The construction of the calorimeter and even the method of

determination vary considerably in different cases. Since the

beginning of the nineties, a large number of determinations

of the heat of combustion have been conducted in closed bombs

containing compressed oxygen. The greatest number of calorimetric

determinations were made by Berthelot and Thomsen. They are

given in their works _Essai de mécanique chimique fondée

sur la thermochimie_, by M. Berthelot, 1879 (2 vols.), and

_thermochemische Untersuchungen_, by J. Thomsen, 1886 (4 vols.)

The most important methods of recent thermochemistry, and all

the trustworthy results of experiment, are given in Prof. P. F.

Louginin's _Description of the Different Modes of Determining

the Heat of Combustion of Organic Compounds_, Moscow, 1894.

The student must refer to works on theoretical and physical

chemistry for a description of the elements and methods of

_thermochemistry_, into the details of which it is impossible to

enter in this work. One of the originators of thermochemistry,

Hess, was a member of the St. Petersburg Academy of Sciences.

Since 1870 a large amount of research has been carried out in

this province of chemistry, especially in France and Germany,

after the investigations of the French Academician, Berthelot,

and Professor Thomsen, of Copenhagen. Among Russians, Beketoff,

Louginin, Cheltzoff, Chroustchoff, and others are known by their

thermochemical researches. The present epoch of thermochemistry

must be considered rather as a collective one, wherein the

material of facts is amassed, and the first consequences arising

from them are noticed. In my opinion two essential circumstances

prevent the possibility of deducing any exact consequences, of

importance to chemical mechanics, from the immense store of

thermochemical data already collected: (1) The majority of the

determinations are conducted in weak aqueous solutions, and, the

heat of solution being known, are referred to the substances

in solution; yet there is much (Chapter I.) which leads to the

conclusion that in solution water does not play the simple

part of a diluting medium, but of itself acts independently

in a chemical sense on the substance dissolved. (2) Physical

and mechanical changes (decrease of volume, diffusion, and

others) invariably proceed side by side with chemical changes,

and for the present it is impossible, in a number of cases, to

distinguish the thermal effect of the one and the other kind of

change. It is evident that the one kind of change (chemical) is

essentially inseparable and incomprehensible without the other

(mechanical and physical); and therefore it seems to me that

thermochemical data will only acquire their true meaning when

the connection between the phenomena of both kinds (on the one

hand chemical and atomic, and on the other hand mechanical and

molecular or between entire masses) is explained more clearly

and fully than is at present the case. As there is no doubt that

the simple mechanical contact, or the action of heat alone,

on substances sometimes causes an evident and always a latent

(incipient) chemical change--that is, a different distribution

or motion of the atoms in the molecules--it follows that purely

chemical phenomena are inseparable from physical and mechanical

phenomena. A mechanical change may be imagined without a physical

change, and a physical without a chemical change, but it is

impossible to imagine a chemical change without a physical and

mechanical one, for without the latter we should not be able

to recognise the former, and it is by their means that we are

enabled to do so.

[39] The flame, or locality where the combustion of gases and vapours

takes place, is a complex phenomenon, 'an entire factory,' as

Faraday says, and therefore we will consider flame in some detail

in one of the following notes.

[40] If 34,500 units of heat are evolved in the combustion of 1 part

of hydrogen, and this heat is transmitted to the resulting 9

parts by weight of aqueous vapour, then we find that, taking the

specific heat of the latter as 0·475, each unit of heat raises

the temperature of 1 part by weight of aqueous vapour 2°·1 and 9

parts by weight (2·1 ÷ 9) O°·23; hence the 34,500 units of heat

raise its temperature 7,935°. If detonating gas is converted

into water in a closed space, then the aqueous vapour formed

cannot expand, and therefore, in calculating the temperature of

combustion, the specific heat at a constant volume must be taken

into consideration; for aqueous vapour it is 0·36. This figure

gives a still higher temperature for the flame. In reality it

is much lower, but the results given by different observers are

very contradictory (from 1,700° to 2,400°), the discrepancies

depending on the fact that flames of different sizes are cooled

by radiation to a different degree, but mainly on the fact that

the methods and apparatus (pyrometers) for the determination

of high temperatures, although they enable relative changes of

temperature to be judged, are of little use for determining their

absolute magnitude. By taking the temperature of the flame of

detonating gas as 2,000°, I give, I think, the average of the

most trustworthy determinations and calculations based upon the

determination of the variation of the specific heat of aqueous

vapour and other gases (_see_ Chapter XLI.)

[41] It is evident that not only hydrogen, but every other combustible

gas, will give an explosive mixture with oxygen. For this reason

coal-gas mixed with air explodes when the mixture is ignited.

The pressure obtained in the explosions serves as the _motive

power of gas engines_. In this case advantage is taken, not only

of the pressure produced by the explosion, but also of that

contraction which takes place after the explosion. On this is

based the construction of several motors, of which Lenoir's was

formerly, and Otto's is now, the best known. The explosion is

usually produced by coal-gas and air, but of late the vapours of

combustible liquids (kerosene, benzene) are also being employed

in place of gas (Chapter IX.) In Lenoir's engine a mixture of

coal-gas and air is ignited by means of sparks from a Ruhmkorff's

coil, but in the most recent machines the gases are ignited by

the direct action of a gas jet, or by contact with the hot walls

of a side tube.

[Illustration: FIG. 34.--Safety burner for detonating gas, described in

text.]

Mixtures of hydrogen and of various other gases with oxygen are taken

advantage of for obtaining high temperatures. By the aid of such high

temperatures metals like platinum may be melted on a large scale,

which cannot be performed in furnaces heated with charcoal and fed by

a current of air. The burner, shown in fig. 34, is constructed for

the application of detonating gas to the purpose. It consists of two

brass tubes, one fixed inside the other, as shown in the drawing. The

internal central tube C C conducts oxygen, and the outside, enveloping,

tube E' E' conducts hydrogen. Previous to their egress the gases

do not mix together, so that there can be no explosion inside the

apparatus. When this burner is in use C is connected with a gas-holder

containing oxygen, and E with a gas-holder containing hydrogen (or

sometimes coal-gas). The flow of the gases can be easily regulated by

the stopcocks O H. The flame is shortest and evolves the greatest heat

when the gases burning are in the proportion of 1 volume of oxygen to

2 volumes of hydrogen. The degree of heat may be easily judged from

the fact that a thin platinum wire placed in the flame of a properly

proportioned mixture easily melts. By placing the burner in the orifice

of a hollow piece of lime, a crucible A B is obtained in which the

platinum may be easily melted, even in large quantities if the current

of oxygen and hydrogen be sufficiently great (Deville). The flame

of detonating gas may also be used for illuminating purposes. It is

by itself very pale, but owing to its high temperature it may serve

for rendering infusible objects incandescent, and at the very high

temperature produced by the detonating gas the incandescent substance

gives a most intense light. For this purpose lime, magnesia, or oxide of

zirconium are used, as they are not fusible at the very high temperature

evolved by the detonating gas. A small cylinder of lime placed in

the flame of detonating gas, if regulated to the required point,

gives a very brilliant white light, which was at one time proposed

for illuminating lighthouses. At present in the majority of cases

the electric light, owing to its constancy and other advantages, has

replaced it for this purpose. The light produced by the incandescence of

lime in detonating gas is called the _Drummond light_ or _limelight_.

The above cases form examples of the combustion of elements in oxygen,

but exactly similar phenomena are observed in the _combustion of

compounds_. So, for instance, the solid, colourless, shiny substance,

naphthalene, C_{10}H_{8}, burns in the air with a smoky flame, whilst

in oxygen it continues to burn with a very brilliant flame. Alcohol,

oil, and other substances burn brilliantly in oxygen on conducting the

oxygen by a tube to the flame of lamps burning these substances. A high

temperature is thus evolved, which is sometimes taken advantage of in

chemical practice.

In order to understand why combustion in oxygen proceeds more rapidly,

and is accompanied by a more intense heat effect, than combustion in

air, it must be recollected that air is oxygen diluted with nitrogen,

which does not support combustion, and therefore fewer particles of

oxygen flow to the surface of a substance burning in air than when

burning in pure oxygen, besides which the reason of the intensity of

combustion in oxygen is the high temperature acquired by the substance

burning in it.[41 bis]

[41 bis] Let us consider as an example the combustion of

sulphur in air and in oxygen. If 1 gram of sulphur burns in air

or oxygen it evolves in either case 2250 units of heat--_i.e._

evolves sufficient heat for heating 2,250 grams of water 1°

C. This heat is first of all transmitted to the sulphurous

anhydride, SO_{2}, formed by the combination of sulphur with

oxygen. In its combustion 1 gram of sulphur forms 2 grams of

sulphurous anhydride--_i.e._ the sulphur combines with 1 gram of

oxygen. In order that 1 gram of sulphur should have access to 1

gram of oxygen in air, it is necessary that 3·4 grams of nitrogen

should simultaneously reach the sulphur, because air contains

seventy-seven parts of nitrogen (by weight) per twenty-three

parts of oxygen. Thus in the combustion of 1 gram of sulphur,

the 2,250 units of heat are transmitted to 2 grams of sulphurous

oxide and to at least 3·4 grams of nitrogen. As 0·155 unit of

heat is required to raise 1 gram of sulphurous anhydride 1°

C., therefore 2 grams require 0·31 unit. So also 3·4 grams of

nitrogen require 3·4 × 0·244 or 0·83 unit of heat, and therefore

in order to raise both gases 1° C. 0·31 + 0·83 or 1·14 unit of

heat is required; but as the combustion of the sulphur evolves

2,250 units of heat, therefore the gases might be heated (if

their specific heats remained constant) to 2250/1·14 or 1,974°

C. That is, the maximum possible temperature of the flame of the

sulphur burning in air will be 1,974° C. In the combustion of the

sulphur in oxygen the heat evolved (2,250 units) can only pass to

the 2 grams of sulphurous anhydride, and therefore the highest

possible temperature of the flame of the sulphur in oxygen will

be = 2250/0·31 or 7258°. In the same manner it may be calculated

that the temperature of charcoal burning in air cannot exceed

2,700°, while in oxygen it may attain 10,100° C. For this reason

the temperature in oxygen will always be higher than in air,

although (judging from what has been said respecting detonating

gas) neither one temperature nor the other will ever approximate

to the theoretical amount.

[Illustration: FIG. 35.--Faraday's experiment for investigating the

different parts of a candle flame.]

Among the phenomena accompanying the combustion of certain substances,

the _phenomenon of flame_ attracts attention. Sulphur, phosphorus,

sodium, magnesium, naphthalene, &c., burn like hydrogen with a flame,

whilst in the combustion of other substances no flame is observed, as,

for instance, in the combustion of iron and of charcoal. The appearance

of flame depends on the capacity of the combustible substance to yield

gases or vapours at the temperature of combustion. At the temperature

of combustion, sulphur, phosphorus, sodium, and naphthalene pass into

vapour, whilst wood, alcohol, oil, &c., are decomposed into gaseous

and vaporous substances. The combustion of gases and vapours forms

flames, and therefore _a flame is composed of the hot and incandescent

gases and vapours produced by combustion_. It may easily be proved that

the flames of such non-volatile substances as wood contain volatile

and combustible substances formed from them, by placing a tube in the

flame connected with an aspirator. Besides the products of combustion,

combustible gases and liquids, previously in the flame as vapours,

collect in the aspirator. For this experiment to succeed--_i.e._ in

order to really extract combustible gases and vapours from the flame

it is necessary that the suction tube should be placed _inside_ the

flame. The combustible gases and vapours can only remain unburnt inside

the flame, for at the surface of the flame they come into contact with

the oxygen of the air and burn.[42] Flames are of different degrees

of _brilliancy_, according to whether _solid_ incandescent particles

occur in the combustible gas or vapour, or not. Incandescent gases

and vapours emit but little light by themselves, and therefore give

a paler flame.[43] If a flame does not contain solid particles it is

transparent, pale, and emits but little light.[44] The flames of

burning alcohol, sulphur, and hydrogen are of this kind. A pale flame

may be rendered luminous by placing fine particles of solid matter in

it. Thus, if a very fine platinum wire be placed in the pale flame

of burning alcohol--or, better still, of hydrogen--the flame emits

a bright light. This is still better seen by sifting the powder of

an incombustible substance, such as fine sand, into the flame, or

by placing a bunch of asbestos threads in it. Every brilliant flame

always contains some kind of solid particles, or at least some very

dense vapour. The flame of sodium burning in oxygen has a brilliant

yellow colour, from the presence of particles of solid sodium oxide.

The flame of magnesium is brilliant from the fact that in burning

it forms solid magnesia, which becomes white hot, and similarly the

brilliancy of the Drummond light is due to the heat of the flame

raising the solid non-volatile lime to a state of incandescence.

The flames of a candle, wood, and similar substances are brilliant,

because they contain particles of charcoal or soot. It is not the flame

itself which is luminous, but the incandescent soot it contains. These

particles of charcoal which occur in flames may be easily observed

by introducing a cold object, like a knife, into the flame.[45] The

particles of charcoal burn at the outer surface of the flame if the

supply of air be sufficient, but if the supply of air--that is, of

oxygen--be insufficient for their combustion the flame smokes, because

the unconsumed particles of charcoal are carried off by the current of

air.[46]

[42] Faraday proved this by a very convincing experiment on

a candle flame. If one arm of a bent glass tube be placed in a

candle flame above the wick in the dark portion of the flame,

then the products of the partial combustion of the stearin will

pass up the tube, condense in the other arm, and collect in a

flask placed under it (fig. 35) as heavy white fumes which burn

when lighted. If the tube be raised into the upper luminous

portion of the flame, then a dense black smoke which will not

inflame accumulates in the flask. Lastly, if the tube be let down

until it touches the wick, then little but stearic acid condenses

in the flask.

[43] All transparent substances which transmit light with great ease

(that is, which absorb but little light) are but little luminous

when heated; so also substances which absorb but few heat rays,

when heated transmit few rays of heat.

[44] There is, however, no doubt but that very heavy dense vapours or

gases under pressure (according to the experiments of Frankland)

are luminous when heated, because, as they become denser they

approach a liquid or solid state. Thus detonating gas when

exploded under pressure gives a brilliant light.

[45] If hydrogen gas be passed through a volatile liquid hydrocarbon--for

instance, through benzene (the benzene may be poured directly

into the vessel in which hydrogen is generated)--then its vapour

burns with the hydrogen and gives a very bright flame, because

the resultant particles of carbon (soot) become incandescent.

Benzene, or platinum gauze, introduced into a hydrogen flame may

be employed for illuminating purposes.

[46] In _flames_ the separate parts may be distinguished with more or

less distinctness. That portion of the flame whither the

combustible vapours or gases flow, is not luminous because its

temperature is still too low for the process of combustion to

take place in it. This is the space which in a candle surrounds

the wick, or in a gas jet is immediately above the orifice from

which the gas escapes. In a candle the combustible vapours and

gases which are formed by the action of heat on the melted

tallow or stearin rise in the wick, and are heated by the high

temperature of the flame. By the action of the heat, the solid or

liquid substance is here, as in other cases, decomposed, forming

products of dry distillation. These products occur in the central

portion of the flame of a candle. The air travels to it from the

outside, and is not able to intermix at once with the vapours

and gases in all parts of the flame equally; consequently, in

the outer portion of the flame the amount of oxygen will be

greater than in the interior portions. But, owing to diffusion,

the oxygen, of course mixed with nitrogen, flowing towards the

combustible substance, does finally penetrate to the interior

of the flame (when the combustion takes place in ordinary air).

The combustible vapours and gases combine with this oxygen,

evolve a considerable amount of heat, and bring about that state

of incandescence which is so necessary both for keeping up the

combustion and also for the uses to which the flame is applied.

Passing from the colder envelope of air through the interior

of the flame, to the source of the combustible vapours (for

instance, the wick), we evidently first traverse layers of higher

and higher temperature, and then portions which are less and

less hot, in which the combustion is less complete, owing to the

limited supply of oxygen.

[Illustrationtion: FIG. 36.--In the candle flame the portion C

contains the vapours and products of decomposition; in the bright

zone A the combustion has commenced, and particles of carbon are

emitted; and in the pale zone B the combustion is completed.]

Thus unburnt products of the decomposition of organic substances

occur in the interior of the flame. But there is always free

hydrogen in the interior of the flame, even when oxygen is

introduced there, or when a mixture of hydrogen and oxygen burns,

because the temperature evolved in the combustion of hydrogen

or the carbon of organic matter is so high that the products

of combustion are themselves partially decomposed--that is,

dissociated--at this temperature. Hence, in a flame a portion

of the hydrogen and of the oxygen which might combine with

the combustible substances must always be present in a free

state. If a hydrocarbon burns, and we imagine that a portion

of the hydrogen is in a free state, then a portion of the

carbon must also occur in the same form in the flame, because,

other conditions being unchanged, carbon burns after hydrogen,

and this is actually observed in the combustion of various

hydrocarbons. Charcoal, or the soot of a common flame, arises

from the dissociation of organic substances contained in the

flame. The majority of hydrocarbons, especially those containing

much carbon--for instance, naphthalene--burn, even in oxygen,

with separation of soot. In that portion of the flame where the

hydrogen burns the carbon remains unburnt, or at least partly so.

It is this free carbon which causes the brilliancy of the flame.

That the interior of the flame contains a mixture which is still

capable of combustion may be proved by the following experiment:

A portion of the gases may be withdrawn by an aspirator from

the central portion of the flame of carbonic oxide, which is

combustible in air. For this purpose Deville passed water

through a metallic tube having a fine lateral orifice, which is

placed in the flame. As the water flows along the tube portions

of the gases of the flame enter, and, passing along the tube

alternately with cylinders of water, are carried away into an

apparatus where they can be investigated. It appears that all

portions of the flame obtained by the combustion of a mixture

of carbonic oxide and oxygen contain a portion of this mixture

still unburnt. The researches of Deville and Bunsen showed that

in the explosion of a mixture of hydrogen and of carbonic oxide

with oxygen in a closed space, complete combustion does not ever

take place immediately. If two volumes of hydrogen and one volume

of oxygen be confined in a closed space, then on explosion the

pressure does not attain that magnitude which it would were there

immediate and complete combustion. It may be calculated that the

pressure should attain twenty-six atmospheres. In reality, it

does not exceed nine and a half atmospheres.

Hence the admixture of the products of combustion with an

explosive mixture prevents the combustion of the remaining

mass, although capable of burning. The admixture of carbonic

anhydride prevents carbonic oxide from burning. The presence of

any other foreign gas interferes in the same manner. This shows

that every portion of a flame must contain combustible, burning,

and already burnt substances--_i.e._ oxygen, carbon, carbonic

oxide, hydrogen, hydrocarbons, carbonic anhydride, and water.

Consequently, _it is impossible to attain instantaneous complete

combustion_, and this is one of the reasons of the phenomenon

of flame. A certain space is required, and the temperature must

be unequal in different parts of it. In this space different

quantities of the component parts are successively subjected

to combustion, or are cooled under the influence of adjacent

objects, and combustion only ends where the flame ends. If

the combustion could be concentrated at one spot, then the

temperature would be incomparably higher than it is under the

actual circumstances.

The various regions of the flame have formed the frequent

subject of experimental research, and the experiments conducted

by Smithells and Ingle (1892) are particularly instructive;

they show that the reducing (interior) and oxidising (exterior)

portions of the flame of a burning gas may be divided by taking a

Bunsen burner and surrounding the flame of the gas burnt in it,

by another wider tube (without the access of air to the annular

space or allowing only a small current of air to pass), when

a gaseous mixture, containing oxide of carbon and capable of

further combustion, will issue from this enveloping tube, so that

a second flame, corresponding to the exterior (oxidising) portion

of an ordinary flame, may be obtained above the enveloping tube.

This division of the flame into two portions is particularly

clear when cyanogen C_{2}N_{2} is burnt, because the interior

portion (where CO is chiefly formed according to the equation

C_{2}N_{2} + O_{2} = 2CO + N_{2}, but a portion of the nitrogen

is oxidised) is of a rose colour, while the exterior portion

(where the CO burns into CO_{2} at the expense of a fresh

quantity of oxygen and of the oxides of nitrogen proceeding from

the interior portions) is of a bluish-grey colour.

The combination of various substances with oxygen may not present

any signs of combustion--that is, the temperature may rise but

inconsiderably. This may either proceed from the fact that the

reaction of the substance (for example, tin, mercury, lead at a

high temperature, or a mixture of pyrogallol with caustic potash at

the ordinary temperature) evolves but little heat, or that the heat

evolved is transmitted to good conductors of heat, like metals, or

that the combination with oxygen takes place so slowly that the heat

evolved succeeds in passing to the surrounding objects. Combustion

is only a particular, intense, and evident case of combination with

oxygen. Respiration is also an act of combination with oxygen; it

also serves, like combustion, for the development of heat by those

chemical processes which accompany it (the transformation of oxygen into

carbonic anhydride). Lavoisier enunciated this in the lucid expression,

'respiration is slow combustion.'

Reactions involving slow combination of substances with oxygen are

termed _oxidations_. Combination of this kind (and also combustion)

often results in the formation of acid substances, and hence the name

_oxygen_ (_Sauerstoff_). Combustion is only rapid oxidation. Phosphorus,

iron, and wine may be taken as examples of substances which slowly

oxidise in air at the ordinary temperature. If such a substance be

left in contact with a definite volume of air or oxygen, it absorbs

the oxygen little by little, as may be seen by the decrease in volume

of the gas. This slow oxidation is not often accompanied by a sensible

evolution of heat; an evolution of heat really does occur, only it

is not apparent to our senses owing to the small rise in temperature

which takes place; this is owing to the slow rate of the reaction and

to the transmission of the heat formed as radiant heat, &c. Thus, in

the oxidation of wine and its transformation into vinegar by the usual

method of preparation of the latter, the heat evolved cannot be observed

because it extends over several weeks, but in the so-called rapid

process of the manufacture of vinegar, when a large quantity of wine is

comparatively rapidly oxidised, the evolution of heat is quite apparent.

Such slow processes of oxidation are always taking place in nature by

the action of the atmosphere. Dead organisms and the substances obtained

from them--such as bodies of animals, wood, wool, grass, &c.--are

especially subject to this action. They _rot_ and _decompose_--that is,

their solid matter is transformed into gases, under the influence of

moisture and atmospheric oxygen, and generally under the influence of

other organisms, such as moulds, worms, micro-organisms (bacteria), and

the like. These are processes of slow combustion, of slow combination

with oxygen. It is well known that manure rots and develops heat,

that stacks of damp hay, damp flour, straw, &c., become heated and

are changed in the process.[47] In all these transformations the same

chief products of combustion are formed as those which are contained

in smoke; the carbon gives carbonic anhydride, and the hydrogen water.

Hence these processes require oxygen just like combustion. This is

the reason why the entire prevention of access of air hinders these

transformations,[48] and an increased supply of air accelerates them.

The mechanical treatment of arable lands by the plough, harrow, and

other similar means has not only the object of facilitating the spread

of roots in the ground, and of making the soil more permeable to water,

but it also serves to facilitate the access of the air to the component

parts of the soil; as a consequence of which the organic remains of soil

rot--so to speak, breathe air and evolve carbonic anhydride. One acre

of good garden land in the course of a summer evolves more than sixteen

tons of carbonic anhydride.

[47] Cotton waste (used in factories for cleaning machines from

lubricating oil) soaked in oil and lying in heaps is

self-combustible, being oxidised by the air.

[48] When it is desired to preserve a supply of vegetable and animal

food, the access of the oxygen of the atmosphere (and also of

the germs of organisms present in the air) is often prevented.

With this object articles of food are often kept in hermetically

closed vessels, from which the air has been withdrawn; vegetables

are dried and soldered up while hot in tin boxes; sardines are

immersed in oil, &c. The removal of water from substances is also

sometimes resorted to with the same object (the drying of hay,

corn, fruits), as also is saturation with substances which absorb

oxygen (such as sulphurous anhydride), or which hinder the growth

of organisms forming the first cause of putrefaction, as in

processes of smoking, embalming, and in the keeping of fishes and

other animal specimens in spirit, &c.

It is not only vegetable and animal substances which are subject to slow

oxidation in the presence of water. Some metals even rust under these

conditions. Copper very easily absorbs oxygen in the presence of acids.

Many metallic sulphides (for example, pyrites) are very easily oxidised

with access of air and moisture. Thus processes of slow oxidation

proceed throughout nature. However, there are many elements which do

not under any circumstances combine directly with gaseous oxygen;

nevertheless their compounds with oxygen may be obtained. Platinum,

gold, iridium, chlorine, and iodine are examples of such elements. In

this case recourse is had to a so-called _indirect method_--_i.e._ the

given substance is combined with another element, and by a method of

double decomposition this element is replaced by oxygen. Substances

which do not directly combine with oxygen, but form compounds with it

by an indirect method, often readily lose the oxygen which they had

absorbed by double decomposition or at the moment of its evolution.

Such, for example, are the compounds of oxygen with chlorine, nitrogen,

and platinum, which evolve oxygen on heating--that is, they may be

used as oxidising agents. In this respect _oxidising agents_, or those

compounds of oxygen which are employed in chemical and technical

practice for transferring oxygen to other substances, are especially

remarkable. The most important among these is nitric acid or _aqua

fortis_--a substance rich in oxygen, and capable of evolving it when

heated, which easily oxidises a great number of substances. Thus nearly

all metals and organic substances containing carbon and hydrogen are

more or less oxidised when heated with nitric acid. If strong nitric

acid be taken, and a piece of burning charcoal be immersed in the acid,

it continues to burn. Chromic acid acts like nitric acid; alcohol burns

when mixed with it. Although the action is not so marked, even water

may oxidise with its oxygen. Sodium is not oxidised in perfectly dry

oxygen at the ordinary temperature, but it burns very easily in water

and aqueous vapour. Charcoal can burn in carbonic anhydride--a product

of combustion--forming carbonic oxide. Magnesium burns in the same gas,

separating carbon from it. Speaking generally, combined oxygen can pass

from one compound to another.

The products of combustion or oxidation--and in general the definite

compounds of oxygen--are termed _oxides_. Some oxides are not capable of

combining with other oxides--or combine with only a few, and then with

the evolution of very little heat; others, on the contrary, enter into

combination with very many other oxides, and in general have remarkable

chemical energy. The oxides incapable of combining with others, or only

showing this quality in a small degree, are termed _indifferent oxides_.

Such are the peroxides, of which mention has before been made.

The class of oxides capable of entering into mutual combination we will

term _saline oxides_. They fall into two chief groups--at least, as

regards the most extreme members. The members of one group combine with

the members of the other group with particular ease. As representative

of one group may be taken the oxides of the metals, magnesium, sodium,

calcium, &c. Representatives of the other group are the oxides formed

by the non-metals, sulphur, phosphorus, carbon. Thus, if we take the

oxide of calcium, or lime, and bring it into contact with oxides of the

second group, combination very readily ensues. For instance, if we mix

calcium oxide with oxide of phosphorus they combine with great facility

and with the evolution of much heat. If we pass the vapour of sulphuric

anhydride, obtained by the combination of sulphurous oxide with oxygen,

over pieces of lime heated to redness, the sulphuric anhydride is

absorbed by the lime with the formation of a substance called calcium

sulphate. The oxides of the first kind, which contain metals, are

termed _basic oxides_ or _bases_. Lime is a familiar example of this

class. The oxides of the second group, which are capable of combining

with the bases, are termed _anhydrides of the acids_ or _acid oxides_.

Sulphuric anhydride, SO_{3}, may be taken as a type of the latter group.

It is a compound of sulphur with oxygen formed not directly but by the

addition of a fresh quantity of oxygen to sulphurous anhydride, SO_{2},

by passing it together with oxygen over incandescent spongy platinum.

Carbonic anhydride (often termed 'carbonic acid'), CO_{2}, phosphoric

anhydride, sulphurous anhydride, are all acid oxides, for they can

combine with such oxides as lime or calcium oxide, magnesia or magnesium

oxide, MgO, soda or sodium oxide, Na_{2}O, &c.

If a given element form but one basic oxide, it is termed the _oxide_;

for example, calcium oxide, magnesium oxide, potassium oxide. Some

indifferent oxides are also called 'oxides' if they have not the

properties of peroxides, and at the same time do not show the properties

of acid anhydrides--for example, carbonic oxide, of which mention

has already been made. If an element forms two basic oxides (or two

indifferent oxides not having the characteristics of a peroxide) then

that of the lower degree of oxidation is called a _suboxide_--that is,

suboxides contain less oxygen than oxides. Thus, when copper is heated

to redness in a furnace it increases in weight and absorbs oxygen,

until for 63 parts of copper there is absorbed not more than 8 parts

of oxygen by weight, forming a red mass, which is suboxide of copper;

but if the roasting be prolonged, and the draught of air increased,

63 parts of copper absorb 16 parts of oxygen, and form black oxide of

copper. Sometimes to distinguish between the degrees of oxidation a

change of suffix is made in the oxidised element, _-ic_ oxide denoting

the higher degree of oxidation, and _-ous_ oxide the lower degree.

Thus ferrous oxide and ferric oxide are the same as suboxide of iron

and oxide of iron. If an element forms one anhydride only, then it is

named by an adjective formed from the name of the element made to end in

_-ic_ and the word _anhydride_. When an element forms two anhydrides,

then the suffixes _-ous_ and _-ic_ are used to distinguish them: _-ous_

signifying less oxygen than _-ic_; for example, sulphurous and sulphuric

anhydrides.[49] When several oxides are formed from the same element,

the prefixes _mon_, _di_, _tri_, _tetra_ are used, thus: chlorine

monoxide, chlorine dioxide, chlorine trioxide, and chlorine tetroxide

or chloric anhydride.

[49] It must be remarked that certain elements form oxides of all three

kinds--_i.e._ indifferent, basic, and acid; for example,

manganese forms manganous oxide, manganic oxide, peroxide of

manganese, red oxide of manganese, and manganic anhydride,

although some of them are not known in a free state but only

in combination. The basic oxides contain less oxygen than the

peroxides, and the peroxides less than the acid anhydrides.

Thus they must be placed in the following general normal

order with respect to the amount of oxygen entering into

their composition--(1) basic oxides, suboxides, and oxides;

(2) peroxides; (3) acid anhydrides. The majority of elements,

however, do not give all three kinds of oxides, some giving only

one degree of oxidation. It must further be remarked that there

are oxides formed by the combination of acid anhydrides with

basic oxides, or, in general, of oxides with oxides. For every

oxide having a higher and a lower degree of oxidation, it might

be said that the intermediate oxide was formed by the combination

of the higher with the lower oxide. But this is not true in all

cases--for instance, when the oxide under consideration forms

a whole series of independent compounds--for oxides which are

really formed by the combination of two other oxides do not give

such independent compounds, but in many cases decompose into the

higher and lower oxides.

The oxides themselves rarely undergo chemical transformations, and in

the few cases where they are subject to such changes a particularly

important part is played by their combinations with water. The majority

of, if not all, basic and acid oxides combine with water, either by a

direct or an indirect method forming _hydrates_--that is, compounds

which split up into water and an oxide of the same kind only. It is

well known that many substances are capable of combining with water.

Oxides possess this property in the highest degree. We have already

seen examples of this (Chapter I.) in the combination of lime, and

of sulphuric and phosphoric anhydrides, with water. The resulting

combinations are basic and acid hydrates. Acid hydrates are called

_acids_ because they have an acid taste when dissolved in water (or

saliva), for then only can they act on the palate. Vinegar, for example,

has an acid taste because it contains acetic acid dissolved in water.

Sulphuric acid, to which we have frequently referred, because it is the

acid of the greatest importance both in practical chemistry and for its

technical applications, is really a hydrate formed by the combination

of sulphuric anhydride with water. Besides their acid taste, dissolved

acids or acid hydrates have the property of changing the blue colour of

certain vegetable dyes to red. Of these dyes _litmus_ is particularly

remarkable and much used. It is the blue substance extracted from

certain lichens, and is used for dyeing tissues blue; it gives a blue

infusion with water. This infusion, on the addition of an acid, _changes

from blue to red_.[50]

[50] Blotting or unsized paper, soaked in a solution of litmus, is

usually employed for detecting the presence of acids. This paper

is cut into strips, and is called _test paper_; when dipped into

acid it immediately turns red. This is a most sensitive reaction,

and may be employed for testing for the smallest traces of

acids. If 10,000 parts by weight of water be mixed with 1 part

of sulphuric acid, the coloration is distinct, and it is even

perceptible on the addition of ten times more water. Certain

precautions must, however, be taken in the preparation of such

very sensitive litmus paper. Litmus is sold in lumps. Take, say,

100 grams of it; powder it, and add it to cold pure water in

a flask; shake and decant the water. Repeat this three times.

This is done to wash away easily-soluble impurities, especially

alkalis. Transfer the washed litmus (it is washed with absolute

alcohol to remove the non-sensitive reddish colouring matter) to

a flask, and pour in 600 c.c. of water, heat, and allow the hot

infusion to remain for some hours in a warm place. Then filter,

and divide the filtrate into two parts. Add a few drops of nitric

acid to one portion, so that a faint red tinge is obtained, and

then mix the two portions. Add spirit to the mixture, and keep

it in a stoppered bottle (it soon spoils if left open to the

air). This infusion may be employed directly; it reddens in the

presence of acids, and turns blue in the presence of alkalis. If

evaporated, a solid mass is obtained which is soluble in water,

and may be kept unchanged for any length of time. The test paper

may be prepared as follows:--Take a strong infusion of litmus,

and soak blotting-paper with it; dry it, and cut it into strips,

and use it as test-paper for acids. For the detection of alkalis,

the paper must be soaked in a solution of litmus just reddened by

a few drops of acid; if too much acid be taken, the paper will

not be sensitive. Such acids as sulphuric acid colour litmus,

and especially its infusion, a brick-red colour, whilst more

feeble acids, such as carbonic, give a faint red-wine tinge.

Test-paper of a yellow colour is also employed; it is dyed by

an infusion of turmeric roots in spirit. In alkalis it turns

brown, but regains its original hue in acids. Many blue and other

vegetable colouring matters may be used for the detection of

acids and alkalis; for example, infusions of cochineal, violets,

log-wood, &c. Certain artificially prepared substances and dyes

may also be employed. Thus rosolic acid, C_{20}H_{16}O_{3} and

phenolphthaleïn, C_{20}H_{14}O_{4} (it is used in an alcoholic

solution, and is not suitable for the detection of ammonia), are

colourless in an acid, and red in an alkaline, solution. Cyanine

is also colourless in the presence of acids, and gives a blue

coloration with alkalis. Methyl-orange (yellow in an aqueous

solution) is not altered by alkalis but becomes pink with acids

(weak acids have no action), &c. These are very sensitive tests.

Their behaviour in respect to various acids, alkalis, and salts

sometimes give the means of distinguishing substances from each

other.

Basic oxides, in combining with water, form hydrates, of which, however,

very few are soluble in water. Those which are soluble in water have

an alkaline taste like that of soap or of water in which wood ashes

have been boiled, and are called _alkalis_. Further, alkalis have the

property of restoring the blue colour to litmus which has been reddened

by the action of acids. The hydrates of the oxides of sodium and

potassium, NaHO and KHO, are examples of basic hydrates easily soluble

in water. They are true alkalis, and are termed _caustic_, because they

act very powerfully on the skin of animals and plants. Thus NaHO is

called 'caustic' soda.

The saline oxides are capable of combining together and with water.

Water itself is an oxide, and not an indifferent one, for it can, as we

have seen, combine with basic and acid oxides; it is a representative

of a whole series of saline oxides, _intermediate oxides_, capable of

combining with both basic and acid oxides. There are many such oxides,

which, like water, combine with basic and acid anhydrides--for instance,

the oxides of aluminium and tin, &c. From this it may be concluded that

all oxides might be placed, in respect to their capacity for combining

with one another, in one uninterrupted series, at one extremity of which

would stand those oxides which do not combine with the bases--that

is, the alkalis--while at the other end would be the acid oxides, and

in the interval those oxides which combine with one another and with

both the acid and basic oxides. The further apart the members of this

series are, the more stable are the compounds they form together, the

more energetically do they act on each other, the greater the quantity

of heat evolved in their reaction, and the more marked is their saline

chemical character.

We said above that basic and acid oxides combine together, but rarely

react on each other; this depends on the fact that the majority of them

are solids or gases--that is, they occur in the state least prone

to chemical reaction. The gaseo-elastic state is with difficulty

destroyed, because it necessitates overcoming the elasticity proper

to the gaseous particles. The solid state is characterised by the

immobility of its particles; whilst chemical action requires contact,

and hence a displacement and mobility. If solid oxides be heated, and

especially if they be melted, then reaction proceeds with great ease.

But such a change of state rarely occurs in nature or in practice.

Only in a few furnace processes is this the case. For example; in the

manufacture of glass, the oxides contained in it combine together in

a molten state. But when oxides combine with water, and especially

when they form hydrates soluble in water, then the mobility of their

particles increases to a considerable extent, and their reaction

is greatly facilitated. Reaction then takes place at the ordinary

temperature--easily and rapidly; so that this kind of reaction belongs

to the class of those which take place with unusual facility, and are,

therefore, very often taken advantage of in practice, and also have

been and are going on in nature at every step. We will now consider the

reactions of oxides in the state of hydrates, not losing sight of the

fact that water is itself an oxide with definite properties, and has,

therefore, no little influence on the course of those changes in which

it takes part.

If we take a definite quantity of an acid, and add an infusion of litmus

to it, it turns red; the addition of an alkaline solution does not

immediately alter the red colour of the litmus, but on adding more and

more of the alkaline solution a point is reached when the red colour

changes to violet, and then the further addition of a fresh quantity

of the alkaline solution changes the colour to blue. This change of

the colour of the litmus is a consequence of the formation of a new

compound. This reaction is termed the _saturation_ or _neutralisation_

of the acid by the base, or _vice versâ_. The solution in which the acid

properties of the acid are saturated by the alkaline properties of the

base is termed a _neutral_ solution. Such a solution, although derived

from the mixture of a base with an acid, does not exhibit either the

acid or basic reaction on litmus, yet it preserves many other signs of

the acid and alkali. It is observed that in such a definite admixture

of an acid with an alkali, besides the changes in the colour of litmus

there is a heating effect--_i.e._ an evolution of heat--which is alone

sufficient to prove that there was chemical action. And, indeed, if the

resultant violet solution be evaporated, there separates out, not the

acid or the alkali originally taken, but a substance which has neither

acid nor alkaline properties, but is usually solid and crystalline,

having a saline appearance; this is a _salt_ in the chemical sense of

the word. Hence a salt is derived from the reaction of an acid on an

alkali, in a certain definite proportion. The water here taken for

solution plays no other part than merely facilitating the progress

of the reaction. This is seen from the fact that the anhydrides of

the acids are able to combine with basic oxides, and give the same

salts as do the acids with the alkalis or hydrates. Hence, a salt is a

compound of definite quantities of an acid with an alkali. In the latter

reaction, water is separated out if the substance formed be the same

as is produced by the combination of anhydrous oxides together.[51]

Examples of the formation of salts from acids and bases are easily

observed, and are very often applied in practice. If we take, for

instance, insoluble magnesium oxide (magnesia) it is easily dissolved

in sulphuric acid, and on evaporation gives a saline substance, bitter,

like all the salts of magnesium, and familiar to all under the name of

Epsom salts, used as a purgative. If a solution of caustic soda--which

is obtained, as we saw, by the action of water on sodium oxide--be

poured into a flask in which charcoal has been burnt; or if carbonic

anhydride, which is produced under so many circumstances, be passed

through a solution of caustic soda, then sodium carbonate or soda,

Na_{2}CO_{3}, is obtained, of which we have spoken several times, and

which is prepared on a large scale and often used in manufactures. This

reaction is expressed by the equation, 2NaHO + CO_{2} = Na_{2}CO_{3} +

H_{2}O. Thus, the various bases and acids form an innumerable number of

different salts.[52] Salts constitute an example of definite chemical

compounds, and both in the history and practice of science are most

often cited as confirming the conception of definite chemical compounds.

Indeed, all the indications of a definite chemical combination are

clearly seen in the formation and properties of salts. Thus, they are

produced with a definite proportion of oxides, heat is evolved in their

formation,[53] and the chemical character of the oxides and many of the

physical properties become hidden in their salts. For example, when

gaseous carbonic anhydride combines with a base to form a solid salt,

the elasticity of the gas quite disappears in its passage into the

salt.[54]

[51] That water really is separated in the reaction of acid on alkaline

hydrates, may be shown by taking some other intermediate

hydrate--for example, alumina--instead of water. Thus, if a

solution of alumina in sulphuric acid be taken, it will have,

like the acid, an acid reaction, and will therefore colour

litmus red. If, on the other hand, a solution of alumina in an

alkali--say, potash--be taken, it will have an alkaline reaction,

and will turn red litmus blue. On adding the alkaline to the

acid solution until neither an alkaline nor an acid reaction is

produced, a salt is formed, consisting of sulphuric anhydride

and potassium oxide. In this, as in the reaction of hydrates,

an intermediate oxide is separated out--namely, alumina. Its

separation will be very evident in this case, as alumina is

insoluble in water.

[52] The mutual interaction of hydrates, and their capacity of forming

salts, may be taken advantage of for determining the character

of those hydrates which are insoluble in water. Let us imagine

that a given hydrate, whose chemical character is unknown, is

insoluble in water. It is therefore impossible to test its

reaction on litmus. It is then mixed with water, and an acid--for

instance, sulphuric acid--is added to the mixture. If the hydrate

taken be basic, reaction will take place, either directly or by

the aid of heat, with the formation of a salt. In certain cases,

the resultant salt is soluble in water, and this will at once

show that combination has taken place between the insoluble basic

hydrate and the acid, with the formation of a soluble saline

substance. In those cases where the resultant salt is insoluble,

still the water loses its acid reaction, and therefore it may he

ascertained, by the addition of an acid, whether a given hydrate

has a basic character, like the hydrates of oxide of copper,

lead, &c. If the acid does not act on the given insoluble hydrate

(at any temperature), then it has not a basic character, and it

should be tested as to whether it has an acid character. This is

done by taking an alkali, instead of the acid, and by observing

whether the unknown hydrate then dissolves, or whether the

alkaline reaction disappears. Thus it may he proved that hydrate

of silica is acid, because it dissolves in alkalis and not in

acids. If it be a case of an insoluble intermediate hydrate, then

it will be observed to react on both the acid and alkali. Hydrate

of alumina is an instance in question, which is soluble both in

caustic potash and in sulphuric acid.

The _degree of affinity_ or chemical _energy_ proper to oxides

and their hydrates is very dissimilar; some extreme members

of the series possess it to a great extent. When acting on

each other they evolve a large quantity of heat, and when

acting on intermediate hydrates they also evolve heat to a

considerable degree, as we saw in the combination of lime and

sulphuric anhydride with water. When extreme oxides combine

they form stable salts, which are decomposed with difficulty,

and often show characteristic properties. The compounds of the

intermediate oxides with each other, or even with basic and acid

oxides, present a very different case. However much alumina we

may dissolve in sulphuric acid, we cannot saturate the acid

properties of the sulphuric acid, the resulting solution will

always have an acid reaction. So also, whatever quantity of

alumina is dissolved in an alkali, the resulting solution will

always present an alkaline reaction.

[53] In order to give an idea of the quantity of heat evolved in the

formation of salts I append a table of data for _very dilute

aqueous solutions_ of acids and alkalis, according to the

determinations of Berthelot and Thomsen. The figures are given

in major calories--that is, in thousands of units of heat. For

example, 49 grams of sulphuric acid, H_{2}SO_{4}, taken in a

dilute aqueous solution, when mixed with such an amount of a weak

solution of caustic soda, NaHO, that a neutral salt is formed

(when all the hydrogen of the acid is replaced by the sodium),

evolves 15,800 units of heat.

49 parts of 63 parts of

H_{2}SO_{4} HNO_{3}

NaHO 15·8 13·7

KHO 15·7 13·8

NH_{3} 14·5 12·5

CaO 15·6 13·9

BaO 18·4 13·9

MgO 15·6 13·8

FeO 12·5 10·7 (?)

ZnO 11·7 9·8

Fe_{2}O_{3} 5·7 5·9

These figures cannot be considered as the heat of neutralisation,

because the water here plays an important part. Thus, for

instance, sulphuric acid and caustic soda in dissolving in water

evolve very much heat, and the resultant sodium sulphate very

little; consequently, the amount of heat evolved in an anhydrous

combination will be different from that evolved in a hydrated

combination. Those acids which are not energetic in combining

with the same quantity of alkalis required for the formation of

normal salts of sulphuric or nitric acids always, however, give

less heat. For instance, with caustic soda: carbonic acid gives

10·2, hydrocyanic, 2·9, hydrogen sulphide, 3·9 major calories.

And as feeble bases (for example, Fe_{2}O_{3}) also evolve less

heat than those which are more powerful, so a certain general

correlation between thermochemical data and the degree of

affinity shows itself here, as in other cases (_see_ Chapter II.,

Note 7); this does not, however, give any reason for measuring

the affinity which binds the elements of salts by the heat

of their formation in dilute solutions. This is very clearly

demonstrated by the fact that water is able to decompose many

salts, and is separated in their formation.

[54] Carbonic anhydride evolves heat in dissolving in water. The

solution easily dissociates and evolves carbonic anhydride,

according to the law of Henry and Dalton (_see_ Chapter I.)

In dissolving in caustic soda, it either gives a normal salt,

Na_{2}CO_{3}, which does not evolve carbonic anhydride, or an

acid salt, NaHCO_{3} which easily evolves carbonic anhydride when

heated. The same gas, when dissolved in solutions of salts, acts

in one or the other manner (_see_ Chapter II., Note 38). Here

it is seen what a successive series of relations exists between

compounds of a different order, between substances of different

degrees of stability. By making a distinction between the

phenomena of solutions and chemical compounds, we overlook those

natural transitions which in reality exist.

Judging from the above, a salt is a compound of basic and acid oxides,

or the result of the action of hydrates of these classes on each other

with separation of water. But salts may be obtained by other methods. It

must not be forgotten that basic oxides are formed by metals, and acid

oxides usually by non-metals. But metals and non-metals are capable of

combining together, and a salt is frequently formed by the oxidation of

such a compound. For example, iron very easily combines with sulphur,

forming iron sulphide FeS (as we saw in the Introduction); this in

air, and especially moist air, absorbs oxygen, with the formation of

the same salt FeSO_{4}, that may be obtained by the combination of the

oxides of iron and sulphur, or of the hydrates of these oxides. Hence,

it cannot be said or supposed that a salt has the properties of the

oxides, or must necessarily contain two kinds of oxides in itself.

The derivation of salts from oxides is merely one of the methods of

their preparation. We saw, for instance, that in sulphuric acid it was

possible to replace the hydrogen by zinc, and that by this means zinc

sulphate was formed; so likewise the hydrogen in many other acids may be

replaced by zinc, iron, potassium, sodium, and a whole series of similar

metals, corresponding salts being obtained. The hydrogen of the acid,

in all these cases, is exchanged for a metal, and a salt is obtained

from the hydrate. Regarding a salt from this point of view, it may be

said that _a salt is an acid in which hydrogen is replaced by a metal_.

This definition shows that a salt and an acid are essentially compounds

of the same series, with the difference that the latter contains

hydrogen and the former a metal. Such a definition is more exact than

the first definition of salts, inasmuch as it likewise includes those

acids which do not contain oxygen, and, as we shall afterwards learn,

there is a series of such acids. Such elements as chlorine and bromine

form compounds with hydrogen in which the hydrogen may be replaced by

a metal, forming substances which, in their reactions and external

characters, resemble the salts formed from oxides. Table salt, NaCl, is

an example of this. It may be obtained by the replacement of hydrogen

in hydrochloric acid, HCl, by the metal sodium, just as sulphate of

sodium, Na_{2}SO_{4}, may be obtained by the replacement of hydrogen in

sulphuric acid, H_{2}SO_{4}, by sodium. The exterior appearance of the

resulting products, their neutral reaction, and even their saline taste,

show their resemblance to one another.

To the fundamental properties of salts yet another must be

added--namely, that they are more or less _decomposed by the action of a

galvanic current_. The results of this decomposition are very different

according to whether the salt be taken in a fused or dissolved state.

But the decomposition may generally be so represented, that the metal

appears at the electro-negative pole or cathode (like hydrogen in the

decomposition of water, or its mixture with sulphuric acid), and the

remaining parts of the salt appear at the electro-positive pole or anode

(where the oxygen of water appears). If, for instance, an electric

current acts on an aqueous solution of sodium sulphate, then the sodium

appears at the negative pole, and oxygen and the anhydride of sulphuric

acid at the positive pole. But in the solution itself the result is

different, for sodium, as we know, decomposes water with evolution of

hydrogen, forming caustic soda; consequently hydrogen will be evolved,

and caustic soda appear at the negative pole: while at the positive

pole the sulphuric anhydride immediately combines with water and forms

sulphuric acid, and therefore oxygen will be evolved and sulphuric acid

formed round this pole.[55] In other cases, when the metal separated is

not able to decompose water, it will be deposited in a free state. Thus,

for example, in the decomposition of copper sulphate, copper separates

out at the cathode, and oxygen and sulphuric acid appear at the anode,

and if a copper plate be attached to the positive pole, then the oxygen

evolved will oxidise the copper, and the oxide of copper will dissolve

and be deposited at the negative pole--that is, a transfer of copper

from the positive to the negative pole ensues. The galvanoplastic art

(electro-typing) is based on this principle.[56] Therefore the most

radical and general properties of salts (including also such salts as

table salt, which contain no oxygen) may be expressed by representing

the salt as composed of a metal M and a haloid X--that is, by expressing

the salt by MX. In common table salt the metal is sodium, and the haloid

an elementary body, chlorine. In sodium sulphate, Na_{2}SO_{4}, sodium

is again the metal, but the complex group, SO_{4}, is the haloid. In

sulphate of copper, CuSO_{4}, the metal is copper and the haloid the

same as in the preceding salt. Such a representation of salts expresses

with great simplicity the _capacity of every salt to enter into saline

double decompositions with other salts_; consisting in the mutual

replacement of the metals in the salts. This exchange of their metals

is the fundamental property of salts. In the case of two salts with

different metals and haloids, which are in solution or fusion, or in

any other manner brought into contact, the metals of these salts will

always partially or wholly exchange places. If we designate one salt

by MX, and the other by NY, then we either partially or wholly obtain

from them new salts, MY and NX. Thus we saw in the Introduction, that

on mixing solutions of table salt, NaCl, and silver nitrate, AgNO_{3},

a white insoluble precipitate of silver chloride, AgCl, is formed and

a new salt, sodium nitrate, NaNO_{3}, is obtained in solution. If the

metals of salts exchange places in reactions of double decomposition,

it is clear that metals themselves, taken in a separate state, are

able to act on salts, as zinc evolves hydrogen from acids, and as iron

separates copper from copper sulphate. When, to what extent, and which

metals displace each other, and how the metals are distributed between

the haloids, will be discussed in Chapter X., where we shall be guided

by those reflections and deductions which Berthollet introduced into the

science at the beginning of this century.

[55] This kind of decomposition may be easily observed by pouring

a solution of sodium sulphate into a U-shaped tube and

inserting electrodes in the two branches. If the solution be

coloured with an infusion of litmus, it will easily be seen

that it turns blue at the cathode, owing to the formation of

sodium hydroxide, and red at the electro-positive pole, from the

formation of sulphuric acid.

[56] In other cases the decomposition of salts by the electric current

may be accompanied by much more complex results. Thus, when the

metal of the salt is capable of a higher degree of oxidation,

such a higher oxide may be formed at the positive pole by the

oxygen which is evolved there. This takes place, for instance,

in the decomposition of salts of silver and manganese by the

galvanic current, peroxides of these metals being formed. Thus in

the electrolysis of a solution of KCl, KClO_{3} is formed, and

of sulphuric acid (corresponding to SO_{3}) persulphuric acid,

corresponding to S_{2}O_{7}. But all the phenomena as yet known

may be expressed by the above law--that the current decomposes

salts into metals, which appear at the negative pole, and into

the remaining component parts, which appear at the positive pole.

According to the above observations, an acid is nothing more than a

salt of hydrogen. Water itself may be looked on as a salt in which

the hydrogen is combined with either oxygen or the aqueous radicle,

OH; water will then be HOH, and alkalis or basic hydrates, MOH. The

group OH, or the _aqueous radicle_, otherwise called _hydroxyl_, may

be looked on as a haloid like the chlorine in table salt, not only

because the element Cl and the group OH very often change places, and

combine with one and the same element, but also because free chlorine is

very similar in many properties and reactions to peroxide of hydrogen,

which is the same in composition as the aqueous radicle, as we shall

afterwards see in Chapter IV. Alkalis and basic hydrates are also

salts consisting of a metal and hydroxyl--for instance, caustic soda,

NaOH; this is therefore termed _sodium hydroxide_. According to this

view, _acid salts_ are those in which a portion only of the hydrogen is

replaced by a metal, and a portion of the hydrogen of the acid remains.

Thus sulphuric acid (H_{2}SO_{4}) not only gives the normal salt

Na_{2}SO_{4}, with sodium, but also an acid salt, NaHSO_{4}. A _basic

salt_ is one in which the metal is combined not only with the haloids

of acids, but also with the aqueous radicale of basic hydrates--for

example, bismuth gives not only a normal salt of nitric acid,

Bi(NO_{3})_{3}, but also basic salts like Bi(OH)_{2}(NO_{3}).

As basic and acid salts of the oxygen acids contain hydrogen and oxygen,

they are able to part with these as water and to give anhydro-salts,

which it is evident will be compounds of normal salts with anhydrides

of the acids or with bases. Thus the above-mentioned acid sodium

sulphate corresponds with the anhydro-salt, Na_{2}S_{2}O_{7}, equal

to 2NaHSO_{4}, less H_{2}O. The loss of water is here, and frequently

in other cases, brought about by heat alone, and therefore such salts

are frequently termed _pyro-salts_--for instance, the preceding is

sodium pyrosulphate (Na_{2}S_{2}O_{7}), or it may be regarded as the

normal salt Na_{2}SO_{4} + sulphuric anhydride, SO_{3}. _Double_ salts

are those which contain either two metals, KAl(SO_{4})_{2}, or two

haloids.[57]

[57] The above-enunciated generalisation of the conception of salts as

compounds of the metals (simple, or compound like ammonium,

NH_{4}), with the haloids (simple, like chlorine, or compound,

like cyanogen, CN, or the radical of sulphuric acid, SO_{4}),

capable of entering into double saline decomposition, which

is in accordance with the general data respecting salts, was

only formed little by little after a succession of most varied

propositions as to the chemical structure of salts.

Salts belong to the class of substances which have been known

since very early times, and have long been investigated in many

directions. At first, however, no distinction was made between

salts, acids, and bases. Glauber prepared many artificial salts

during the latter half of the seventeenth century. Up to that

time the majority of salts were obtained from natural sources,

and that salt which we have referred to several times--namely,

sodium sulphate--was named Glauber's salt after this chemist.

Rouelle distinguished normal, acid, and basic salts, and showed

their action on vegetable dyes, still he confounded many salts

with acids (even now every acid salt ought to be regarded as an

acid, because it contains hydrogen, which may be replaced by

metals--that is, it is the hydrogen of an acid). Baumé disputed

Rouelle's opinion concerning the subdivision of salts, contending

that normal salts only are true salts, and that basic salts are

simple mixtures of normal salts with bases and acid salts with

acids, considering that washing alone could remove the base or

acid from them. Rouelle, in the middle of the last century,

however, rendered a great service to the study of salts and the

diffusion of knowledge respecting this class of compounds in

his attractive lectures. He, like the majority of the chemists

of that period, did not employ the balance in his researches,

but satisfied himself with purely qualitative data. The first

quantitative researches on salts were carried on about this

time by Wenzel, who was the director of the Freiburg mines, in

Saxony. Wenzel studied the double decomposition of salts, and

observed that in the double decomposition of neutral salts a

neutral salt was always obtained. He proved, by a method of

weighing, that this is due to the fact that the saturation of a

given quantity of a base requires such relative quantities of

different acids as are capable of saturating every other base.

Having taken two neutral salts--for example, sodium sulphate and

calcium nitrate--let us mix their solutions together. Double

decomposition takes place, because calcium sulphate is formed,

which is almost insoluble. However much we might add of each

of the salts, the neutral reaction will still be preserved,

consequently the neutral character of the salts is not destroyed

by the interchange of metals; that is to say, that quantity of

sulphuric acid which saturated the sodium is sufficient for the

saturation of the calcium, and that amount of nitric acid which

saturated the calcium is enough to saturate the sodium contained

in combination with sulphuric acid in sodium sulphate. Wenzel

was even convinced that matter does not disappear in nature, and

on this principle he corrects, in his _Doctrine of Affinity_,

the results of his experiments when he found that he obtained

less than he had originally taken. Although Wenzel deduced the

law of the double decomposition of salts quite correctly, he did

not determine those quantities in which acids and bases act on

each other. This was carried out at the end of the last century

by Richter. He determined the quantities by weight of the bases

which saturate acids and of the acids which saturate bases, and

obtained comparatively correct results, although his conclusions

were not correct, for he states that the quantity of a base

saturating a given acid varies in arithmetical progression, and

the quantity of an acid saturating a given base in geometrical

progression. Richter studied the deposition of metals from their

salts by other metals, and observed that the neutral reaction

of the solution is not destroyed by this exchange. He also

determined the quantities by weight of the metals replacing one

another in salts. He showed that copper displaces silver from

its salt, and that zinc displaces copper and a whole series of

other metals. Those quantities of metals which were capable of

replacing one another were termed equivalents.

Richter's teaching found no followers, because, although he

fully believed in the discoveries of Lavoisier, yet he still

held to the phlogistic reasonings which rendered his expositions

very obscure. The works of the Swedish savant Berzelius freed

the facts discovered by Wenzel and Richter from the obscurity

of former conceptions, and led to their being explained in

accordance with Lavoisier's views, and in the sense of the law of

multiple proportions which had already been discovered by Dalton.

On applying to salts those conclusions which Berzelius arrived at

by a whole series of researches of remarkable accuracy, we arrive

at the following law of equivalents--_one part by weight of

hydrogen in an acid is replaced by the corresponding equivalent

weight of any metal_; and, therefore, when metals replace each

other their weights are in the same ratio as their equivalents.

Thus, for instance, one part by weight of hydrogen is replaced by

23 parts of sodium, 39 parts of potassium, 12 parts of magnesium,

20 parts of calcium, 28 parts of iron, 108 parts of silver, 33

parts of zinc, &c.; and, therefore, if zinc replaces silver, then

33 parts of zinc will take the place of 108 parts of silver, or

33 parts of zinc will he substituted by 23 parts of sodium, &c.

The doctrine of equivalents would be precise and simple did

every metal only give one oxide or one salt. It is rendered

complicated from the fact that many metals form several oxides,

and consequently offer different equivalents in their different

degrees of oxidation. For example, there are oxides containing

iron in which its equivalent is 28--this is in the salts formed

by the suboxide; and there is another series of salts in which

the equivalent of iron equals 18-2/3--which contain less iron,

and consequently more oxygen, and correspond with a higher degree

of oxidation--ferric oxide. It is true that the former salts are

easily formed by the direct action of metallic iron on acids, and

the latter only by a further oxidation of the compound formed

already; but this is not always so. In the case of copper,

mercury, and tin, under different circumstances, salts are formed

which correspond with different degrees of oxidation of these

metals, and many metals have two equivalents in their different

salts--that is, in salts corresponding with the different degrees

of oxidation. Thus it is impossible to endow every metal with one

definite equivalent weight. Hence the conception of equivalents,

while playing an important part from an historical point of view,

appears, with a fuller study of chemistry, to be but subordinate

to a higher conception, with which we shall afterwards become

acquainted.

The fate of the theoretical views of chemistry was for a

long time bound up with the history of salts. The clearest

representation of this subject dates back to Lavoisier, and was

systematically developed by Berzelius. This representation is

called the _binary_ theory. All compounds, and especially salts,

are represented as consisting of two parts. Salts are represented

as compounds of a basic oxide (a base) and an acid (that is, an

anhydride of an acid, then termed an acid), whilst hydrates are

represented as compounds of anhydrous oxides with water. Such

an expression was employed not only to denote the most usual

method of formation of these substances (where it would be quite

true), but also to express that internal distribution of the

elements by which it was proposed to explain all the properties

of these substances. Copper sulphate was supposed to contain

two most intimate component parts--copper oxide and sulphuric

anhydride. This is an hypothesis. It arose from the so-called

_electro-chemical hypothesis_, which supposed the two component

parts to be held in mutual union, because one component (the

anhydride of the acid) has electro-negative properties, and the

other (the base in salts) electro-positive. The two parts are

attracted together, like substances having opposite electrical

charges. But as the decomposition of salts in a state of fusion

by an electric current always gives a metal, that representation

of the constitution and decomposition of salts called the

_hydrogen theory_ of acids is nearer the truth than that which

considers salts as made up of a base and an anhydride of an acid.

But the hydrogen theory of acids is also a binary hypothesis,

and does not contradict the electro-chemical hypothesis, but

is rather a modification of it. The binary theory dates from

Rouelle and Lavoisier, the electro-chemical aspect was zealously

developed by Berzelius, and the hydrogen theory of acids is due

to Davy and Liebig.

These hypothetical views simplified and generalised the study of

a complicated subject, and served to support further arguments,

but when salts were in question it was equally convenient to

follow one or the other of these hypotheses. But these theories

were brought to bear on all other substances, on all compound

substances. Those holding the binary and electro-chemical

hypotheses searched for two anti-polar component parts, and

endeavoured to express the process of chemical reactions by

electro-chemical and similar differences. If zinc replaces

hydrogen, they concluded that it is more electro-positive

than hydrogen, whilst they forgot that hydrogen may, under

different circumstances, displace zinc--for instance, at a

red heat. Chlorine and oxygen were considered as being of

opposite polarity to hydrogen because they easily combine with

it, nevertheless both are capable of replacing hydrogen, and,

what is very characteristic, in the replacement of hydrogen

by chlorine in carbon compounds not only does the chemical

character often remain unaltered, but even the external form

may remain unchanged, as Laurent and Dumas demonstrated. These

considerations undermine the binary, and more especially the

electro-chemical theory. An explanation of known reactions then

began to be sought for not in the difference of the polarity of

the different substances, but in the joint influences of all the

elements on the properties of the compound formed. This is the

reverse of the preceding hypothesis.

This reversal was not, however, limited to the destruction of

the tottering foundations of the preceding theory; it proposed a

new doctrine, and laid the foundation for the modern course of

our science. This doctrine may be termed the unitary theory--that

is, it strictly acknowledges the joint influences of the elements

in a compound substance, denies the existence of separate and

contrary components in them, regards copper sulphate, for

instance, as a strictly definite compound of copper, sulphur,

and oxygen; then seeks for compounds which are analogous in

their properties, and, placing them side by side, endeavours

to express the influence of each element in determining the

united properties of its compound. In the majority of cases it

arrives at conclusions similar to those which are obtained by

the above-mentioned hypotheses, but in certain special cases the

conclusions of the unitary theory are in entire opposition to

those of the binary theory and its corollaries. Cases of this

kind are most often met with in the consideration of compounds of

a more complex nature than salts, especially organic compounds

containing hydrogen. But it is not in this change from an

artificial to a natural system, important as it is, that the

chief service and strength of the unitary doctrine lies. By a

simple review of the vast store of data regarding the reactions

of typical substances, it succeeded from its first appearance

in establishing a new and important law, it introduced a new

conception into science--namely, the conception of molecules,

with which we shall soon become acquainted. The deduction of

the law and of the conception of molecules has been verified by

facts in a number of cases, and was the cause of the majority of

chemists of our times deserting the binary theory and accepting

the unitary theory, which forms the basis of the present work.

Laurent and Gerhardt must be considered as the founders of this

doctrine.

Inasmuch as oxygen compounds predominate in nature, it should be

expected from what has been said above, that salts, rather than acids

or bases, would occur most frequently in nature, for these latter would

always tend to combine forming salts, especially through the medium

of the all-pervading water. And, as a matter of fact, salts are found

everywhere in nature. They occur in animals and plants, although in

but small quantity, because, as forming the last stage of chemical

reaction, they are capable of only a few chemical transformations. And

organisms are bodies in which a series of uninterrupted, varied, and

active chemical transformations proceed, whilst salts, which only enter

into double decompositions between each other, are little prone to such

changes. But organisms always contain salts. Thus, for instance, bones

contain calcium phosphate, the juice of grapes potassium tartrate (cream

of tartar), certain lichens calcium oxalate, and the shells of mollusca

calcium carbonate, &c. As regards water and soil, portions of the earth

in which the chemical processes are less active, they are full of

salts. Thus the waters of the oceans, and all others (Chap. I.), abound

in salts, and in the soil, in the rocks of the earth's crust, in the

upheaved lavas, and in the falling meteorites the salts of silicic acid,

and especially its double salts, predominate. Saline substances also

make up the composition of those limestones which often form mountain

chains and whole thicknesses of the earth's strata, these consisting of

calcium carbonate, CaCO_{3}.

Thus we have seen oxygen in a free state and in various compounds

of different degrees of stability, from the unstable salts, like

Berthollet's salt and nitre, to the most stable silicon compounds, such

as exist in granite. We saw an entirely similar gradation of stability

in the compounds of water and of hydrogen. In all its aspects oxygen,

as an element, or single substance, remains the same however varied

its chemical states, just as a substance may appear in many different

physical states of aggregation. But our notion of the immense variety of

the chemical states in which oxygen can occur would not be completely

understood if we did not make ourselves acquainted with it in the form

in which it occurs in ozone and peroxide of hydrogen. In these it is

most active, its energy seems to have increased. They illustrate fresh

aspects of chemical correlations, and the variety of the forms in which

matter can appear stand out clearly. We will therefore consider these

two substances somewhat in detail.