, this process is not an
economical one for the preparation of potassium chlorate. The
commercial method is the following. Chlorine is passed into hot
solutions of calcium hydroxide, a compound which is very cheap.
The resulting calcium chloride and chlorate are both very
soluble. To the solution of these salts potassium chloride is
added, and as the solution cools the sparingly soluble
potassium chlorate crystallizes out:
Ca(ClO_{3})_{2} + 2KCl = 2KClO_{3} + CaCl_{2}.
Electro-chemical processes are also used.
~Potassium nitrate~ (_saltpeter_) (KNO_{3}). This salt was formerly made
by allowing animal refuse to decompose in the open air in the presence
of wood ashes or earthy materials containing potassium. Under these
conditions the nitrogen in the organic matter is in part converted into
potassium nitrate, which was obtained by extracting the mass with water
and evaporating to crystallization. This crude and slow process is now
almost entirely replaced by a manufacturing process in which the
potassium salt is made from Chili saltpeter:
NaNO_{3} + KCl = NaCl + KNO_{3}.
This process has been made possible by the discovery of the Chili niter
beds and the potassium chloride of the Stassfurt deposits.
The reaction depends for its success upon the apparently
insignificant fact that sodium chloride is almost equally
soluble in cold and hot water. All four factors in the equation
are rather soluble in cold water, but in hot water sodium
chloride is far less soluble than the other three. When hot
saturated solutions of sodium nitrate and potassium chloride
are brought together, sodium chloride precipitates and can be
filtered off, leaving potassium nitrate in solution, together
with some sodium chloride. On cooling, potassium nitrate
crystallizes out, leaving small amounts of the other salts in
solution.
Potassium nitrate is a colorless salt which forms very large crystals.
It is stable in the air, and when heated is a good oxidizing agent,
giving up oxygen quite readily. Its chief use is in the manufacture of
gunpowder.
~Gunpowder.~ The object sought for in the preparation of
gunpowder is to secure a solid substance which will remain
unchanged under ordinary conditions, but which will explode
readily when ignited, evolving a large volume of gas. When a
mixture of carbon and potassium nitrate is ignited a great deal
of gas is formed, as will be seen from the equation
2KNO_{3} + 3C = CO_{2} + CO + N_{2} + K_{2}CO_{3}.
By adding sulphur to the mixture the volume of gas formed in
the explosion is considerably increased:
2KNO_{3} + 3C + S = 3CO_{2} + N_{2} + K_{2}S.
Gunpowder is simply a mechanical mixture of these three
substances in the proportion required for the above reaction.
While the equation represents the principal reaction, other
reactions also take place. The gases formed in the explosion,
when measured under standard conditions, occupy about two
hundred and eighty times the volume of the original powder.
Potassium sulphide (K_{2}S) is a solid substance, and it is
largely due to it that gunpowder gives off smoke and soot when
it explodes. Smokeless powder consists of organic substances
which, on explosion, give only colorless gases, and hence
produce no smoke. Sodium nitrate is cheaper than potassium
nitrate, but it is not adapted to the manufacture of the best
grades of powder, since it is somewhat deliquescent and does
not give up its oxygen so readily as does potassium nitrate. It
is used, however, in the cheaper grades of powder, such as are
employed for blasting.
~Potassium cyanide~ (KCN). When animal matter containing nitrogen is
heated with iron and potassium carbonate, complicated changes occur
which result in the formation of a substance commonly called yellow
prussiate of potash, which has the formula K_{4}FeC_{6}N_{6}. When this
substance is heated with potassium, potassium cyanide is formed:
K_{4}FeC_{6}N_{6} + 2 K = 6KCN + Fe.
Since sodium is much cheaper than potassium it is often used in place of
it:
K_{4}FeC_{6}N_{6} + 2Na = 4KCN + 2NaCN + Fe.
The mixture of cyanides so resulting serves most of the purposes of the
pure salt. It is used very extensively in several metallurgical
processes, particularly in the extraction of gold. Potassium cyanide is
a white solid characterized by its poisonous properties, and must be
used with extreme caution.
~Potassium carbonate~ (_potash_) (K_{2}CO_{3}). This compound occurs in
wood ashes in small quantities. It cannot be prepared by the Solvay
process, since the acid carbonate is quite soluble in water, but is made
by the Le Blanc process. Its chief use is in the manufacture of other
potassium salts.
~Other salts of potassium.~ Among the other salts of potassium frequently
met with are the sulphate (K_{2}SO_{4}), the acid carbonate (KHCO_{3}),
the acid sulphate (KHSO_{4}), and the acid sulphite (KHSO_{3}). These
are all white solids.
LITHIUM, RUBIDIUM, CÆSIUM
Of the three remaining elements of the family--lithium, rubidium, and
cæsium--lithium is by far the most common, the other two being very
rare. Lithium chloride and carbonate are not infrequently found in
natural mineral waters, and as these substances are supposed to increase
the medicinal value of the water, they are very often added to
artificial mineral waters in small quantities.
COMPOUNDS OF AMMONIUM
~General.~ As explained in a previous chapter, when ammonia is passed into
water the two compounds combine to form the base NH_{4}OH, known as
ammonium hydroxide. When this base is neutralized with acids there are
formed the corresponding salts, known as the ammonium salts. Since the
ammonium group is univalent, ammonium salts resemble those of the alkali
metals in formulas; they also resemble the latter salts very much in
their chemical properties, and may be conveniently described in
connection with them. Among the ammonium salts the chloride, sulphate,
carbonate, and sulphide are the most familiar.
~Ammonium chloride~ (_sal ammoniac_) (NH_{4}Cl). This substance is
obtained by neutralizing ammonium hydroxide with hydrochloric acid. It
is a colorless substance crystallizing in fine needles, and, like most
ammonium salts, is very soluble in water. When placed in a tube and
heated strongly it decomposes into hydrochloric acid and ammonia. When
these gases reach a cooler portion of the tube they at once recombine,
and the resulting ammonium chloride is deposited on the sides of the
tube. In this way the salt can be separated from nonvolatile impurities.
Ammonium chloride is sometimes used in preparation of ammonia; it is
also used in making dry batteries and in the laboratory as a chemical
reagent.
~Ammonium sulphate~ ((NH_{4})_{2}SO_{4}). This salt resembles the chloride
very closely, and, being cheaper, is used in place of it when possible.
It is used in large quantity as a fertilizer, the nitrogen which it
contains being a very valuable food for plants.
~Ammonium carbonate~ ((NH_{4})_{2}CO_{3}). This salt, as well as the acid
carbonate (NH_{4}HCO_{3}), is used as a chemical reagent. They are
colorless solids, freely soluble in water. The normal carbonate is made
by heating ammonium chloride with powdered limestone (calcium
carbonate), the ammonium carbonate being obtained as a sublimate in
compact hard masses:
2NH_{4}Cl + CaCO_{3} = (NH_{4})_{2}CO_{3} + CaCl_{2}.
The salt always smells of ammonia, since it slowly decomposes, as shown
in the equation
(NH_{4})_{2}CO_{3} = NH_{4}HCO_{3} + NH_{3}.
The acid carbonate, or bicarbonate, is prepared by saturating a solution
of ammonium hydroxide with carbon dioxide:
NH_{4}OH + CO_{2} = NH_{4}HCO_{3}.
It is a well-crystallized stable substance.
~Ammonium sulphide~ ((NH_{4})_{2}S). Ammonium sulphide is prepared by the
action of hydrosulphuric acid upon ammonium hydroxide:
2NH_{4}OH + H_{2}S = (NH_{4})_{2}S + 2H_{2}O.
If the action is allowed to continue until no more hydrosulphuric acid
is absorbed, the product is the acid sulphide, sometimes called the
hydrosulphide:
NH_{4}OH + H_{2}S = NH_{4}HS + H_{2}O.
If equal amounts of ammonium hydroxide and ammonium acid sulphide are
brought together, the normal sulphide is formed:
NH_{4}OH + NH_{4}HS = (NH_{4})_{2}S + H_{2}O
It has been obtained in the solid state, but only with great difficulty.
As used in the laboratory it is always in the form of a solution. It is
much used in the process of chemical analysis because it is a soluble
sulphide and easily prepared. On exposure to the air ammonium sulphide
slowly decomposes, being converted into ammonia, water, and sulphur:
(NH_{4})_{2}S + O = 2NH_{3} + H_{2}O + S.
As fast as the sulphur is liberated it combines with the unchanged
sulphide to form several different ammonium sulphides in which there are
from two to five sulphur atoms in the molecule, thus: (NH_{4})_{2}S_{2},
(NH_{4})_{2}S_{3}, (NH_{4})_{2}S_{5}. These sulphides in turn decompose
by further action of oxygen, so that the final products of the reaction
are those given in the equation. A solution of these compounds is yellow
and is sometimes called _yellow ammonium sulphide_.
FLAME REACTION--SPECTROSCOPE
When compounds of either sodium or potassium are brought into
the non-luminous flame of a Bunsen burner the flame becomes
colored. Sodium compounds color it intensely yellow, while
those of potassium color it pale violet. When only one of these
elements is present it is easy to identify it by this simple
test, but when both are present the intense color of the sodium
flame entirely conceals the pale tint characteristic of
potassium compounds.
It is possible to detect the potassium flame in such cases,
however, in the following way. When light is allowed to shine
through a very small hole or slit in some kind of a screen,
such as a piece of metal, upon a triangular prism of glass, the
light is bent or refracted out of its course instead of passing
straight through the glass. It thus comes out of the prism at
some angle to the line at which it entered. Yellow light is
bent more than red, and violet more than yellow. When light
made up of the yellow of sodium and the violet of potassium
shines through a slit upon such a prism, the yellow and the
violet lights come out at somewhat different angles, and so two
colored lines of light--a yellow line and a violet line--are
seen on looking into the prism in the proper direction. The
instrument used for separating the rays of light in this way is
called a _spectroscope_ (Fig. 79). The material to be tested is
placed on a platinum wire and held in the colorless Bunsen
flame. The resulting light passes through the slit in the end
of tube B, and then through B to the prism. The resulting
lines of light are seen by looking into the tube A, which
contains a magnifying lens. Most elements give more than one
image of the slit, each having a different color, and the
series of colored lines due to an element is called its
spectrum.
[Illustration: Fig. 79]
The spectra of the known elements have been carefully studied, and any
element which imparts a characteristic color to a flame, or has a
spectrum of its own, can be identified even when other elements are
present. Through the spectroscopic examination of certain minerals a
number of elements have been discovered by the observation of lines
which did not belong to any known element. A study of the substance then
brought to light the new element. Rubidium and cæsium were discovered in
this way, rubidium having bright red lines and cæsium a very intense
blue line. Lithium colors the flame deep red, and has a bright red line
in its spectrum.
EXERCISES
1. What is an alkali? Can a metal itself be an alkali?
2. Write equations showing how the following changes may be brought
about, giving the general principle involved in each change: NaCl -->
Na_{2}SO_{3}, Na_{2}SO_{3} --> NaCl, NaCl --> NaBr, Na_{2}SO_{4} -->
NaNO_{3}, NaNO_{3} --> NaHCO_{3}.
3. What carbonates are soluble?
4. State the conditions under which the reaction represented by the
following equation can be made to go in either direction:
Na_{2}CO_{3} + H_{2}O + CO_{2} <--> 2 NaHCO_{3}.
5. Account for the fact that solutions of sodium carbonate and potassium
carbonate are alkaline.
6. What non-metallic element is obtained from the deposits of Chili
saltpeter?
7. Supposing concentrated hydrochloric acid (den. = 1.2) to be worth six
cents a pound, what is the value of the acid generated in the
preparation of 1 ton of sodium carbonate by the Le Blanc process?
8. What weight of sodium carbonate crystals will 1 kg. of the anhydrous
salt yield?
9. Write equations for the preparation of potassium hydroxide by three
different methods.
10. What would take place if a bit of potassium hydroxide were left
exposed to the air?
11. Write the equations for the reactions between sodium hydroxide and
bromine; between potassium hydroxide and iodine.
12. Write equations for the preparation of potassium sulphate; of
potassium acid carbonate.
13. What weight of carnallite would be necessary in the preparation of 1
ton of potassium carbonate?
14. Write the equations showing how ammonium chloride, ammonium
sulphate, ammonium carbonate, and ammonium nitrate may be prepared from
ammonium hydroxide.
15. Write an equation to represent the reaction involved in the
preparation of ammonia from ammonium chloride.
16. What substances already studied are prepared from the following
compounds? ammonium chloride; ammonium nitrate; ammonium nitrite; sodium
nitrate; sodium chloride.
17. How could you prove that the water in crystals of common salt is not
water of crystallization?
18. How could you distinguish between potassium chloride and potassium
iodide? between sodium chloride and ammonium chloride? between sodium
nitrate and potassium nitrate?
[Illustration: ROBERT WILHELM BUNSEN (German) (1811-1899)
Invented many lecture-room and laboratory appliances (Bunsen burner);
invented the spectroscope and with it discovered rubidium and cæsium;
greatly perfected methods of electrolysis, inventing a new battery; made
many investigations among metallic and organic substances]