The progress of the science of chemistry forms one phase of the
progress of human thought. While at first mankind was contented to
observe certain phenomena, and to utilize them for industrial purposes,
if they were found suitable, “philosophers,” as the thinking portion
of our race loved to call themselves, have always attempted to assign
some explanation for observed facts, and to group them into similars
and dissimilars. It was for long imagined, following the doctrines of
the Greeks and of their predecessors, that all matter consisted of four
elements or principles, names which survive to this day in popular
language. These were “fire,” “air,” “water,” and “earth.” It was not
until the seventeenth century that Boyle in his _Sceptical Chymist_
(1661) laid the foundations of the modern science, by pointing out
that it was impossible to explain the existence of the fairly numerous
chemical substances known in his day, or the changes which they can be
made to undergo, by means of the ancient Greek hypotheses regarding the
constitution of matter. He laid down the definition of the modern
meaning of the word “element”; he declined to accept the current view
that the properties of matter could be modified by its assimilating
the qualities of fire, air, earth, or water, and he defined an element
as the _constituent_ of a compound body. The first problem, then, to be
solved, was to determine which of the numerous forms of matter were
to be regarded as elementary, and which are compound, or composed of
two or more elements in a state of combination; and to produce such
compounds by causing the appropriate elements to unite with each other.
One of the first objects to excite curiosity and interest was the air
which surrounds us, and in which we live and move and have our being.
It was, however, endowed with a semi-spiritual and scarcely corporeal
nature in the ideas of our ancestors, for it does not affect the
senses of sight, smell, or taste, and though it can be felt, yet it
eludes our grasp. The word “gas,” moreover, was not invented until Van
Helmont devised it to designate various kinds of “airs” which he had
observed. The important part which gases play in the constitution of
many chemical compounds was accordingly overlooked; and, indeed, it
appeared to be almost as striking a feat of necromancy to produce a
quantity of a gas of great volume from a small pinch of solid powder
as for a “Jinn” of enormous stature but of delicate texture to issue
from a brass pot, as related in the _Arabian Nights Entertainments_.
Gradually, however, it came to be recognized, not merely that gases
have corporeal existence, but that they even possess weight. This,
though foreshadowed by Torricelli, Jean Rey, and others, was first
clearly proved by Black, professor of chemistry in Edinburgh, in 1752,
through his masterly researches, as carbonic acid.
The ignorance of the material nature of gases and of their weight
lies at the bottom of the “Phlogistic Theory,” a theory devised by
Stahl about the year 1690, to account for the phenomena of combustion
and respiration and the recovery or “reduction” of metals from their
“earths” by heating with charcoal or allied bodies. According to
this inverted theory, a substance capable of burning was imagined
to contain more or less phlogiston, a principle which it parted
with on burning, leaving an earth deprived of phlogiston, or
“dephlogisticated,” behind if a metal. This earth, when heated with
substances rich in phlogiston, such as coal, wood, flour, and similar
bodies, recovered the phlogiston, which it had lost on burning, and,
with the added phlogiston, its metallic character. Other substances,
such as phosphorus and sulphur, gave solids or acid liquids, to
which phlogiston was not so easy to add; but even they could be
rephlogisticated. On this hypothesis, it was the earths, and such acid
liquids as sulphuric or phosphoric acids, which were the elements; the
metals and sulphur and phosphorus were their compounds with phlogiston.
The discovery of oxygen by Priestley and by Scheele in 1774, and
the explanation of its functions by Lavoisier during the following
ten years, gave their true meaning to these phenomena. It was then
recognized that combustion was union with oxygen; that an “earth” or
“calx” was to be regarded as the compound of a metal with oxygen; that
when a metal becomes tarnished, and converted into such an earthy
powder, it is being oxidized; that this oxide, on ignition with
charcoal or carbon, or with compounds such as coal, flour, or wood,
of which carbon is a constituent, gives up its oxygen to the carbon,
forming an oxide of carbon, carbonic oxide on the one hand, or carbonic
“acid” on the other, while the metal is reproduced in its “reguline”
or metallic condition, and that the true elements are metals, carbon,
sulphur, phosphorus, and similar bodies, and not the products of their
oxidation.
The discovery that air is in the main a mixture of nitrogen, an inert
gas, and oxygen, an active one, together with a small proportion of
carbonic “acid” (or, as it is now termed, anhydride)—a discovery
perfected by Rutherford, Black, and Cavendish—and that water is a
compound with oxygen of hydrogen, previously known as inflammable air,
by Cavendish and by Watt, finally overthrew the theory of phlogiston;
but at the beginning of this century it still lingered on, and was
defended by Priestley until his death in 1804. Such, in brief, was
the condition of chemical thought in the year 1800. Scheele had died
in 1786, at the early age of forty-four; Lavoisier was one of the
victims of the French Revolution, having been guillotined in 1794;
Cavendish had ceased to work at chemical problems, and was devoting his
extraordinary abilities to physical problems of the highest importance,
while living the life of an eccentric recluse, and Priestley, driven
by religious persecution from England to the more tolerant shores of
America, was enjoying a peaceful old age, enlivened by occasional
incursions into the region of sectarian controversy.
The first striking discovery of our century was that of the compound
nature of the alkalies and of the alkaline earths. This discovery was
made by Humphry Davy. Born in Cornwall in 1778, he began the study of
chemistry, self-taught, in 1796; and in 1799 he became director of
the “Pneumatic Institution,” an undertaking founded by Dr. Beddoes,
at Bristol, for the purpose of experiments on the curative effects
of gases in general. Here he at once made his mark by the discovery
of the remarkable properties of “laughing gas,” or nitrous oxide. At
the same time he constructed a galvanic battery, and began to perform
experiments with it in attempting to decompose chemical compounds by
its means. In 1801 Davy was appointed professor of chemistry at the
Royal Institution, a society or club which had been founded a few
years previously by Benjamin Thompson, Count Rumford, for the purpose
of instructing and amusing its members with recent discoveries in
chemistry and natural philosophy. In 1807 Davy applied his galvanic
battery to the decomposition of damp caustic potash and soda, using
platinum poles. He was rewarded by seeing globules of metal, resembling
mercury in appearance, at the negative pole; and he subsequently proved
that these globules, when burned, reproduced the alkali from which
they had been derived. They also combined with “oxymuriatic acid,” as
chlorine (discovered by Scheele) was then termed, forming ordinary
salt, if sodium be employed, and the analogous salt, “muriate of
potash,” if the allied metal, potassium, were subjected to combustion.
By using mercury as the negative pole, and passing a current through a
strong solution of the chloride of calcium, strontium, or barium, Davy
succeeded in procuring mixtures with mercury or “amalgams” of their
metals, to which he gave the names calcium, strontium, and barium.
Distillation removed most of the mercury, and the metal was left
behind in a state of comparative purity. The alkali metals, potassium
and sodium, were found to attack glass, liberating “the basis of the
silex,” to which the name silicon has since been given.
Thus nearly the last of the “earths” had been decomposed. It was proved
that not merely were the “calces” of iron, copper, lead, and other
well-known metals compounds of the respective metals with oxygen, but
Davy showed that lime, and its allies, strontia and baryta, and even
silica or flint, were to be regarded as oxides of elements of metallic
appearance. To complete our review of this part of the subject,
suffice it to say that aluminum, a metal now produced on an industrial
scale, was prepared for the first time in 1827 by Wöhler, professor of
chemistry at Göttingen, by the action of potassium on its chloride,
and alumina, the earthy basis of clay, was shown to be the oxide of
the metal aluminum. Indeed, the preparation of this metal in quantity
is now carried out at Schoffhausen-on-the-Rhine and at the Falls of
Foyers, in Scotland, by electrolysis of the oxide dissolved in melted
cryolite, a mineral consisting of the fluorides of sodium and aluminum,
by a method differing only in scale from that by means of which Davy
isolated sodium and potassium in 1806.
To Davy, too, belongs the merit of having dethroned oxygen from its
central position among the elements. Lavoisier gave to this important
gas the name “oxygen,” because he imagined it to be the constituent
of all acids. He renamed the common compounds of oxygen in such a
manner that the term oxygen was not even represented in the name—only
inferred. Thus a “nitrate” is a compound of an oxide of nitrogen
and an oxide of a metal; a “sulphate,” of the oxide of a metal with
one of the oxides of sulphur, and so on. Davy, by discovering the
elementary nature of chlorine, showed, first, that it is not an
oxide of hydrochloric acid (or muriatic acid as it was then called);
and, second, that the latter acid is the compound of the element
chlorine with hydrogen. This he did by passing chlorine over white-hot
carbon—a substance eminently suited to deprive oxy-compounds of their
oxygen—and proving that no oxide of carbon is thereby produced; by
acting on certain chlorides, such as those of tin or phosphorus with
ammonia, and showing that no oxide of tin or phosphorus is formed; and,
lastly, by decomposing “muriatic acid gas” (gaseous hydrogen chloride)
with sodium, and showing that the only product besides common salt is
hydrogen. Instead, therefore, of the former theory that a chloride was
a compound of the unknown basis of oxymuriatic acid with oxygen and
the oxide of a metal, he introduced the simpler and correct view that
a chloride is merely a compound of the element chlorine with a metal.
In 1813 he established the similar nature of fluorine, pointing out
that on the analogy of the chlorides it was a fair deduction that the
fluorides are compounds of an undiscovered element, fluorine, with
metals; and that hydrofluoric acid is the true analogue of hydrochloric
acid. The truth of this forecast has been established of recent years
by Henri Moissan, who isolated gaseous fluorine by subjecting a mixture
of hydrofluoric acid and hydrogen potassium fluoride contained in a
platinum U tube to the action of a powerful electric current. He has
recently found that the tube may be equally well constructed of copper;
and this may soon lead to the industrial application of the process.
The difficulty of isolating fluorine is due to its extraordinary
chemical energy; for there are few substances, elementary or compound,
which resist the action of this pale yellow, suffocating gas. In 1811
iodine, separated by Courtois from the ashes of sea-plants, was shown
by Davy to be an element analogous to chlorine. Gay-Lussac subsequently
investigated it and prepared many of its compounds; and in 1826 the
last of these elements, bromine, was discovered in the mother-liquor
of sea-salt by Balard. The elements of this group have been termed
“halogens,” or “salt producers.”
While Davy was pouring his researches into the astonished ears of
the scientific and dilettante world, John Dalton, a Manchester
school-master, conceived a theory that has proved of the utmost service
to the science of chemistry, and which bids fair to outlast our day. It
had been noticed by Wenzel, by Richter, by Wollaston, and by Cavendish,
towards the end of the last century, that the same compounds contain
the same constituents in the same proportions, or, as the phrase runs,
“possess constant composition.” Wollaston, indeed, had gone one step
farther, and had shown that when the vegetable acid, oxalic acid, is
combined with potash, it forms two compounds, in one of which the acid
is contained in twice as great an amount relatively to the potash as
in the other. The names _monoxalate_ and _binoxalate_ of potash were
applied to these compounds, to indicate the respective proportions of
the ingredients. Dalton conceived the happy idea that by applying the
ancient Greek conception of atoms to such facts the relative weights
of the atoms could be determined. Illustrating his views with the two
compounds of carbon with hydrogen, marsh gas and olefiant gas, and with
the two acids of carbon, carbonic oxide, carbonic “acid,” he regarded
the former as a compound of one atom of carbon and one of hydrogen, and
the second as a compound of one atom of carbon and two of hydrogen, and
similarly for the two oxides of carbon. Knowing the relative weights
in which these elements enter into combination, we can deduce the
relative weights of the atoms. Placing the relative weight of an atom
of hydrogen equal to unity, we have:
Marsh Olefiant | Carbonic Carbonic
Gas Gas | Oxide Acid
Carbon 6 6 | Carbon 5 6
Hydrogen 1 2 | Oxygen 8 16
Thus the first compound, marsh gas, was regarded by Dalton as
composed of an atom of carbon in union with an atom of hydrogen; or,
to reproduce his symbols, as ⊜◉; while the second, olefiant gas, on
this hypothesis, was a compound of two atoms of hydrogen with one of
carbon, or ◉◍◉. Similarly the symbols ◍○, and ○◍○ were given to the two
compounds of carbon with oxygen. So water was assigned the symbol ◉○,
for Dalton imagined it to be a compound of one atom of hydrogen with
one of oxygen. Compounds containing only two atoms were termed by him
“binary”; those containing three, “ternary”; four, “quaternary,” and
so on. The weight of an atom of oxygen was eight times that of an atom
of hydrogen; while that of an atom of carbon was six times as great
as the unit. By assigning symbols to the elements, consisting of the
initial letters of their names, or of the first two letters, formulas
were developed, indicating the composition of the compound, the atomic
weights of the elements being assured. Thus, NaO signified a compound
of an atom of sodium (natrium), weighing twenty-three times as much as
a similar atom of hydrogen, with an atom of oxygen, possessing eight
times the weight of an atom of hydrogen. Therefore, thirty-one pounds
of soda should consist of twenty-three pounds of sodium in combination
with eight pounds of oxygen, for, according to Dalton, each smallest
particle of soda contains an atom of each element, and the proportion
is not changed, however many particles be considered.
It has been pointed out by Judge Stallo, of Philadelphia, in his
_Concepts of Physics_, that such a hypothesis as that of Dalton is
no explanation; that a fact of nature, as, for example, the fact of
simple and multiple proportions, is not _explained_ by being minified.
Allowing the general truth of this statement, it is, nevertheless,
undoubted that chemistry owes much to Dalton’s hypothesis—a lucky
guess at first, it represents one of the fundamental truths of nature,
although its form must be somewhat modified from that in which Dalton
conceived it. Dalton’s work was first expounded by Thomas Thomson,
professor at Glasgow, in his _System of Chemistry_, published in 1805;
and subsequently in Dalton’s own _New System of Chemical Philosophy_,
the three volumes of which were published in 1808, in 1810, and in 1827.
The determination of these “Constants of Nature” was at once followed
out by many chemists, Thomson among the first. But chief among the
chemists who have pursued this branch of work was Jacob Berzelius,
a Swede, who devoted his long life (1779–1848) to the manufacture
of compounds, and to the determination of their composition, or, as
it is still termed, the determination of the “atomic weights”—more
correctly, “equivalents”—of the elements of which they are composed.
It is to him that we owe most of our analytical methods, for, prior to
his time, there were few, if any, accurate analyses. Although Lavoisier
had devised a method for the analysis of compounds of carbon, viz.,
by burning the organic compounds in an atmosphere of oxygen contained
in a bell-jar over mercury, and measuring the volume of carbon dioxide
produced, as well as that of the residual oxygen, Berzelius achieved
the same results more accurately and more expeditiously by heating the
substance, mixed with chlorate of potassium and sodium chloride, and
then estimating the hydrogen as well as the carbon; this process was
afterwards perfected by Liebig. Berzelius, however, was able to show
that compounds of carbon, like those of other elements, were instances
of combination in constant and in multiple proportions.
In 1815 two papers were published in the _Annals of Philosophy_ by Dr.
Prout, which have had much influence on the progress of chemistry. They
dealt with the figures which were being obtained by Thomson, Berzelius,
and others, at that time supposed to represent the “atomic weights” of
the elements. Prout’s hypothesis, based on only a few numbers, was that
the atomic weights of all elements were multiples of that of hydrogen,
taken as unity. There was much dispute regarding this assertion at the
time, but as it was contradicted by Berzelius’s numbers, the balance
of opinion was against it. But about the year 1840 Dumas discovered an
error in the number (12.12) given by Berzelius as the atomic weight of
carbon; and with his collaborator, Stas, undertook the redetermination
of the atomic weights of the commoner elements—for example, carbon,
oxygen, chlorine, and calcium. This line of research was subsequently
pursued alone by Stas, whose name will always be remembered for the
precision and accuracy of his experiments. At first Dumas and Stas
inclined to the view that Prout’s hypothesis was a just one, but it
was completely disproved by Stas’s subsequent work, as well as by that
of numerous other observers. It is, nevertheless, curious that a much
larger proportion of the atomic weights approximate to whole numbers
than would be foretold by the doctrine of chances, and perhaps the last
has not been heard of Prout’s hypothesis, although in its original
crude form it is no longer worthy of credence.
One of the most noteworthy of the discoveries of the century was
made by Gay-Lussac (1778–1850) in the year 1808. In conjunction with
Alexander von Humboldt, Gay-Lussac had rediscovered about three years
before what had previously been established by Cavendish—namely,
that, as nearly as possible, two volumes of hydrogen combine with one
volume of oxygen to form water, the gases having been measured at the
same temperature and pressure. Humboldt suggested to Gay-Lussac that
it would be well to investigate whether similar simple relations exist
between the volumes of other gaseous substances when they combine with
each other. This turned out to be the case; it appeared that almost
exactly two volumes of carbonic oxide unite with one volume of oxygen
to form carbon dioxide; that equal volumes of chlorine and hydrogen
unite to form hydrochloric acid gas; that two volumes of ammonia
gas consist of three volumes of hydrogen in union with one volume
of nitrogen, and so on. From such facts, Gay-Lussac was led to make
the statement that: The weights of equal volumes of both simple and
compound gases, and therefore their densities, are proportional to
their empirically found combining weights, or to rational multiples
of the latter. Gay-Lussac recognized this discovery of his to be a
support for the atomic theory; but it did not accord with many of the
then received atomic weights. The assumption that equal volumes of
gases contain equal numbers of particles, or, as they were termed by
him, _molécules intégrantes_, was made in 1811 by Avogadro, professor
of physics at Turin (1776–1856). This theory, which has proved of the
utmost importance to the sciences both of physics and of chemistry,
had no doubt occurred to Gay-Lussac, and had been rejected by him for
the following reasons: A certain volume of hydrogen, say one cubic
inch, may be supposed to contain an equal number of particles (atoms)
as an equal volume of chlorine. Now these two gases unite in equal
volumes. The deduction appears so far quite legitimate that one atom of
hydrogen has combined with one atom of chlorine. But the resulting gas
occupies two cubic inches, and must therefore contain the same number
of particles of hydrogen chloride, the compound of the two elements,
as one cubic inch originally contained of hydrogen, or of chlorine.
Thus we have two cubic inches containing, of uncombined gases, twice
as many particles as is contained in that volume, after combination.
Avogadro’s hypothesis solved the difficulty. By premising two different
orders of particles, now termed _atoms_ and _molecules_, the solution
was plain. According to him, each particle, or molecule, of hydrogen is
a complex, and contains two atoms; the same is the case with chlorine.
When these gases combine, or rather _react_, to form hydrogen chloride,
the phenomenon is one of a change of partners; the molecule, the double
atom, of hydrogen splits; the same is the case with the molecule of
chlorine; and each liberated atom of hydrogen unites with a liberated
atom of chlorine, forming a compound, hydrogen chloride, which equally
consists of a molecule, or double atom. Thus two cubic inches of
hydrogen chloride consist of a definite number of molecules, equal
in number to those contained in a cubic inch of hydrogen, plus those
contained in a cubic inch of chlorine. The case is precisely similar,
if other compounds of gases be considered.
Berzelius was at first inclined to adopt this theory, and indeed went
so far as to change many of his atomic weights to make them fit it.
But later he somewhat withdrew from his position, for it appeared to
him that it was hazardous to extend to liquids and solids a theory
which could be held only of gases. Avogadro’s suggestion, therefore,
rested in abeyance until the publication, in 1858, by Cannizzaro, now
professor of chemistry in Rome, of an essay in which all the arguments
in favor of the hypothesis were collected and stated in a masterly
manner. It will be advisable to revert to this hypothesis at a later
point, and to consider other guides for the determination of atomic
weights.
In 1819, Dulong (1785–1838), director of the Ecole Polytechnique at
Paris, and Petit (1791–1820), professor of physics there, made the
discovery that equal amounts of heat are required to raise equally the
temperature of solid and liquid elements, provided quantities are taken
proportional to their atomic weights. Thus, to raise the temperature of
56 grammes of iron through one degree requires approximately the same
amount of heat as is required to raise through one degree 32 grammes of
sulphur, 63.5 grammes of copper, and so on; these numbers representing
the atomic weights of the elements named. In other words, _equal
numbers of atoms have equal capacity for heat_. The number of heat
units, or calories (one calory is the amount of heat required to raise
the temperature of 1 gramme of water through 1° C.), which is necessary
to raise the atomic weight expressed in grammes of any solid or liquid
element through 1° C. is approximately 6.2; it varies between 5.7 and
6.6 in actual part. This affords a means of determining the true value
of the atomic weight of an element, as the following example will show:
The analysis of the only compound of zinc and chlorine shows that it
contains 47.49 per cent. of zinc and 52.16 per cent. of chlorine. Now
one grain of hydrogen combines with 35.5 grains of chlorine to form
36.5 grains of hydrogen chloride; and, as already remarked, one volume
of hydrogen and one volume of chlorine combine, forming two volumes of
hydrogen chloride. Applying Avogadro’s hypothesis, one molecule of
hydrogen and one molecule of chlorine react to yield two molecules of
hydrogen chloride; and as each molecule is supposed to consist in this
case of two atoms, hydrogen chloride consists of one atom of each of
its constituent elements. The amount of that element, therefore, which
combines with 35.5 grains of chlorine may give the numerical value of
the atomic weight of the element, if the compound contains one atom
of each element; in that case the formula of the above compound would
be zinc, and the atomic weight of zinc, 32.7; but if the formula is
ZuCl3, the atomic weight of zinc would be 32.7 × 2; if ZuCl3, 32.7 × 3,
and so on. The specific heat of metallic zinc enables this question to
be solved. For it has been found, experimentally, to be about 0.095;
and 6.2 ÷ 0.095 = 65.2, a close approximation to 32.7 × 2 = 65.4. The
conclusion is therefore drawn that zinc chloride is composed of one
atom of zinc in combination with two atoms of chlorine, that the atomic
weight of zinc is 65.4, and that the molecular weight of zinc chloride
is 65.4 + (35.5 × 2) = 136.4. Inasmuch as the relative weight of a
molecule of hydrogen is 2 (that of an atom being 1), zinc chloride in
the gaseous state should be 136.4 ÷ 2 = 68.2 times that of hydrogen,
measured at the same temperature and pressure. This has been found,
experimentally, to be the case.
The methods of determining the vapor densities, or relative weights of
vapors, are three in number; the first method, due to Dumas (1827),
consists in vaporizing the substance in question in a bulb of glass or
of porcelain, at a known temperature, closing the bulb while still hot,
and weighing it after it is cold. Knowing the capacity of the bulb,
the weight of hydrogen necessary to fill it at the desired temperature
can be calculated, and the density of the vapor thus arrived at.
A second method was devised by Gay-Lussac and perfected by A. W.
Hofmann (1868); and a third, preferable for its simplicity and ease of
execution, is due to Victor Meyer (1881).
In 1858, as already remarked, Cannizzaro showed the connection between
these known facts, and for the first time attention was called to
the true atomic weights, which were, up to that time, confused with
equivalents, or weights of elements required to replace one unit weight
of hydrogen. These were generally regarded as atomic weights by Dalton
and his contemporaries.
Some exceptions had been observed to the law of Dulong and Petit, viz.,
beryllium, or glucinium, an element occurring in emeralds; boron,
of which borax is a compound; silicon, the component of quartz and
flint, and carbon. It was found by Weber that at high temperatures
the specific heats of these elements are higher, and the atomic heats
approximate to the number of 6.2; but this behavior is not peculiar to
these elements, for it appears that the specific heat of all elements
increases with rise of temperature.
A certain number of exceptions have also been noticed to the law of
Gay-Lussac, which may be formulated: the molecular weight of a compound
in a gaseous state is twice its density referred to hydrogen. Thus
equal volumes of ammonia and hydrogen chloride unite to form ammonium
chloride. It was to be expected that the density should be half the
molecular weight, thus:
NH3 + HCl = NH4Cl; and 53.5 ÷ 2 = 26.75 = density.
(14+3) (1+35.5) 53.5
But the density actually found is only half that number, viz., 13.37;
and for long this and similar cases were supposed to be exceptions
to the law of Gay-Lussac, viz., that equal volumes of gases at the
same pressure expand equally for equal rise of temperature. In other
instances the gradual decrease in density with rise of temperature can
be followed, as with chloral hydrate, the products of which are chloral
and water.
It was recognized by St. Claire Deville (1857) that the decrease
in density of such mixtures of gases was due, not to their being
exceptions to Avogadro’s law, but to the gradual decomposition of the
compound body with rise of temperature. To this gradual decomposition
he gave the name _dissociation_. This conception has proved of the
utmost importance to the science, as will be seen in the sequel. To
take the above instance of ammonium chloride, its abnormal density is
due to its dissociation into ammonia and hydrogen chloride; and the gas
which is obtained on raising its temperature consists, not of gaseous
ammonium chloride, but of a mixture of ammonia and hydrogen chloride,
which, as is easily seen, occupy, when separate, twice the volume that
would be occupied by the gaseous compound. Of recent years it has been
shown by Brereton Baker that, if perfectly free from moisture, ammonium
chloride gasifies as such, and that its density in the state of vapor
is, in fact, 26.75.
The molecular complexity of gases has thus gradually become
comprehended, and the truth of Avogadro’s law has gained acceptance.
And as a means of picturing the behavior of gaseous molecules, the
“Kinetic Theory of Gases” has been devised by Joule, Clausius, Maxwell,
Thomson (Lord Kelvin), and others. On the assumption that the pressure
of a gas on the walls of the vessel which contains it is due to the
continued impacts of its molecules, and that the temperature of a gas
is represented by the product of the mass of the molecules, or the
square of their velocity, it has been possible to offer a mechanical
explanation of Boyle’s law, that at constant temperature the volume
of a gas diminishes in proportion as the pressure increases; of
Gay-Lussac’s law, that all gases expand equally for equal rise of
temperature, provided pressure is kept constant; the condition being
that equal volumes of gases contain equal numbers of molecules. A
striking support is lent to this chain of reasoning by the facts
discovered by Thomas Graham (1805–1869), professor at University
College, London, and subsequently master of the Royal Mint. Graham
discovered that the rate of diffusion of gases into each other is
inversely as the square roots of their densities. For instance, the
density of hydrogen being taken as unity, that of oxygen is sixteen
times as great; if a vessel containing hydrogen be made to communicate
with one containing oxygen, the hydrogen will pass into the oxygen and
mix with it; and, conversely, the oxygen will pass into the hydrogen
vessel. This is due to the intrinsic motion of the molecule of each
gas. And Graham found, experimentally, that for each volume of oxygen
which enters the hydrogen vessel four volumes of hydrogen will enter
the oxygen vessel. Now, 4 = √16; and as these masses are relatively
1 and 16, and their temperatures are equal, the square of their
velocities are respectively 1 and 16.
The question of the molecular complexity of gases being thus disposed
of, it remains to be considered what are the relative complexity of
liquid molecules. The answer is indicated by a study of the capillary
phenomena of liquids, one method of measuring which is the height of
their ascent in narrow or capillary tubes. We shall not enter here into
detail as to the method and arguments necessary; suffice it to say that
the Hungarian physicist Eötvös was the first to indicate the direction
of research, and that Ramsay and Shields succeeded in proving that the
complexity of the molecules of most liquids is not greater than that
of the gases which they form on being vaporized; and also that certain
liquids, _e.g._, water, the alcohols, and other liquids, are more or
less “associated,” _i.e._, their molecules occur in couplices of two,
three, four, or more, and as the temperature is raised the complexity
of molecular structure diminishes.
As regards the molecular complexity of solids, nothing definite
is known, and, moreover, there appears to be no method capable of
revealing it.
While the researches of which a short account has now been given have
led to knowledge regarding the nature of molecules, the structure
of the molecule has excited interest since the early years of the
century, and its investigation has led to important results. The
fact of the decomposition of acidified water by an electric current,
discovered by Nicholson and Carlisle, and of salts into “bases” and
“acids” by Berzelius and Hisinger in 1803, led to the belief that
a close connection exists between electric energy, or, as it was
then termed, “electric force,” and the affinity which holds the
constituents of chemical compounds in combination. In 1807 Davy
propounded the theory that all compounds consist of two portions, one
electro-positive and the other electro-negative. This idea was the
result of experiments on the behavior of substances, such, for example,
as copper and sulphur—if portions of these elements be insulated and
then brought into contact they become oppositely electrified. The
degree of electrification is intensified by rise of temperature until,
when combination ensues, the electrification vanishes. Combination,
therefore, according to Davy, is concurrent with the equalization of
potentials. In 1812 Berzelius brought forward an electro-chemical
theory which for the following twenty years was generally accepted.
His primary assumption was that the atoms of elements, or, in certain
cases, groups of atoms, are themselves electrified; that each atom, or
group of atoms, possesses two poles, one positive, the other negative;
that the electrification of one of these poles predominates over
that of the other, so that the atom or group is itself, as a whole,
electro-positive, or electro-negative; that combination ensued between
such oppositely electrified bodies by the neutralization, partial or
complete, of their electric charges; and, lastly, that the polarity
of an element or group could be determined by noting whether the
element or group separated at the positive or at the negative pole of
the galvanic battery, or electrolysis. For Berzelius, oxygen was the
most electro-negative and potassium the most electro-positive of the
elements, the bridge between the “non-metals” and the “metals” being
hydrogen, which, with nitrogen, forms a basic, or electro-positive,
group, while with chlorine, etc., it forms electro-negative groups.
The fact that an electric current splits compounds in solution
into two portions led Berzelius to devise his “dualistic” system,
which involved the assumption that all compounds consist of two
portions, one electro-positive, the other electro-negative.
Thus sulphate of magnesium and potassium was to be regarded as
composed of electro-positive potassium sulphate in combination
with electro-negative magnesium sulphate; the former in its turn
consisted of electro-negative sulphur trioxide (SO3) in combination
with electro-positive oxide of potassium (K2O); while each of these
proximate constituents of potassium sulphate were themselves composed
of the electro-negative oxygen in combination with electro-positive
sulphur, or potassium. On contrasting sulphur with potassium, however,
the former was considered more electro-negative than the latter; so
that the group SO3 as a whole was electro-negative, while K2O was
electro-positive. The symbols given above, which are still in universal
use, were also devised by Berzelius for the purpose of illustrating and
emphasizing his views. These views, however, met with little acceptance
at the time in England.
Lavoisier’s idea, that oxygen was the necessary constituent of all
acids, began about this time to lose ground. For Davy had proved the
elementary nature of chlorine; and hydrochloric acid, one of the
strongest, was thus seen to contain no oxygen, and Davy expressed the
view, founded on his observation, that iodic “acid,” I2O5, was devoid
of acid properties until dissolved in water, and that the essential
constituent of all acids was hydrogen, not oxygen. The bearing of this
theory on the dualistic theory is, that while, _e.g._, sulphuric acid
was regarded by Berzelius as SO3, containing no hydrogen, and was
supposed to be separated as such at the positive pole of a battery,
Davy’s suggestion led to the opposite conclusion that the formula
of sulphuric acid is H2SO4, and that by the current it is resolved
into H2 and SO4. Faraday’s electrolytic law, that when a current is
passed through electrolytes in solution the elements are liberated in
quantities proportional to their equivalents, led to the abandonment
of the dualistic theory. For when a current is passed in succession
through acidified water, fused lead chloride, and a solution of
potassium sulphate, the quantities of hydrogen and oxygen from the
water, of lead and chlorine from the lead chloride, and the potassium
of the sulphate are in accordance with Faraday’s law. But in addition
to the potassium there is liberated at the same pole an equivalent of
hydrogen. Now, if Berzelius’s theory be true, the products should be
SO3 and K2O, but if the opposite view be correct, then K2 is liberated
first and by its subsequent action on water it yields potash and
its equivalent of hydrogen. This was pointed out first by Daniell,
professor at King’s College, London, and it was regarded as a powerful
argument against Berzelius’s system. In 1833, too, Graham investigated
the phosphoric acids, and prepared the salts of three, to which he
gave the names, _ortho-_, _pyro-_, and _meta-_ phosphoric acids. To
understand the bearing of this on the doctrine of dualism it must be
remembered that P2O5, pentoxide of phosphorus, was at that date named
phosphoric acid. When dissolved in water it reacts with bases, forming
salts—the phosphates. But the quantity of water necessary was not then
considered essential; Graham, however, showed that there exist three
series of salts—one set derived from P2O5,3H2O, one from P2O5,2H2O,
and a third from P2O5,H2O. His way of stating the fact was that water
could play the part of a base; for example, the ordinary phosphate
of commerce possessed, according to him, the formula P2O5,2Na2O,H2O,
two-thirds of the “water of constitution” being replaced by oxide of
sodium. Liebig, then professor at Giessen (1803–1873), founded on these
and on similar observations of his own the doctrine of poly-basic
acids—acids in which one, two, three, or more atoms of hydrogen
were replaceable by metals. Thus, instead of writing, as Graham did,
P2O5,2Na2O,H2O, he wrote, PO4Na2H; and for orthophosphoric acid
PO4H3. The group of atoms (PO4), therefore, existed throughout the
whole series of orthophosphates, and could exist in combination with
hydrogen, with hydrogen and metals, or with metals alone. Similarly
the group (P2O7) was characteristic of pyrophosphates and (PO3) of
metaphosphates, for P2O5,2H2O=(P2O7)H4; and P2O5,H2O=2(PO3)H.
The first clear ideas of the structure of the molecule were, however,
gained from the study of the compounds of carbon. It was difficult to
apply the dualistic theory to them. For few of them are electrolytes,
and therefore their products of electrolysis, being non-existent,
could not be classified. Nevertheless, Gay-Lussac regarded alcohol,
C2H6O, as a compound of C2H4, ethylene, and H2O, water; and oxalic
acid (anhydrous), C2O3, as one of CO2 with CO. The discovery of
“isomeric compounds,” _i.e._, of compounds which possess the same
ultimate formula and yet differ entirely in their properties,
forced upon chemists the necessity of attending to the structure of
the molecule; for only by such a supposition could the difference
between two isomeric bodies be explained. In 1823 Liebig discovered
that silver fulminate and silver cyanate both possessed the empirical
formula AgCNO; in 1825 this was followed by the discovery by Faraday
that oil gas contains a hydrocarbon identical in composition with
ethylene, C2H4, yet differing from it in properties; and in 1829
Wöhler, professor in Göttingen (1800–1882), discovered that urea, a
constituent of urine, could be produced by heating ammonium cyanate,
NH4CNO, a substance of the same formula. It therefore became clear that
the identity of a compound must depend on some other cause than its
ultimate composition.
In 1833 Liebig and Wöhler took an important step in elucidating this
question by their investigations on benzoic acid and acid obtainable
by distilling a resin named gum benzoin. They showed that this acid,
C7H6O2, could be conceived as consisting of the group C7H5O, to which
they gave the name “benzoyl,” in combination with OH; that benzoic
aldehyde, C7H6O, might be regarded as its compound with hydrogen; that
it also formed compounds with chlorine, and bromine, and sulphur,
and replaced hydrogen in ammonia (C7H6O,NH2). They termed this
group, benzoyl, a “compound element” or a “radical.” This research
was followed by one by Robert Bunsen, professor at Heidelberg, born
in 1811, and recently (1899) dead, which bore reference to cacodyl,
a compound of arsenic, carbon and hydrogen, in which the idea of a
radical was confirmed and amplified.
The idea of a radical having thus become established, Jean Baptiste
Andrée Dumas, professor in Paris (1800–1884), propounded the theory
of “substitution,” _i.e._, that an element such as chlorine or oxygen
(which, be it noticed, is electro-negative on Berzelius’s scale) could
replace hydrogen in carbon compounds, atom for atom, the resulting
compound belonging to the same “type” as the one from which it was
derived. And Laurent, warden of the mint at Paris (1807–1853), and
Gerhardt, professor at Montpelier and at Strasburg (1816–1856),
emphasized the fact that one element, be it what it may, can replace
another without fundamentally altering its chemical character, and
also that an atom of hydrogen can be replaced by a group of atoms or
radical, behaving for the occasion like the atom of an element. It is
to Laurent and Gerhardt that we owe the definition of an atom—_the
smallest quantity of an element which can be present in a compound_; an
equivalent—_that weight of an element which combines with or replaces
one part by weight of hydrogen_; and a molecule—_the smallest quantity
which can exist in a free state, whether of an element or a compound_.
They recognized, too, that a molecule of hydrogen, chlorine, etc.,
consists of two atoms.
In 1849 Wurtz, professor in Paris (1817–1884), and Hofmann, then
professor in the College of Chemistry in London, afterwards at Berlin
(1818–1892), discovered a series of compounds allied to ammonia, NH3,
in which one or more atoms of hydrogen were replaced by a group or
radical, such as methyl (CH3), ethyl (C2H5), or phenyl (C6H5). Wurtz
referred such compounds to the ammonia “type.” They all resemble
ammonia in their physical properties—smell, taste, etc.—as well as
in their power of uniting with acids to form salts resembling ammonium
chloride (NH4Cl), and other ammonium compounds. Shortly afterwards
Williamson, professor at University College, London, added the “water
type,” in consequence of his researches on “mixed ethers”—bodies in
which the hydrogen of water might be regarded as replaced by organic
radicals. Thus we have the series:
H. O. H.; CH3. O. H.; CH3. O. CH3; and NH3; NH2; H3; NH(CH3)2; and
N(CH3)3; the first representing compounds following the water type,
the latter the ammonia type. This suggestion had been previously made
by Laurent, in 1846. But Williamson extended his views to inorganic
compounds; thus, sulphuric acid was represented as constructed on the
double water type—HO. SO2. OH, being derived from H. O. (H. H) O.
H, the two hydrogen atoms enclosed in brackets being replaced by the
radical SO2. To these types Gerhardt added the hydrogen and hydrogen
chloride types, H.H. and H.Cl; and, later, Kekulé, professor in Bonn
(1829), added the marsh gas type C(H)4. The next important step was
taken by Frankland, professor in the Royal School of Mines, London;
his work, however, had been anticipated by Cunn Brown, professor at
Edinburgh University, in a pamphlet even yet little known. It was to
attribute to elements one or more powers of combination. To these
he gave the name “valency,” and the capacity of possessing valency
was called “quantivalence.” Thus hydrogen was taken as a “monad,” or
monovalent. Chlorine, because it unites with hydrogen atom to atom,
is also a monad. Oxygen, having the power to combine with two atoms
of hydrogen, was termed a dyad, or divalent; nitrogen a triad, or
trivalent; carbon a tetrad, or tetravalent, and so on. This is evident
from inspection of the formulas of their compounds with hydrogen, thus:
H H H
/ \ /
H——Cl; H——O——H; H——N ; C
\ / \
H H H
Instances of penta, hexa, and even hepta-valency are not wanting.
This was the key to unlock the structure of chemical compounds; and
Frankland’s views, just stated, are still held by chemists. The
determination of the constitution of compounds, chiefly those of
carbon, occupied the attention of chemists, almost exclusively, until
1880. The plan of action is much the same as that of a mechanician who
wishes to imitate a complicated mechanism. He must first dissect it
into groups of mechanical contrivances; these are next constructed; and
they are finally built together into the complete machine. In certain
cases the atoms of carbon are arranged in “chains,” as, for example, in
pentyl alcohol:
H3C——C——C——C——C——O——H
H2 H2 H2 H2
each atom being tetrad, and its “affinities,” or powers of combination,
saturated either with hydrogen or with those of neighboring atoms of
carbon; in others they are in the form of a “ring,” as in benzene, the
formula of which was first suggested by Kekulé, viz.:
H H
C——C
/ \
HC CH;
\ /
C==C
H H
or in both, as in ethyl benzene,
H H
C——C
/ \ H H
HC C——C——CH.
\ / H H
C==C
H H
One or more atoms of nitrogen, or of oxygen, may form part of the
circle, as in pyridine:
H H H H
C——C C C
/ \ /
N CH and furfurane, O == ,
\ / \
C==C C C
H H H H
and so on. By means of conceptions such as these many interesting
compounds have been built up out of the elements which they contain;
_e.g._, urea and uric acid, constituents of urine; theobromine and
caffeine, the essential principles of cocoa and tea; alizarine and
indigo, valuable dyestuffs; and several of the alkaloids, bitter
principles contained in plants, of great medicinal value.
They have led, too, to the discovery of many brilliant colors, now
almost universally employed, to the exclusion of those less brilliant,
because less pure, derived from plants, and in one or two cases
from animals; the manufacture of gun-cotton, dynamite, and similar
high explosives; and to the development of the candle industry; the
sugar manufacture; to improvement in tanning, in brewing, and in the
preparation of gas and oils for illuminating purposes. In short, it may
be said that the industrial progress of the latter half of the century
has been due to the theoretical views of which a short sketch has just
been given.
Such formulas, however, can evidently not represent the true
constitution of matter, inasmuch as the atoms are imagined to lie
on a plane, whereas it is evident that they must occupy space
of three dimensions and possess the attributes of solidity. The
conception which led to the formulation of such views was due first
to Pasteur, in his later years director of the institute known by
his name at Paris, and more directly to LeBel and Van’t Hoff, now
professor at Berlin, independently of each other. In 1848 Pasteur
discovered that it was possible to separate the two varieties of
tartaric acid from each other; and that that one which rotated the
plane of polarized light to the right gave crystals with an extra
face, unsymmetrically disposed with regard to the other faces of the
crystal. The variety, the solution of which in water was capable of
producing left-handed rotation, also possessed a similar face, but so
placed that its reflection in a mirror reproduced the right-handed
variety. Pasteur also showed that a mixture of these acids gave
crystals not characterized by an unsymmetrically placed face; and
also that the solution was without action on polarized light. These
observations remained unexplained, until LeBel and Van’t Hoff, in
1874, simultaneously and independently devised a theory which has,
up till now, stood the test of research. It is briefly this: Imagine
two regular tetrahedra, or three-sided pyramids, standing each on its
triangular base. An idea can best be got by a model, easily made by
laying on a table three lucifer matches so as to form an equilateral
triangle, and erecting a tripod with three other matches, so that
each leg of the tripod stands on one corner of the triangle. At the
centre of such a tetrahedron, an atom of carbon is supposed to be
placed. Marsh gas, CH4, is supposed to have such a structure, each
corner, or solid angle of the structure (of which there are four),
being occupied by an atom of hydrogen. This represents the solid or
_stereochemical_ formula of methane or marsh gas. Now, suppose one of
the atoms of hydrogen in each of these structures to be replaced by
chlorine, the group (OH), or any other monovalent element or group. It
is evident that if not exactly similar (owing to the replacement not
having been made at similar corners in each), the two structures can be
made similar by turning one of them round, until the position of the
substituting atom or group (which we will term X) coincides in position
with X in the stationary one. If two such replacements be made, say,
with X and Y in each, coincidence can again be made to take place; but
the same is not the case if X, Y, and Z replace three atoms of hydrogen
in the structure; for there is one way of replacement which is the
optical image of the other, and represents the other’s reflection in a
mirror.
[Illustration: (Tetrahedron XYZ) and (Tetrahedron XZY)]
Now, it is found that when the four corners of such a structure are
occupied by four separate atoms or groups, or when (as the expression
goes) the body contains an “asymmetrical carbon atom,” if the substance
or one of its derivations can be obtained in a crystalline form, the
crystals are also asymmetric, _i.e._, arc develops a face which is the
mirror-reflection of a similar face developed on the other variety;
and if a beam of polarized light be passed through the solution of
the substance, its plane is rotated to the left if one variety be
used, and, if the other, to the right. This hypothesis of LeBel’s
and Van’t Hoff’s has had an enormous influence on the progress of
organic chemistry. By its means Fischer, now professor at Berlin, has
explained the reason of the existence of the enormous number of bodies
analogous to grape and cane sugar, and has prepared many new varieties;
and it appears likely that the terpenes, a class of bodies allied to
turpentine, and comprising most of the substances to which the odor of
flowers is due, may thereby find their explanation. It may be mentioned
in passing that Pasteur, having found that ordinary mould destroyed
one variety of tartaric acid rather than the other in a mixture of
the two, and made use of this observation in order to prepare the
unattached variety in a state of purity, was led to study the action
of organisms more or less resembling mould; and that this has led
to the development of the science of bacteriology, which has had an
enormous influence on our views regarding fermentation in general,
and guides the work of our physicians, our surgeons (witness Lister’s
antiseptic treatment), our sanitary engineers in their estimate of
the purity of drinking-water and of the disposal of sewage, of our
manufacturers of beer and spirits, of wine-growers, and more recently
of farmers. All these processes depend upon the action of organisms
in producing chemical changes, whether in the tissues of the body,
causing or curing disease, or in the production of flavored alcohol
from sugar, or in the manufacture of butter and cheese, or in preparing
the land for the reception of crops. We also owe to the genius of
Van’t Hoff the most important advance of recent times in the region
of physical chemistry. It has been observed by Raoult, professor at
Grenoble, that the freezing-point of a solvent as a general rule is
lowered to the same extent if there be dissolved in it quantities of
substances proportional to their molecular weights. Thus, supposing
1.80 grams of grape-sugar be dissolved in 100 grams of water and the
solution cooled below 0° with constant stirring, ice separates suddenly
in thin spicules, and the temperature rises to −0.185°. If 3.42 grams
of cane-sugar be similarly dissolved in 100 grams of water, the
freezing-point of the solution is again −0.185°. Now, 1.80 and 3.42 are
respectively the hundredth part of the molecular weights of grape-sugar
(C6H12O6) and cane-sugar (C12H22O11). Similarly, Raoult found that
quantities proportional to molecular weights dissolved in a solvent
depress the vapor pressure of that solvent equally, or, what comes to
the same thing, raise its boiling-point by an equal number of degrees.
But ordinary salts, such as sodium chloride, potassium nitrate, etc.,
dissolved in water, give too great a depression of the freezing-point
and too high a boiling-point. Next, it has been observed by botanists,
Devries, Pfeffer, and others, who had examined the ascent of sap in
plants, that if a vessel of unglazed porcelain, so treated as to cause
a film of cupric ferrocyanide (a slimy red compound) to deposit in the
pores of its walls, be filled with a weak (about 1 per cent.) solution
of sugar or similar substance, and plunged in a vessel of pure water,
water entered through the pores. By attaching a monometer to the porous
vessel the pressure exerted by the entering water could be measured.
Such pressure was termed “osmotic pressure,” referring to the “osmosis”
or passage through the walls of the vessel. Such prepared walls are
permeable freely to water, but not to sugar or similar bodies. Van’t
Hoff pointed out that the total pressure registered is proportional to
the amount of substance in solution, and that it is proportional to
the absolute temperature, and he showed, besides, that the pressure
exerted by the sugar molecules is the same as that which would be
exerted at the same temperature were an equal number of molecules of
hydrogen to occupy the same volume as the sugar solution. This may
be expressed by stating that when in dilute solution sugar molecules
behave _as if_ they were present in the gaseous state. Here again,
however, it was noticed that salts tended to give a higher pressure; it
was difficult to construct a semi-permeable diaphragm, however, which
would resist the passage of salt molecules, while allowing those of
water to pass freely. Lastly, Arrhenius, of Stockholm, had shown that
the conductivity of salt solutions for electricity may be explained on
the assumption that when a salt, such as KNO3 is dissolved in water,
it dissociates into portions similar in number and kind to those it
would yield if electrolyzed (and if no secondary reactions were to
take place). Such portions (K and NO3, for example) had been named
ions by Faraday. The conductivity of such solutions becomes greater,
per unit of dissolved salt, the weaker the solution, until finally a
limit is reached, after which further dilution no longer increases
conductivity. Now Van’t Hoff united all these isolated observations
and showed their bearing on each other. Stated shortly, the hypothesis
is as follows: When a substance is dissolved in a large quantity of
a solvent, its molecules are separated from each other to a distance
comparable with that which obtains in gases. They are, therefore,
capable of independent action; and when placed in a vessel the walls of
which are permeable to the solvent, but not to the dissolved substance
(“semi-permeable membrane”), the imprisoned molecules of the latter
exert pressure on the interior surface of these walls as if they
were gaseous. Van’t Hoff showed the intimate connection between this
phenomenon and the depression of freezing-point and the use of vapor
pressure already alluded to. He pointed out further that the exceptions
to this behavior, noticed in the case of dissolved salts, are due to
their “electric dissociation,” or “ionization,” as it is now termed;
and that in a sufficiently dilute solution of potassium nitrate,
for example, the osmotic pressure, and the correlated depression of
freezing-point and rise of boiling-point, are practically equal to what
would be produced were the salt to be split into its ions, K and NO3.
These views were vigorously advocated by Ostwald, professor at Leipzig,
in his _Zeitschrift für physikalische Chemie_, and he and his pupils
have done much to gather together facts in confirmation of this theory,
and in extending its scope.
It must be understood that the ions K and NO3 are not, strictly
speaking, atoms; they are _charged_ atoms; the K retains a +, and the
NO3 a − charge. On immersing into the solution the poles of a battery,
one charged + and the other −, the + K atoms are attracted to the −
pole, and are there discharged; as soon as they lose their charge they
are free to act on the water, when they liberate their equivalent of
hydrogen. Similarly, the − NO3 groups are discharged at the + pole, and
abstract hydrogen from the water, liberating an equivalent quantity
of oxygen. Thus the phenomenon of electrolysis, so long a mysterious
process, finds a simple explanation. The course of ordinary chemical
reactions is also readily realized when viewed in the light of this
theory. Take, for example, the ordinary equation:
AgNO3.Aq + NaCl.Ag = AgCl + NaNO3.Aq;
_i.e._, solutions of silver nitrate and sodium chloride give a
precipitate of silver chloride, leaving sodium nitrate in solution. By
the new views, such an equation must be written:
+ − + − + −
Ag.Aq + NO3.Aq + Na.Aq + Cl.Aq = AgCl + Na.Aq + NO3.Aq.
The compound, silver chloride, being insoluble in water, is formed
by the union of the ions Ag and Cl, and their consequent discharge,
forming an electrically neutral compound; while the sodium ions,
charged positively together with the NO3 ions, negatively charged,
remain in solution.
One more application of the principle may be given. Many
observers—Andrews, Favre, and Silbermann, but especially Julius
Thomsen, of Copenhagen, and M. Berthelot, of Paris—have devoted much
labor and time to the measurement of the heat evolved during chemical
reactions. Now, while very different amounts of heat are evolved when
chlorine, bromine, or iodine combine respectively with sodium or
potassium, the number of heat units evolved on neutralizing sodium
or potassium hydroxide with hydrochloric, hydrobromic, hydriodic, or
nitric acids is always about 13,500. How can this fact be explained? It
finds its explanation as follows: These acids and bases are ionized in
solution as shown in the equation:
+ − + − + −
H.Aq + Cl.Aq. + Na.Aq + OH.Aq = H.OH + Na.Aq + Cl.Aq.
Water is the only compound formed; and it is produced by the union
of the hydrogen-ion originally belonging to the acid, and the OH or
hydroxyl-ion originally belonging to the base. No further change has
occurred; hence the uniform evolution of heat by the interaction of
equivalent quantities of these acids and bases.
It now remains to give a short account of the greatest generalization
which has as yet been made in chemistry. It has been termed the
“Periodic Arrangement of the Elements.”
In 1864 Newlands, of London, and Lothar Meyer, late of Tübingen, found
that by arranging the elements in the order of their atomic weights
certain regularities were to be observed between each element, and
in general the eighth in succession from it, in the order of their
numerical value. Such similar elements formed groups or quantities;
while the elements separating them belong to a _period_, hence the name
“periodic arrangement.” Commencing with lithium, a light, lustrous
metal found in silicate in certain minerals, we have the following
series:
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
7 9.2 11 12 14 16 19 20
Sodium Magnesium Aluminum Silicon Phosphorus Sulphur Chlorine Argon
23 24.3 27 28 31 34 35.5 40
and so on. It is unnecessary to point out in detail the resemblances
between the elements which stand in the vertical columns; but it
may be stated that the resemblance extends also to the formulas and
properties of their compounds. Thus the chlorides of lithium and sodium
are each white soluble salts, of the formulas LiCl and NaCl; oxides of
magnesium and of beryllium are both insoluble white earthy powders,
MgO and BeO (GeO), and so on. Newlands, in his preliminary sketch,
termed this order the “Law of Octaves,” and predicted the existence of
certain undiscovered elements which should occupy unfilled positions in
the table. Mendeléef, professor at St. Petersburg, in 1869 amplified
and extended these relations; and he and Meyer pointed out that the
volume occupied by equal numbers of atoms of such elements underwent
a periodic variation when the elements are classified as above.
The prediction of undiscovered elements was made by Mendeléef in a
more assured manner; and in several cases they have been realized.
Thus what Mendeléef called “ekaboron” has since been discovered by
Lecoq de Boisbandron and named, patriotically, “gallium”; Mendeléef’s
“eka-silicon” is now known as “germanium,” discovered by Winkler; and
“eka-aluminum” is now Cléve’s “scandium.” Moreover, the atomic weights
of cæsium, beryllium, molybdenium, and mercury have been altered so
that they fit the periodic table; and further research has justified
the alteration.
The valency of these elements increases from right to left, as will be
seen by inspection of the following series:
LiCl BeCl3 BCl3 CCl4 NH4Cl
Na2O MgO B2O3 SiO2 PCl3
Monad. Dyad. Triad. Tetrad. Triad and Pentad.
OH2 FH Ne——
SO3 Cl(OH)O3 A——
Dyad and Hexad. Monad and Heptad. No valency.
The elements of no valency are of recent discovery. In 1894 Lord
Rayleigh had determined the density of the nitrogen of the atmosphere,
having separated from it the oxygen and carbon dioxide which is mixed
with nitrogen in air. He found it to be of somewhat higher density
than that obtainable from ammonia and other compounds of nitrogen.
In conjunction with Ramsay he investigated atmospheric nitrogen; it
was absorbed either by a method devised by Cavendish, or by making it
combine with magnesium at a red heat. They found that the unabsorbable
residue possessed an unknown spectrum, and that its density was nearly
20. To this new gas they gave the name “argon,” or inactive, seeing
that all attempts to cause it to enter into combination had failed.
In 1895 Ramsay, searching for possible combinations of argon in
minerals, experimented with one which had been previously examined by
Hillebrand, of Baltimore, and obtained from it helium, a gas of density
2, possessing a spectrum which had been previously discovered in 1868
in the chromosphere of the sun, by Jannsen, of Paris, and named helium
by Frankland and Lockyer. Subsequent liquefaction of crude argon by
means of liquid air, prepared by a process invented simultaneously by
Linde and Hampson, gave a residue which was named by its discoverers,
Ramsay and Travers, “neon.” Liquid argon has yielded two other gases
also, “krypon” and “xenon.” These elements form a separate group in the
Periodic Table, commencing with helium, with atomic weight, 4; neon,
20; argon, 40; krypon, 82; and xenon, 128. They all agree in being
mono-atomic, _i.e._, their molecules consist of single atoms; and they
have no tendency to form compounds, _i.e_., they possess no valency.
In this sketch of the progress of chemistry during the century which
has just passed, attention has been paid chiefly to the progress
of thought. Allusions must, however, be made to the applications
of chemistry to industrial purposes. The development of the soda
industry, the preparation of carbonate of soda and caustic from common
salt—initiated in France by LeBlanc (1742–1806)—has been developed
by Tennant, in Scotland, and Muspeath and Gossage, and by Hargreaves,
Weldon, and Maetea, in England; this process has at present a serious
rival in the ammonia-soda process, developed by Solway, in Belgium,
and by Brunner and Mond, in England. The main action of sulphuric
acid, so long associated with the alkali process, has made enormous
strides during the present century, but is still, in the main, the
original process of causing sulphur dioxide in presence of water to
absorb the oxygen of the air through nitric oxide. But the saving of
the oxides of nitrogen through the invention of a sulphuric acid power
by Gay-Lussac, known by his name, and the re-utilization of these
oxides in the “Glover” power, invented by John Glover, of Newcastle,
have greatly lessened the cost of the acid. Concentration of the acid
in iron vessels is now common, the cost of platinum or of fragile
glass vessels being thereby saved. The desulphurization of iron and
the removal of silicon, carbon, and phosphorus by Bessemer’s process,
modified by Thomas and Gilchrist through the introduction of a “basic
magnesia lining” for the convertors, has made it possible to obtain
pure iron and steel from ores previously regarded as of little value.
The use of artificial manures, prepared by mixing refuse animal
matters with tetra-hydrogen, calcium phosphate, and nitrate of soda,
or sulphate of ammonia, first introduced by Liebig, has created
a revolution in agricultural methods and in the weight of crops
obtainable from a given area of soil. The influence of manures on
crops has been fully studied by Lawes and Gilbert for more than fifty
years in their experimental farm at Rothampstead. The most remarkable
advances which have been made, however, are due to cheap electric
current. The electrolysis of alumina, dissolved in fused cryolite to
obtain aluminum, an operation carried out at Schaffhausen-on-the-Rhine,
and at the Falls of Foyers, in Scotland; the electro-deposition of pure
copper for electric wires and cables, electro-silvering, gilding, and
nickelling, all these are instances where decomposition of a compound
by the electric current has led to important industrial results. At
present soda and chlorine are being manufactured by the electrolysis
of salt solution contained in rocking trays, one of the electrodes
being mercury, by the Castner-Kellner process. This manufacture is
being carried on at Niagara, as well as in England. But electricity
as a heating agent finds ever-extending application. Louis Moisson,
professor at Paris, led the way by utilizing the enormous heat of
the ore in his electric furnace, thereby, among other interesting
reactions, manufacturing diamonds, small, it is true, though none the
less real. The use of electricity as a heating agent has received
new applications. Phosphorus is now made by distilling a mixture of
phosphates of lime and alumina with coke; a new polishing agent has
been found in “carborundum,” a compound of carbon and silicon, produced
by heating in an electric furnace a mixture of sand and coke; and
cyanide of potassium, almost indispensable for the extraction of gold
from ores poor in gold, is now manufactured by heating a mixture of
carbon and carbonate of barium in an electric furnace in a current
of carbon monoxide. These are but some of the instances in which
electricity has been adopted as an agent in effecting chemical changes;
and it may be confidently predicted that the earlier years of the
twentieth century will witness a great development in this direction.
It may be pointed out that the later developments of industrial
chemistry owe their success entirely to the growth of chemical theory;
and it is obvious that that nation which possesses the most competent
chemists, theoretical and practical, is destined to succeed in the
competition with other nations for commercial supremacy and all its
concomitant advantages.