salts, some chemical substances, are more soluble in water than others,
and that their solubilities under different circumstances of temperature
vary in different ways. However, some salts and compounds are
practically insoluble in water under any circumstances. We now arrive at
the important result known to chemists as the precipitation of insoluble
compounds from solutions. In order to define this result, however, we
must, of course, first consider the circumstances of causation of the
result. Let us take a simple case of chemical decomposition resulting in
the deposition or precipitation of a substance from solution in the
insoluble state. We will take a salt you are probably acquainted
with--sulphate of copper, or bluestone, and dissolve it in water, and we
have then the sulphate of copper in solution in water. Now suppose it is
our desire to obtain from that solution all the copper by depositing it
in some insoluble form. We may accomplish this in several different
ways, relying on certain methods of decomposing that sulphate of copper.
One of the simplest and most economical is that adopted in a certain
so-called wet method of extracting copper. It is based on the fact that
metallic iron has a greater tendency to combine in water solutions, with
the acids of copper salts, than the copper has in those salts. We
simply need to place some scraps of iron in the copper sulphate solution
to induce a change which may be represented as follows: Copper sulphate,
consisting of a combination of copper oxide with sulphuric acid, yields
with iron, iron sulphate, a combination of iron oxide with sulphuric
acid, and metallic copper. The metallic copper produced separates in the
form of a red coating on the iron scraps. But we may also, relying on
the fact that oxide of copper is insoluble in water, arrange for the
deposition of the copper in that form. This we can do by adding caustic
soda to a hot solution of copper sulphate, when we get the following
change: Copper sulphate, consisting of a combination of copper oxide
with sulphuric acid, yields with caustic soda, sulphate of soda, a
combination of soda with sulphuric acid and oxide of copper. Oxide of
copper is black, and so in this decomposition what is called a "black
precipitate" of that oxide is produced on adding the caustic soda. But
it might not suit us thus to deposit the copper from our solution; we
might desire to remove the sulphuric acid from the copper sulphate, and
leave the copper dissolved, say in the form of a chloride. We select,
then, a compound which is a chloride, and a chloride of a metal which
forms an insoluble combination with sulphuric acid--chloride of barium,
say. On adding this chloride of barium to sulphate of copper solution,
we get then a change which we might represent thus: Copper sulphate,
consisting of a combination of copper oxide with sulphuric acid, yields
with barium chloride, which is a combination of barium and chlorine,
insoluble barium sulphate, a combination of barium oxide with sulphuric
acid, and soluble copper chloride, a combination of copper and chlorine.
This is called a double interchange. Now these are a few illustrations
to show you what is meant by chemical decompositions. One practical
lesson, of course, we may draw is this: We must have a care in
dissolving bluestone or copper sulphate, not to attempt it in iron pans,
and not to store or put verdigris into iron vessels, or the iron will
be acted upon, and to some extent the copper salt will become
contaminated with iron. It will now be clear to you that, as a solvent
for bodies usually soluble in water, water that is perfectly pure will
be most suitable and not likely to cause any deposition or precipitation
through chemical decompositions, for there are no salts or other
compounds in pure water to cause such changes. Such pure water is called
soft water. But the term is only a comparative one, and water that is
not quite, but nearly pure--pure enough for most practical purposes--is
also called soft water. Now rain is the purest form of natural water,
for it is a kind of distilled water. Water rises in vapour from the
ocean as from a still, and the salt and other dissolved matters remain
behind. Meeting cold currents of air, the vapours condense in rain, and
fall upon the earth. After coming in contact with the earth, the
subsequent condition of that water entirely depends upon the character,
as regards solubility or insolubility, of the substances composing the
strata or layers of earth upon which it falls, and through which it
sinks. If it meets with insoluble rocks--for all rocks are not
insoluble--it remains, of course, pure and soft, and in proportion as
the constituents of rock and soil are soluble, in that proportion does
the water become hard. We all know how dangerous acid is in water,
causing that water to act on many substances, the iron of iron vessels,
the lime in soil or rock, etc., bringing iron and lime respectively into
solution. Now the atmosphere contains carbonic acid, and carbonic acid
occurs in the earth, being evolved by decomposing vegetation, etc.
Carbonic acid is also soluble to a certain, though not large extent, in
water. As we shall see, water charged with carbonic acid attacks certain
substances insoluble in pure water, and brings them into solution, and
thus the water soon becomes hard. About the close of the last lecture, I
said that lime is, to a certain extent, soluble in cold water. The
solution is called lime-water; it might be called a solution of caustic
lime. When carbonic acid gas first comes in contact with such a
solution, chalk or carbonate of lime, which is insoluble in water, is
formed, and the lime is thus precipitated as carbonate. Supposing,
however, we continued to pass carbonic acid gas into that water,
rendered milky with chalk powder, very soon the liquid would clear, and
we should get once more a solution of lime, but not caustic lime as it
was at first, simply now a solution of carbonate of lime in carbonic
acid, or a solution of bicarbonate of lime. I will take some lime-water,
and I will blow through; my breath contains carbonic acid, and you will
see the clear liquid become milky owing to separation of insoluble
carbonate of lime, or chalk. I now continue blowing, and at length that
chalk dissolves with the excess of carbonic acid, forming bicarbonate of
lime. This experiment explains how it is that water percolating through
or running over limestone strata dissolves out the insoluble chalk. Such
water, hard from dissolved carbonate of lime, can be softened by merely
boiling the water, for the excess of carbonic acid is then expelled, and
the chalk is precipitated again. This would be too costly for the
softening of large quantities of water, the boiling process consuming
too much coal, and so another process is adopted. Quicklime, or milk of
lime, is added to the water in the proper quantity. This lime unites
with the excess of carbonic acid holding chalk in solution, and forms
with it insoluble chalk, and so all deposits together as chalk. By this
liming process, also, the iron of the water dissolved likewise in
ferruginous streams, etc., by carbonic acid, would be precipitated. To
show this deposition I will now add some clear lime-water to the
solution I made of chalk with the carbonic acid of my breath, and a
precipitate is at once formed, all the lime and carbonic acid together
depositing as insoluble chalk. Hence clear lime-water forms a good test
for the presence of bicarbonates of lime or iron in a water. But water
may be hard from the presence of other salts, other lime salts. For
example, certain parts of the earth contain a great deal of gypsum, or
natural sulphate of lime, and this is soluble to some extent in water.
Water thus hardened is not affected by boiling, or the addition of lime,
and is therefore termed permanently hard water, the water hardened with
dissolved chalk being termed temporarily hard water. I have said nothing
of solid or undissolved impurities in water, which are said to be in
suspension, for the separation of these is a merely mechanical matter of
settling, or filtration and settling combined. As a general rule, the
water of rivers contains the most suspended and vegetable matter and the
least amount of dissolved constituents, whereas spring and well waters
contain the most dissolved matters and the least suspended. Serious
damage may be done to the dyer by either of these classes of impurities,
and I may tell you that the dissolved calcareous and magnesian
impurities are the most frequent in occurrence and the most injurious. I
told you that on boiling, the excess of carbonic acid holding chalk or
carbonate of lime in solution as bicarbonate, is decomposed and
carbonate of lime precipitated. You can at once imagine, then, what
takes place in your steam boilers when such water is used, and how
incrustations are formed. Let us now inquire as to the precise nature of
the waste and injury caused by hard and impure waters. Let us also take,
as an example, those most commonly occurring injurious constituents, the
magnesian and calcareous impurities. Hard water only produces a lather
with soap when that soap has effected the softening of the water, and
not till then. In that process the soap is entirely wasted, and the
fatty acids in it form, with the lime and magnesia, insoluble compounds
called lime and magnesia soaps, which are sticky, greasy, adhesive
bodies, that precipitate and fix some colouring matters like a mordant.
We have in such cases, then, a kind of double mischief--(i) waste of
soap, (ii) injury to colours and dyes on the fabrics. But this is not
all, for colours are precipitated as lakes, and mordants also are
precipitated, and thus wasted, in much the same sense as the soaps are.
Now by taking a soap solution, formed by dissolving a known weight of
soap in a known volume of water, and adding this gradually to hard water
until a permanent lather is just produced, we can directly determine the
consumption of soap by such a water, and ascertain the hardness. Such a
method is called Clark's process of determination or testing, or Clark's
soap test. We hear a great deal just now of soaps that will wash well in
hard water, and do wonders under any conditions; but mark this fact,
none of them will begin to perform effective duty until such hard water
has been rendered soft at the expense of the soap. Soaps made of some
oils, such as cocoa-nut oil, for example, are more soluble in water than
when made of tallow, etc., and so they more quickly soften a hard water
and yield lather, but they are wasted, as far as consumption is
concerned, to just the same extent as any other soaps. They do not,
however, waste so much time and trouble in effecting the end in view,
and, as you know, "Time is money" in these days of work and competition.
After making a soap test as described above, and knowing the quantity of
water used, it is, of course, easy to calculate the annual loss of soap
caused by the hardness of the water. The monthly consumption of soap in
London is 1,000,000 kilograms (about 1000 tons), and it is estimated
that the hardness of the Thames water means the use of 230,000 kilograms
(nearly 230 tons) more soap per month than would be necessary if soft
water were used. Of course the soap manufacturers around London would
not state that fact on their advertising placards, but rather dwell on
the victorious onslaught their particular brand will make on the dirt in
articles to be washed, in the teeth of circumstances that would be
hopeless for any other brand of soap! I have referred to the sticky and
adhesive character of the compounds called lime soaps, formed in hard
waters. Now in washing and scouring wool and other fibres, these sticky
lime soaps adhere so pertinaciously that the fibres, be they of wool,
silk, or any other article, remain in part untouched, impermeable to
mordant or colouring matter, and hence irregular development of colour
must be the consequence. Also an unnatural lustre or peculiar bloom may
in parts arise, ruining the appearance of the goods. In some cases the
lime soaps act like mordants, attracting colouring matter unequally, and
producing patchy effects. In the dye-baths in which catechu and tannin
are used, there is a waste of these matters, for insoluble compounds are
formed with the lime, and the catechu and tannin are, to a certain
extent, precipitated and lost. Some colours are best developed in an
acid bath, such as Cochineal Scarlet, but the presence of the
bicarbonate of lime tends to cause neutralisation of the acidity, and so
the dyeing is either retarded or prevented. Such mordants as "red
liquor" and "iron liquor," which are acetates of alumina and iron
respectively, are also wasted, a portion of them being precipitated by
the lime, thus weakening the mordant baths.
_Ferruginous Impurities in Water._--Iron in solution in water is very
objectionable in dyeing operations. When the iron is present as
bicarbonate, it acts on soap solutions like the analogous lime and
magnesia compounds, producing even worse results. In wool scouring,
cotton bleaching, and other processes requiring the use of alkaline
carbonates, ferric oxide is precipitated on the fibre. A yellowish tinge
is communicated to bleached fabrics, and to dye bright and light colours
is rendered almost out of the question. You may always suspect iron to
be present in water flowing from or obtained directly out of old coal
pits, iron mines, or from places abounding in iron and aluminous shales.
Moreover, you sometimes, or rather generally, find that surface water
draining off moorland districts, and passing over ochre beds, contains
iron, and on its way deposits on the beds of the streamlets conveying
it, and on the stones, red or brown oxide of iron. All water of this
kind ought to be avoided in dyeing and similar operations. The iron in
water from old coal pits and shale deposits is usually present as
sulphate due to the oxidation of pyrites, a sulphuret or sulphide of
iron. Water from heaths and moorlands is often acid from certain
vegetable acids termed "peaty acids." This acidity places the water in
the condition of a direct solvent for iron, and that dissolved iron may
cause great injury. If such water cannot be dispensed with, the best way
is to carefully neutralise it with carbonate of soda; the iron is then
precipitated as carbonate of iron, and can be removed.
_Contamination of Water by Factories._--You may have neighbours higher
up the stream than yourselves, and these firms may cast forth as waste
products substances which will cause immense waste and loss. Amongst
these waste products the worst are those coming from chemical works,
paper works, bleach works, etc. If the paper works be those working up
wood pulp, the pollutions of effluent water will be about as noxious as
they well can be. You will have gums and resins from the wood, calcium
chloride from the bleach vats, acids from the "sours"; resin, and
resin-soaps; there may also be alumina salts present. Now alumina, lime,
resin, and resin-soaps, etc., precipitate dyestuffs, and also soap; if
the water is alkaline, some of the mordants used may be precipitated and
wasted, and very considerable damage done.
Permanent hardness in water, due to the presence of gypsum or sulphate
of lime in solution, may be remedied by addition of caustic soda. Of
course, if an alkaline water is objectionable in any process, the alkali
would have to be neutralised by the addition of some acid. For use in
boilers, water might thus be treated, but it would become costly if
large quantities required such treatment. Water rendered impure by
contaminations from dyehouses and some chemical works can be best
purified, and most cheaply, by simple liming, followed by a settling
process. If space is limited and much water is required, instead of the
settling reservoirs, filtering beds of coke, sand, etc., may be used.
The lime used neutralises acids in the contaminated and impure water,
precipitates colouring matters, mordants, soap, albuminous matters, etc.
_Tests of Purity._--I will now describe a few tests that may be of value
to you in deciding as to what substances are contaminating any impure
waters that may be at hand.
_Iron._--If to a water you suspect to be hard from presence of carbonate
of lime or carbonate of iron in solution in carbonic acid, _i.e._ as
bicarbonates, you add some clear lime-water, and a white precipitate is
produced, you have a proof of carbonate of lime--hardness. If the
precipitate is brownish, you may have, also, carbonate of iron. I will
now mention a very delicate test for iron. Such a test would be useful
in confirmation. If a very dilute solution of such iron water be treated
with a drop or two of pure hydrochloric acid, and a drop or so of
permanganate of potash solution or of Condy's fluid, and after that a
few drops of yellow prussiate of potash solution be added, then a blue
colour (Prussian blue), either at once or after standing a few hours,
proves the presence of iron.
_Copper._--Sometimes, as in the neighbourhood of copper mines or of some
copper pyrites deposits, a water may be contaminated with small
quantities of copper. The yellow prussiate once more forms a good test,
but to ensure the absence of free mineral acids, it is first well to add
a little acetate of soda solution. A drop or two of the prussiate
solution then gives a brown colour, even if but traces of copper are
present.
_Magnesia._--Suppose lime and magnesia are present. You may first
evaporate to a small bulk, adding a drop of hydrochloric acid if the
liquid becomes muddy. Then add ammonia and ammonium oxalate, when lime
alone is precipitated as the oxalate of lime. Filter through blotting
paper, and to the clear filtrate add some phosphate of soda solution. A
second precipitation proves the presence of magnesia.
_Sulphates._--A solution of barium chloride and dilute hydrochloric acid
gives a white turbidity.
_Chlorides._--A solution of silver nitrate and nitric acid gives a white
curdy precipitate.
_Test for Lead in Drinking Water._--I will, lastly, give you a test that
will be useful in your own homes to detect minute quantities of lead in
water running through lead pipes. Place a large quantity of the water in
a glass on a piece of white paper, and add a solution of sulphuretted
hydrogen and let stand for some time. A brown colour denotes lead. Of
course copper would also yield a brown coloration, but I am supposing
that the circumstances preclude the presence of copper.
I have already said that rain water is the purest of natural waters; it
is so soft, and free from dissolved mineral matters because it is a
distilled water. In distilling water to purify it, we must be very
careful what material we use for condensing the steam in, since it is a
fact probably not sufficiently well known, that the softer and purer a
water is, the more liable it is to attack lead pipes. Hence a coil of
lead pipe to serve as condensing worm would be inadmissible. Such water
as Manchester water, and Glasgow water from Loch Katrine still more so,
are more liable to attack lead pipes than the hard London waters. To
illustrate this fact, we will distil some water and condense in a leaden
worm, then, on testing the water with our reagent, the sulphuretted
hydrogen water, a brown colour is produced, showing the presence of
lead. On condensing in a block tin worm, however, no tin is dissolved,
so tin is safer and better as the material for such a purpose than lead.
_Filtration._--We hear a great deal about filtration or filters as
universal means of purifying water. Filtration, we must remember, will,
as a rule, only remove solid or suspended impurities in water. For
example, if we take some ivory black or bone black, and mix it with
water and afterwards filter the black liquid through blotting-paper, the
bone black remains on the paper, and clear, pure water comes through.
Filtering is effective here. If we take some indigo solution, however,
and pour it on to the filter, the liquid runs through as blue as it was
when poured upon the filter. Filtering is ineffective here, and is so
generally with liquids containing matters dissolved in them. But I said
"generally," and so the question is suggested--Will filtration of any
kind remove matters in solution? This question I will, in conclusion,
try to answer. Bone charcoal, or bone black, has a wonderful attraction
for many organic matters such as colours, dyes, and coloured impurities
like those in peat water, raw sugar solutions, etc. For example, if we
place on a paper filter some bone black, and filter through it some
indigo solution, after first warming the latter with some more of the
bone black, the liquid comes through clear, all the indigo being
absorbed in some peculiar way, difficult to explain, by the bone black,
and remaining on the filter. This power of charcoal also extends to
gases, and to certain noxious dissolved organic impurities, but it is
never safe to rely too much on such filters, since the charcoal can at
length become charged with impurities, and gradually cease to act. These
filters need cleaning and renewing from time to time.
LECTURE V
ACIDS AND ALKALIS
_Properties of Acids and Alkalis._--The name acids is given to a class
of substances, mostly soluble in water, having an acid or sour taste,
and capable of turning blue litmus solution red. All acids contain one
or more atoms of hydrogen capable of being replaced by metals, and when
such hydrogen atoms are completely replaced by metals, there result
so-called neutral or normal salts, that is, neutral substances having no
action on litmus solution. These salts can also be produced by the union
of acids with equivalent quantities of certain metallic oxides or
hydroxides, called bases, of which those soluble in water are termed
alkalis. Alkalis have a caustic taste, and turn red litmus solution
blue.
In order to explain what is called the law of equivalence, I will remind
you of the experiment of the previous lecture, when a piece of bright
iron, being placed in a solution of copper sulphate, became coated with
metallic copper, an equivalent weight of iron meanwhile suffering
solution as sulphate of iron. According to the same law, a certain
weight of soda would always require a certain specific equivalent weight
of an acid, say hydrochloric acid, to neutralise its alkaline or basic
properties, producing a salt.
The specific gravities of acids and alkalis in solution are made use of
in works, etc., as a means of ascertaining their strengths and
commercial values. Tables have been carefully constructed, such that
for every degree of specific gravity a corresponding percentage strength
of acidity and alkalinity may be looked up. The best tables for this
purpose are given in Lunge and Hurter's _Alkali-Makers' Pocket-Book_,
but for ordinary purposes of calculation in the works or factory, a
convenient relationship exists in the case of hydrochloric acid between
specific gravity and percentage of real acid, such that specific gravity
as indicated by Twaddell's hydrometer directly represents percentage of
real acid in any sample of hydrochloric acid.
The point at which neutralisation of an acid by alkali or _vice versâ_
just takes place is ascertained very accurately by the use of certain
sensitive colours. At first litmus and cochineal tinctures were used,
but in testing crude alkalis containing alumina and iron, it was found
that lakes were formed with these colours, and they become precipitated
in the solution, and so no longer sensitive. The chemist was then
obliged to resort to certain sensitive coal-tar colours, which did not,
as the dyer and printer knew, form lakes with alumina and iron, such as
methyl orange, fluorescein, Congo red, phenolphthalein, and so forth.
For determining the alkalimetric strength of commercial sodas, a known
weight of the sample is dissolved in water, and a few drops of a
solution of methyl orange are added, which colour the solution yellow or
orange. Into this solution is then run, from a burette or graduated
tube, a standard solution of an acid, that is, a solution prepared by
dissolving a known weight of an acid, say hydrochloric acid, in a known
volume of water. The acid is run in gradually until the yellow colour
changes to pink, at which point the volume of acid used is noted.
Knowing the weight of acid contained in this volume of standard acid,
and having regard to the law of equivalence mentioned above, it is an
easy matter to calculate the amount of alkali equivalent to the acid
used, and from this the alkali contained in the sample.
_Sulphuric Acid._--The first process for manufacturing sulphuric acid or
vitriol was by placing some burning sulphur in a closed vessel
containing some water. The water absorbed the acid formed by the burning
sulphur. It was next discovered that by mixing with the sulphur some
nitre, much more sulphuric acid could be produced per given quantity of
brimstone. At first large glass carboys were used, but in 1746 the
carboys were replaced by chambers of lead containing water at the
bottom, and in these lead chambers the mixture of sulphur and nitre was
burnt on iron trays. Next, although gradually, the plant was divided
into two portions--a furnace for burning the sulphur, and a chamber for
receiving the vapours. The system was thus developed into the one
followed at the present time. The sulphur, or, in most cases, cupreous
iron pyrites (a combination of iron and copper with sulphur), is burned
in specially constructed kilns or furnaces, and the hot gases,
consisting essentially of sulphur dioxide with the excess of air, pass
through flues in which are placed cast-iron "nitre pots" containing a
mixture of nitre (sodium nitrate) and vitriol. The gases thus become
mixed with nitrous fumes or gaseous oxides of nitrogen, and, after
cooling, are ready for mixing with steam or water spray in the lead
chambers in which the vitriol is produced. These oxides of nitrogen
enable the formation of sulphuric acid to take place more quickly by
playing the part of oxygen-carriers. Sulphuric acid is formed by the
union of oxygen with sulphur dioxide and water; the oxides of nitrogen
combine with the oxygen of the air present in the chambers, then give up
this oxygen to the sulphur dioxide and water or steam to form sulphuric
acid, again combine with more oxygen, and so on. The exact processes or
reactions are of course much more complicated, but the above represents
what is practically the ultimate result. It is evident that the gases
leaving the last lead chamber in which the formation of vitriol is
effected, must still contain nitrous fumes, and it becomes a matter of
importance to recover them, so that they can be used over again. To
effect this object, use is made of the solubility of nitrous fumes in
strong vitriol. The gases from the last lead chamber of the series are
passed through what is called a Gay-Lussac tower (the process was
invented by the eminent French chemist Gay-Lussac), which is a tower
made of lead, supported by a wooden framework, and filled with coke or
special stoneware packing, over which strong vitriol is caused to flow.
The vitriol dissolves the nitrogen oxides, and so-called "nitrous
vitriol" flows out at the base of the tower. The recovery of the
nitrogen compounds from the nitrous vitriol is effected in Glover towers
(the invention of John Glover of Newcastle), which also serve to
concentrate to some extent the weak acid produced in the lead chambers,
and to cool the hot gases from the sulphur burners or pyrites kilns. The
weak chamber acid is mixed with the nitrous vitriol from the Gay-Lussac
tower, and the mixture is pumped to the top of the Glover tower, which
is of similar construction to the Gay-Lussac tower, but is generally
packed with flints. This Glover tower is placed between the sulphur
burners or pyrites kilns and the first lead chamber. The nitrous vitriol
passing down the tower meets the hot gases from the kilns, and a
threefold object is effected: (1) The nitrous fumes are expelled from
the nitrous vitriol, and are carried into the chambers, to again play
the part of oxygen-carriers; (2) the weak chamber acid which was mixed
with the nitrous vitriol is concentrated by the hot kiln gases; and (3)
the hot gases themselves are cooled. The acid from the Glover tower is
purified by special treatment--for example, the arsenic may be removed,
after precipitation with sulphuretted hydrogen, in the form of insoluble
arsenic sulphide,--and the purified acid is concentrated by heating in
glass or platinum vessels.
A considerable amount of sulphuric acid is now made by the so-called
"contact process," in which sulphur dioxide and oxygen unite to form
sulphuric acid in presence of a heated "contact" substance, usually some
form of finely-divided platinum.
_Nitric Acid._--This acid is usually prepared by distilling a mixture of
sodium nitrate and vitriol in cast-iron retorts or pots, the nitric acid
being collected in stoneware vessels connected one with another, or, as
is more generally the case at the present time, in condensing apparatus
consisting of stoneware pipes or coils cooled by water. The effluent
gases are passed through a scrubber in order to free them from the last
traces of acid before discharging them into the atmosphere.
_Hydrochloric Acid._--The greater part of the hydrochloric acid
manufactured in Great Britain is obtained as an intermediate product in
the Leblanc alkali process, which will presently be described, being
produced by heating common salt with vitriol. A large quantity is,
however, also produced by the so-called direct process of Hargreaves &
Robinson, which is, in principle, the same method as that employed in
the Leblanc process, except that the intermediate product, vitriol, is
not separated. It consists essentially in passing the hot gases from
pyrites kilns, as used in the manufacture of vitriol, through large
cast-iron vessels containing common salt heated to a high temperature.
Various physical conditions must be complied with in order to make the
process a success. For example, the salt is used in the form of moulded
hard porous cakes made from a damp mixture of common salt and rock salt.
The cast-iron vessels must be heated uniformly, and the hot pyrites kiln
gases must be passed downwards through the salt in order to ensure
uniform distribution. The hydrochloric acid is condensed in stoneware
pipes connected with towers packed with coke or stoneware.
_Alkali: Leblanc Process._--The manufacture of vitriol, as I have
described it to you, is the first step in the Leblanc process. The next
stage consists in the manufacture of sodium sulphate (salt-cake) and
hydrochloric acid from the sulphuric acid and common salt; this is
called the salt-cake process. The production of salt-cake or crude
sodium sulphate is carried out in two stages. A large covered iron pan,
called the decomposing pan or salt-cake pot, is mounted in one part of
the salt-cake furnace, and alongside it is the hearth or bed on which
the second stage of the process, the drying or roasting, is effected.
The mixture of common salt and vitriol is charged into the salt-cake
pot, which is heated by a fire below. When from two-thirds to
three-quarters of the hydrochloric acid has been expelled from the
charge, the mass acquires the consistence of thick dough, and at this
stage it is raked out of the pan on to the roasting hearth alongside,
where the decomposition is completed by means of flames playing directly
on to the top of the charge. The hydrochloric acid evolved during the
process is condensed in much the same manner as in the process of
Hargreaves & Robinson previously described. It is a curious fact that in
the earlier years of the Leblanc process, hydrochloric acid, or "spirits
of salt," as it is frequently called, was a by-product that required all
the vigilance of the alkali-works inspectors to prevent it being allowed
to escape from the chimneys in more than a certain small regulated
amount. Now, it is the principal product; indeed, the Leblanc alkali
maker may be said to subsist on that hydrochloric acid, as his chief
instrument for producing chloride of lime or bleaching powder.
Mechanical furnaces are now used to a large extent for the salt-cake
process. They consist broadly of a large revolving furnace-hearth or
bed, on to which the mixture of salt and vitriol is charged, and on
which it is continuously agitated, and gradually moved to the place of
discharge, by rakes or the like, operated by suitable machinery.