THE TEMPLE PRIMERS
MODERN CHEMISTRY Systematic
By WILLIAM RAMSAY, D.Sc.
JOHN DALTON
mODERH
CHEmiSTRY
SYSTEI12ATN
LAlttSAY'DS?
1900* ±9 &5O BEDFORD-STREET*
All rights reserved
MODERN CHEMISTRY
SECOND PART SYSTEMATIC CHEMISTRY
CHAPTER I
Methods of Preparing Elements — Their Physical Properties.
Mixtures and Compounds. — In the olden days, no distinction was drawn between a compound and a mixture. Indeed, all " impure " substances artificially prepared were termed " mixts." It was only after the true idea of ele- ments had been arrived at, and indeed not until Dalton had formulated the laws which go by his name, that the distinc- tion was drawn. The ultimate criterion for combination is definiteness of proportion, and this is generally connected with uniformity in properties, or homogeneity. A sub- stance is said to be homogeneous when no one part of it differs from any other part in composition. But this may be predicated of glass, or of air, which are mixtures, and not compounds. A mixture may be homogeneous ; a com- pound must.
Again, it is usually accepted that the separation of the constituents of a mixture may be effected by mechanical, or at least by physical means ; whereas the separation of the elements from a compound require chemical treatment. Here it is difficult to draw a sharp distinction. The
VOL. II. A
385213
2 MODERN CHEMISTRY
separation cf carbon dioxide from soda-water by the appli- . cation of heat is similar in character to the separation of sugar from water by evaporation of the water ; yet we believe that a solution of carbon dioxide in water consti- tutes a compound, while that of sugar in water is a mere mixture of the two. It is necessary to be guided by analogy in the former case ; and it is probable that the compound named carbonic acid is really contained in a solution of carbon dioxide in water, on account of the formulas and behaviour of the carbonates.
The Atmosphere. — In the case of mixtures of gases, the problem becomes an easier one. For in this case, each gas retains its individual properties. The atmosphere, for example, is believed to be a mixture of the gases
Nitrogen, . . 78.16 per cent. Oxygen, . 20.90 „
Argon, &c., . . 0.94 „
100.00
if small amounts of water-vapour, of carbon dioxide, and of ammonia, all of which vary considerably in amount, be subtracted.
This can be shown by several lines of argument.
First, The density of air agrees with the mean of the densities of its constituents, taken in the proportion in which they occur. Thus, the density of the mixture of atmos- pheric nitrogen and argon differs by only I part in 40,000 from that calculated from their relative weights, and the proportion in which they occur. This is the case with compound gases only when the constituents are present in equal proportions by volume, as in hydrogen chloride, HC1. The above mixture is far from fulfilling that requirement.
Second, The constituents of air can be separated by diffusion. Thomas Graham discovered that the rate of escape of gases through an opening, or of passage through
THE ATMOSPHERE 3
a porous partition is inversely in the order of the square roots of their relative densities. Now, air has been enriched in oxygen and in argon by diffusion ; the lighter nitrogen
passes more rapidly in the proportion of
— /= : ,— : —7—, 0/14 \/i6 V20
the last two fractions referring to the rates of oxygen and argon respectively ; the oxygen and argon, being more slowly diffusible, are left to the last.
Third, The constituents of air may be separated by solution in water. While oxygen is soluble at atmospheric temperature in the proportion of about 3 volumes in 100 of water, nitrogen is much less soluble — about 1.5 volumes ; and argon about 4.1 volumes* Hence, on shaking air with water, the relative volumes dissolved are :
Oxygen, 3 x 20.90 ; Nitrogen, 1.5 x 78.16 ; and Argon, 4.1 x 0.9 4,
or in the proportion of 63 : 1 17 : 3.8. It is evident that the relative proportion of nitrogen has considerably decreased.
Fourth, The elements contained in air are not present in any atomic ratio. To ascertain the relative number of atoms of these elements it is necessary to divide the per- centage amount of each by its atomic weight ; thus we have
Nitrogen, Z_: _ = 5.58 ; Oxygen, ^9 = 1.31 ;
Argon, 2^51 = 0.024 ; 40
and these numbers bear to each other no simple ratio.
Lastly, it is possible by distilling liquid air to separate the more volatile nitrogen from the less volatile oxygen and argon.
For these reasons, and other similar ones, it is concluded that air is a mixture.
4 MODERN CHEMISTRY
The Analysis of the Atmosphere is, however, always performed by chemical means, for the difference in physical properties of its constituents is not sufficiently marked to allow of their being utilised for purposes of separation. Many common elements unite easily with oxygen to form non-volatile compounds, when they are heated in air. One of the most convenient for this purpose is metallic copper. By passing a known volume of air over copper turnings, contained in a counter- poised tube of hard glass, and heated to redness, the oxygen of the air is removed, for it combines with the copper to form non- volatile black oxide of copper. The in- crease in weight of this tube gives the weight of the oxygen in the measured volume of air. But it is customary to analyse air volumetrically by absorbing the oxygen from a known volume by means of burning phosphorus, or of a solution of potassium pyrogallate : the remainder consists of a mixture of nitrogen, argon and its congeners. The separation of these gases from each other is described in the next paragraph.
Reference has already been made in Part I. to the different processes which may be used for the isolation of elements from their compounds. But there exists a group of elements, that of which the first member is helium, which form no com- pounds, and which therefore are found only in a free state. It is, therefore, convenient to begin with these.
The HELIUM Group. — These elements are all gases at the ordinary temperature of the atmosphere, and they are consequently all to be found in atmospheric air. They are colourless, even in the liquid condition, and are devoid of smell and taste. They are very sparingly soluble in water ; for example, 100 volumes of water dissolve only 4.1 volumes of argon at 1 5°. Their preparation consists, first, in the separation of the other constituents of air from them, and, second, in their separation from each other.
Air, which is a mixture, and not a compound, of nitro- gen, oxygen, carbon dioxide, ammonia, water-vapour, and the gases of the helium group, is a supporter of combustion,
THE HELIUM GROUP 5
owing to the combination of the oxygen which it contains with most other elements. Now, when air passed through a tube full of a mixture of caustic soda and lime, to remove carbon dioxide, and then through a U-tube containing sul- phuric acid, to deprive it of water-vapour and ammonia, is led over red-hot copper, or over some other red-hot metal which unites with oxygen, the oxygen is retained, and nitrogen with members of the helium group alone passes on. The nitrogen can be removed in one of two ways. The first plan is due to Cavendish, who attempted to prove that atmospheric nitro- gen was a homogeneous substance. He mixed atmospheric nitrogen with oxygen, and passed electric sparks through the mixture, having a little caustic soda present in the tube. Under the influence of the sparks, the nitrogen and oxygen combine, giving nitride peroxide, NO9 ; this com- pound is absorbed by the soda, with formation of sodium nitrate and nitrite, NaNO3 and NaNO2. Cavendish obtained a residue of not more than one-hundred-and- twentieth of the nitrogen ; and he concluded that if atmospheric nitrogen was not homogeneous, it contained only a trace of another gas. The second plan is to pass the atmospheric nitrogen over red-hot magnesium, or, better, over a mixture of magnesium powder and lime, which gives calcium ; the magnesium or the calcium unites with the nitrogen, and the inert gases pass on.
To separate these gases from each other, they are compressed into a bulb, cooled to -185° by being immersed in liquid air. The argon, krypton, and xenon condense to a liquid, in which the neon and helium are dissolved. On removing the bulb from the liquid air, its temperature rises, and the helium and neon escape first, mixed with a large amount of argon. Argon distils next, and krypton and xenon remain till the last. By frequently repeating this process of " fractional distillation," the argon, krypton, and xenon can be separated from each other, and from the helium and neon which still remain mixed with each other, for both are gases at the temperature of boiling air.
6 MODERN CHEMISTRY
To separate helium from neon, recourse must be had to liquid hydrogen. To liquefy hydrogen, the process is in principle the same as that for liquefying air, described on p. 26. The hydrogen, compressed by a pressure of 200 atmospheres, is cooled to —205° by passing through a coil of copper pipe, immersed in liquid air boiling under low pressure. On expanding, its temperature is still further lowered, and the still colder gas, in passing upwards, cools the tubes through which the compressed gas is passing. The hydrogen finally issues in the liquid state, as a colour- less, mobile liquid, of the approximate temperature —240°. By its aid, if a mixture of neon and helium is cooled to —240°, the former freezes, while the latter remains gaseous. The gaseous helium can be removed with the pump ; and the neon, after it has been warmed, may also be pumped off in a pure state.
Helium can also be prepared by heating certain specimens of pitchblende or uraninite, a mineral consisting chiefly of oxide of uranium. The gas, which appears to exist in some sort of combination with the uranium oxide, escapes ; it contains a trace of argon. All these gases give very striking spectra, and that of helium was observed during the solar eclipse of 1 868 in the chromosphere, or coloured atmosphere, of the sun. Although at that time it had not been dis- covered on the earth, the name " helium " was given to the bright yellow line, which is the most characteristic of its spectrum.
As regards the relative amount of these gases contained in air, 100 volumes of air contain 0.937 volume of the mixture. By far the largest portion of this mixture is argon ; probably the volume of all the others taken together does not exceed one-four-hundredth part of that of the argon. Indeed, it may be said with truth that there is less xenon in air than there is gold in sea-water.
Methods of Separating Elements from their Compounds* — The methods of preparation of the remain- ing elements depend on considerations of the cost of the
SEPARATION PROCESSES 7
compound from which the element is to be prepared, and on the ease of preparation. In the case of those elements which are required on a commercial scale, like iron, for example, the process of manufacture is regulated chiefly by the cost of the ore, and of the operations necessary to produce the metal in a state of purity sufficient for commercial purposes. But if perfectly pure iron is required for scientific purposes — for example, in order to determine its electrical properties — then the question of cost does not come into consideration, and processes are adopted which are necessarily very costly. In the description which follows, however, we shall give only the ordinary methods of preparation.
Again, the process chosen depends greatly on the physical and chemical properties of the element which it is desired to isolate. Some elements are volatile, and are more or less easily separated by distillation from the material from which they are produced ; some elements are attacked by water, while others resist attack ; some fuse at comparatively low temperatures, and can thus be separated, while others are producible in a compact state only at the enormously high temperature of the electric arc. It is necessary, therefore, to know the properties of the element required before deciding on a process for its isolation. The preparation of the remaining elements will therefore be considered from this point of view.
( i ) Separation of the element by means of an electric current.
(a) From a fused salt. — One condition is that the salt shall fuse at a convenient temperature — that is, at or below a red heat. Another is that, in the case of metals which are commercially used, the salts must be cheaply obtainable, and the metals easily separated from the salts.
It is interesting to note that this process led, in the hands of Sir Humphry Davy, to the discovery of the metals of the alkalies, potassium and sodium ; he first prepared them by passing a current from a battery of high voltage
8 MODERN CHEMISTRY
through the hydroxide, melted on a piece of platinum foil. The metal was visible only for an instant ; for it floated up from the electrode of platinum wire, and burst into flame as soon as it came into the air.
As a rule, however, the chlorides are the most con- venient salts for electrolysis. From the known fact that the melting-point of a compound is lowered by the presence of an "impurity," it is often found advantageous to electro- lyse a mixture of chlorides rather than a pure chloride ; in this case one of the elements is liberated in preference to the other. As the anode has to withstand the action of chlorine, it is always made of carbon, which does not unite with chlorine directly ; the kathode may be of iron, a metal which has no tendency to form alloys with those which are prepared in this way, at least at the temperatures required. The kathode may be the iron pot in which the chloride is kept fused.
The elements which are prepared in this way are : lithium, sodium, potassium, rubidium, caesium, beryllium, magnesium, calcium, strontium, and barium. The first five are easily fusible white soft metals, which take fire when heated in air, and must therefore be kept in an atmosphere free from oxygen ; they also attack water, liberating hydrogen, with formation of the hydroxide MOH. Their density is so low that they float on their fused chlorides ; they must, therefore, be liberated in the interior of a bell-shaped iron electrode or of a fireclay receptacle, down which an iron kathode passes. Beryllium and magnesium are better prepared from a mixture of their chlorides with potassium chloride ; the latter melts and collects at the bottom of the pot, which, in this case, may be the kathode. They are hard white metals, magnesium melting at about 750°, and beryllium about 1 200°. They, too, take fire when heated in air, and burn with a brilliant flame ; indeed, the chief use of magnesium is for signalling purposes. The metal is drawn, while hot, into wire, which is then rolled into ribbon; this ribbon burns with an exceedingly bright flame,
METALS OF THE ALKALIES 9
producing the oxide MgO. Calcium, strontium, and barium are also white metals ; they have been produced by electrolysis of their cyanides, M(CN)2, compounds which fuse at a lower temperature than the chlorides. They are very readily attacked by water, yielding the hydroxides M(OH)0. The only two of these metals which find commercial use are sodium and magnesium.
Aluminium, which is also manufactured on a large scale, is produced from its ore, bauxite, from which pure alumina, the oxide, is first prepared. The alumina is dissolved in fused cryolite, a fluoride of aluminium and sodium of the formula NagAlF6, deposits of which occur in Greenland. The aluminium sinks to the bottom of the crucible, and when a sufficient quantity accumulates it is tapped out. The "flux," as the cryolite is termed, is again melted, and a further quantity of alumina is dissolved in it. The metal is fairly hard, white, susceptible of a high polish, ductile and malleable. It is also very light (about two and a half times as heavy as water), and not easily oxidised in air at the ordinary temperature, nor is it attacked by water.
(b) From a dissolved salt. — Gallium, a tin- white, hard metal, very rare, contained in some zinc ores, is deposited from a solution of its hydroxide in caustic potash. Copper prepared, as will be seen below, in a crude state by displacement, is purified by electrolysis. It is of the utmost importance to employ pure copper for the conduction of electric currents ; for although copper is one of the best conductors, its resistance is enormously increased by the presence of a very small trace of impurity. To purify it, large rectangular blocks of crude copper are suspended close to thin sheets of pure copper in an acid bath of copper sulphate, CuSO4.Aq. The heavy block is made the anode and the thin sheet the cathode; the
sulphation, SO4, in discharging at the anode, dissolves copper from the thick block as sulphate; while the cuprion, + + Cu, in yielding up its charge at the kathode, deposits on
TO MODERN CHEMISTRY
the latter and increases its thickness. The impurities, arsenic, antimony, and iron, remain in solution, and a sludge is deposited containing silver and gold, besides traces of many other elements. Copper is a very malleable, ductile red metal, melting at 1330°.
Objects of iron are often " nickel-plated," or covered with a thin film of nickel, a white, hard metal which pre- serves its lustre in air, for it is not easily oxidisable. This is done by making the object to be coated with nickel the kathode and a bar of nickel the anode ; the liquid is a solution of oxalate of nickel and potassium. Iron objects are first coated with copper before nickelling. Silver and gold are best deposited from their double cyanides with potassium ; these salts are used because the deposit is harder and more uniform than if a halide be used. In thus coating objects, it is of importance that the current density, i.e. the ratio of the current to the area of the surface of the object to be coated, should be considered ; if this be too high, the metal will be deposited in a loose, flocculent condition.
As an illustration of the changes which take place during such electrolysis, the deposition of silver may be chosen. The compound employed is, as stated, the double cyanide (see p. 187) ; its formula is KAg(CN)2, and the ions
+ are K and Ag(CN)2. There are, however, at the same
+
time a few ions of Ag and CN. From the last, metallic silver is deposited on the kathode ; and as soon as its amount is reduced, a fresh quantity is formed by the
decomposition of the complex ion, Ag(CN)9. The formation and deposition of the silver ion goes on con- tinuously until all the silver required has been deposited. Similar changes take place during the electro-deposition of nickel and of gold.
Modern electrolytic processes for obtaining chlorine and caustic soda (NaOH) from salt result in the liberation of enormous quantities of hydrogen. The salt, dissolved in
ELECTRO-DEPOSITION 11
water, is placed in a tank divided into two compartments by a porous diaphragm ; the anode, which consists of carbon rods, dips into one, and the kathode, which may be formed of copper plates, in the other. The ions, of
+
course, are Na.Aq, and Cl.Aq. The chlorine is liber- ated at the anode, and the sodium at the kathode. But as soon as the sodion is discharged, it reacts with the water, forming caustic soda, thus : 2Na + 2HOH = aNaOH + H2. Hence the production of hydrogen. Bromine and iodine may be liberated in the same way as chlorine, the bromide or iodide of sodium or potassium being substituted for the chloride. As fluorine at once acts on water, liberating • oxygen in the form of ozone, O3, it cannot be produced from an aqueous solution of a fluoride ; but it has been found that liquid hydrogen fluoride has ionising power, so that on passing a current between poles of platinum-iridium (an alloy of metals which is less attacked by fluorine than any other conductor) through a solution of hydrogen-potas- sium fluoride, HKF, in pure liquid hydrogen fluoride, H^F^, at -30°, fluorine is evolved from the anode as a pale yellow gas with a strong characteristic smell, somewhat resembling that of the other halogens, chlorine, bromine, and iodine ; while hydrogen is evolved at the kathode, having been produced by the action of the potas- sium on the hydrogen fluoride. Fluorine boils at —195°, chlorine at -35°, bromine at 59°, and iodine, which is a solid at atmospheric temperature, melts at 114° and boils at 184°. The colours of these elements also show a gradation. Chlorine is greenish-yellow; bromine, red both as gas and liquid ; iodine is a blue-black solid and a violet gas. These three elements are somewhat soluble in water, and more so in a solution of their soluble salts. It has recently been found that another ionising agent than water may be used. Lithium chloride is soluble in pyridine, a compound of the formula C5H5N, and may be electro- deposited on a platinum kathode from such a solution.
12 MODERN CHEMISTRY
The metal is not attacked by pyridine ; the chlorine, however, is rapidly absorbed.
(2) Separation of an element from a compound "by rise of temperature.
This method is applied in practice only to the prepara- tion of oxygen, and of chlorine, bromine, and iodine ; but many other elements may be thus made, where the com- pound heated does not tend to re-form on cooling. These cases will be considered first.
Ordinary coal-gas consists chiefly of methane, CH4, ethylene, C2H4, carbon monoxide, CO, and hydrogen, the last amounting to nearly 50 per cent, of the volume of the gas. This hydrogen owes its origin, at least in part, to the decomposition of its compounds with carbon, by their coming into contact with the red-hot walls of the retort in which the coal is distilled. Carbon deposits in a dense black mass on the iron, and is removed from time to time with a chisel. Hydrogen escapes and mixes with the coal-gas. This form of carbon is used for the pencils for arc-lights, and for the anodes of Bunsen's and other forms of cells, and also for anodes in electro-chemical processes.
The compounds of hydrogen with nitrogen (ammonia, NH3), sulphur, selenium, and tellurium (sulphuretted, seleniuretted, or telluretted hydrogen, H9S, H2Se, H0Te), all of which are gases at the ordinary temperature, are de- composed if passed through a red-hot tube, giving hydrogen, which escapes along with nitrogen if ammonia be heated ; or a deposit of the sulphur, &c., in the cold part of the tube if one of the other gases mentioned be employed.
The oxides of the metals ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, gold, and mercury are decomposed at a red heat ; and the chlorides, bromides, iodides, and sulphides are also decomposed, except those of silver and mercury.
But none of these methods are practical plans of prepar- ing the elements. On the other hand, as already stated,
DECOMPOSITION BY HEATING 13
this method is generally used for the production of oxygen. This gas, although it had probably been obtained in an impure state by the older experimenters, was first pro- duced in approximate purity by Priestley and simul- taneously by Scheele in 1774. Priestley produced it by heating mercuric oxide, HgO, which decomposes thus : 2HgO = 2Hg + O2. And Lavoisier showed that it was possible to produce mercuric oxide by heating mercury to its boiling-point in a confined portion of air, and by sepa- rating and weighing the oxide, and subsequently heating it till it decomposed again, he proved that the oxygen had really been extracted from the air.
Certain oxides are not wholly decomposed into oxygen and element when heated, but leave an oxide containing less oxygen than that originally heated. Among these is black manganese dioxide, a mineral named pyrolusite ; 3MnO2 = Mn3O4 + O0. Lead dioxide undergoes a similar change: 2PbO9= zPbO + O9. The most important ap- plication of this method, however, is the commercial plan of producing oxygen carried out in the " Brin Company's " works. In their process, barium oxide, BaO, is heated in iron tubes under pressure, air being pumped in. The barium oxide absorbs the oxygen of the air, the nitrogen being allowed to escape. After the operation has gone on for about five minutes, a considerable amount of oxygen is absorbed, barium dioxide, BaO2, being formed. The stopcocks of the pipes leading to the pump are then reversed, so that gas is exhausted from the hot iron tubes. When the pressure is reduced, the barium dioxide loses oxygen, and again returns to the state of monoxide : 2BaO2 = 2 BaO + O2. The pumping is continued for about five minutes, and the valves are again reversed. The pro- cess is thus a continuous one ; the oxygen is not pure, for it contains about 7 per cent, of nitrogen ; but for medical use in cases of pneumonia, and for the oxy-hydrogen blow- pipe, its purity is sufficient.
This method of preparing oxygen is an instance of what
I4 MODERN CHEMISTRY
is termed " mass-action." The temperature is kept con- stant, but the pressure is raised when it is desired to cause the oxide to absorb oxygen, and lowered when it is neces- sary to remove the oxygen. When pressure is raised, the number of molecules of oxygen in unit volume of the space ifor the mass) is increased, and hence the number in contact with the absorbing medium, the barium oxide. Combina- tion, therefore, takes place between the two. On reducing pressure, the number per unit volume is reduced, and the compound decomposes. The phenomenon is analogous with the behaviour of a vapour when it is compressed ; after a certain pressure has been reached — the vapour pressure — the vapour condenses to a liquid, and if more vapour be compressed into the same space, the pressure does not rise further, but more vapour is condensed : this is analogous to the formation of more BaO2. On pumping out vapour, the pressure does not fall, but the liquid evaporates : this is the analogue of the decomposition of the BaO2 into BaO. The law of mass-action is very generally applicable.
Certain oxides, for instance, pentoxide of iodine, I2O5, and of nitrogen, N0O5, decompose when heated. These oxides form combinations with the oxides of many other ele- ments, such as sodium or potassium oxide, e.g. Na2O.I2O5 or NaIO3, K2O.N2O5 or KNO3 ; a similar compound is potassium chlorate, KClOg or K^O-C^O^ although the simple oxide of chlorine is unknown. Now, potassium and sodium oxides are not decomposed by heat, and when these salts are heated oxygen is evolved from the pentoxide of chlorine or iodine. These elements, however, do not escape, but replace the oxygen combined with the sodium or potassium, forming chloride of the metal, thus : K2O.C12O5 = K2O+ C12 +50, and K2O + CJ2=2KC1 + O, or, summing up both changes in one equation, 2KC1O3 = 2KC1 + 3O2. Nitrate of potassium, on the other hand, loses only one atom of oxygen, leaving nitrite:
DECOMPOSITION BY HEATING 15
Oxygen is a colourless gas, without smell or taste ; it can be liquefied, at a high pressure and a* low temperature, to a pale blue liquid boiling at —182°. Most elements unite directly with it, often with such a rise of temperature that incandescence is produced ; in such a case the pheno- menon is termed "combustion/' In many instances, for example when iron rusts, the oxidation is not attended by any measurable rise of temperature, although; in all cases heat is evolved, but in some cases extremely slowly.
Chlorine, bromine, and iodine are generally prepared by heating together a chloride, bromide, or iodide with man- ganese dioxide an^ sulphuric acid diluted with water. Here the first change is the formation of the halogen hydride, HC1, HBr, or HI. The hydride, however, is
+ -
ionised in water, and the HCl.Aq., for example, at once reacts with the MnO9, forming non-ionised water and
MnCl4. Aq, thus : MnO2 + 4HC1. Aq. - M n C 14. Aq. + 2H>7O. Tetrad manganese, however, appears not to be able to co-exist with chlorine in solution ; hence the manganese
loses an electron and becomes Mn, the lost charge neutralis- ing one of the charged chlorine ions, which escapes in an
electrically neutral state. Even then,, however, the Mn, though capable of existence at low temperature, still loses a charge, and a second chlorine atom is liberated in a non-
ionised state. Hence the whole change is: Mn Cl4.Aq.
+ + -
= MnCl2.Aq. + C12. Summing all these changes in one equation, we have: MnO2 + zNaCl.Aq. + 2H2SO4.Aq. - MnSO4. Aq. + Na2SO4. Aq. + 2H2O + C12 ; or, if hydro- chloric acid alone be warmed with manganese dioxide, MnO2 + 4HC1. Aq. « MnCl.,. Aq. + 2H2O + C12.
( 3 ) Separation of an element from a compound by displacement. — This is by far the most general method
16 MODERN CHEMISTRY
of preparing elements. The elements commonly used as displacing agents are : —
(a) Hydrogen at a red heat. — The oxide or chloride is placed in a tube of hard glass, heated to 600° or 700° in a tube-furnace, and a stream of dry hydrogen is passed through the tube. Water or hydrogen chloride is formed, and is carried on by the current of hydrogen, and the element i-s left. Indium, thallium, germanium, tin, lead, antimony, and bismuth are left in fused globules, solidifying to white lustrous metallic beads ; arsenic gasifies and con- denses in the unheated part of the tube as a grey deposit ; tellurium, which is also volatile, condenses as a lustrous metallic solid ; while iron, cobalt, nickel, copper, and silver do not fuse at that temperature. The first three remain as grey powders, the copper as a red powder, and the silver in a white spongy condition. These metals can be fused by heating them in a crucible to a sufficiently high tem- perature ; it is well to use a " flux," or substance to make them flow, such as sodium carbonate or borax ; the flux fuses, and dissolves any film of oxide off the surface of the metallic beads, and they then join up to form a single mass of molten metal.
(1} Displacement by means of sodium at a red heat. — The chlorides of beryllium, magnesium, calcium, strontium, barium, aluminium, scandium, yttrium, lantha- num, ytterbium, cerium, thorium, vanadium, niobium, and tantalum are all reduced when added to sodium kept melted in an iron crucible. For boron, silicon, and titanium the double fluoride is more convenient, for the chlorides are volatile liquids. The process for manufacturing magne- sium, which is carried out on a large scale, may be more minutely described as an example. The double chloride of magnesium and potassium, MgCl2.KCl, carefully dried, is mixed with sodium in proportion to unite with the chlorine of the MgCl2, the sodium being in small lumps. The iron crucible containing the mixture is heated ; a violent reaction takes place, and magnesium is liberated :
DISPLACEMENT 17
MgCl.2. KC1 + 2Na - Mg + zNaCl + KC1. As magnesium is volatile, and can be distilled, it is purified by this operation. The contents of the crucible are treated with water ; the potassium and sodium chlorides dissolve, and the globules of magnesium are collected, dried, and placed in a crucible, through the bottom of which a tube is fixed reaching nearly to the lid, and projecting some distance below the bottom. This crucible is placed in a furnace, and on raising the temperature, the magnesium volatilises up, passes down the tube, and the vapour condenses in the cooler part of the tube which projects below the furnace. This particular method of distillation is called destillatio per descensum. The other elements mentioned are too little volatile to admit of purification by this means. In their case, the cooled mass is treated with alcohol in order to remove the excess of sodium, and then with water to dissolve the resulting salt ; the element is left in the state of powder.
(c) Displacement by means of magnesium at a red heat — This process is sometimes used to prepare the element from its oxide. A mixture is made of magnesium filings with the oxide of the element, and it is heated in an iron crucible. The resulting mass is then treated with hydrochloric acid to remove the oxide of magnesium, which is thus converted into the soluble chloride. It is, of course, essential that the liberated element shall not be attacked by hydrochloric acid. The process works for the preparation of boron, silicon, and titanium.
(</) Displacement by heating the oxide with car- bon.— This process is of the most general application. If the element is volatile, it is distilled from an iron or fire- clay retort ; in this way sodium, potassium, rubidium, arsenic, zinc, and cadmium are prepared. If non-volatile at a red heat, a mixture of the oxide with charcoal is heated to bright redness in a clay crucible. On a manu- facturing scale, coal or coke is substituted for the charcoal. The process is applicable to the production of indium,
VOL. II. B
i8 MODERN CHEMISTRY
thallium, germanium, tin, lead, manganese, iron, cobalt, nickel, and copper. To exemplify this method, four instances will be described — the preparation of phosphorus, sodium, zinc, and iron.
Phosphorus. — The commonest natural compounds of phosphorus are phosphorite or calcium phosphate, Ca3(PO4)2, and gibbsite or aluminium phosphate, A1PO4. It is accordingly convenient and economical to prepare phosphorus from one of them. The process depends on the displacing action of carbon on the oxide at a high temperature. There are two methods of effecting this. The first is : the phosphorite is mixed with dilute sulphuric acid ; the hydrogen of the sulphuric acid replaces the cal- cium of the calcium phosphate: Ca3(PO4)9 + 3H9SO4.Aq = 3CaSO4+ 2H3PO4.Aq. Coke or charcoal is impreg- nated with the phosphoric acid and heated to redness, when the phosphoric acid loses water : HgPO4 = HPOg + H2O. The mixture of metaphosphoric acid, HPO3, with carbon is charged into retorts of Stourbridge clay, the mouths of which are attached to a vertical copper tube, the lower end of which dips under water. On raising the retorts to a white heat, phosphorus distils over and condenses in the water. The final equation is : 4HPO3+ I2C = 2H2 + P4 + I2CO. By the second method, the calcium and alumi- nium phosphates are mixed with silica and carbon, and distilled from an electric furnace heated to whiteness by an arc in its interior.
Sodium. — A mixture is made of "spongy iron" (see p. 19) and pitch. This mixture is heated to redness in order to decompose the pitch, which consists of compounds of carbon and hydrogen. These compounds are decom- posed, and a part of the carbon is left mixed with the spongy iron, while the hydrogen escapes in combination with the rest of the carbon. To this mixture, placed in an iron crucible, caustic soda is added ; the lid of the crucible, which is furnished with a curved tube sloping downwards to a condenser, is fixed in place, and the
ZINC AND IRON 19
crucible is heated in a furnace to bright redness. The carbon removes oxygen both from the hydrogen and the sodium, and sodium and hydrogen pass over into the condenser along with carbon monoxide, the sodium alone condensing, for the others are gaseous and escape. The equation is: 2NaOH + zC = 2CO + H2+ 2Na. The con- denser consists of a flat hollow copper vessel ; the sodium is raked out as it accumulates.
Zinc. — The chief ore of zinc is the sulphide. To convert it into the oxide, it is roasted on a flat hearth in a current of air : 2ZnS + $O2= 2ZnO + 2SO2. The oxide is mixed with small coal (slack) and placed in cylindrical retorts of fireclay. These retorts have pipes of rolled sheet-iron luted to the open ends with fireclay ; they are packed into a furnace in tiers, and the temperature is raised to bright redness. The coal distils first, giving off coal- gas, which expels air from the retorts. When the tem- perature exceeds 1000°, the zinc distils and condenses in the iron pipes. It happens that almost all zinc ores contain cadmium sulphide, which, like zinc sulphide, is converted into oxide by roasting ; and on distillation, the cadmium, which is the more volatile metal, distils over first and condenses in the outer portion of the tubes. These are untwisted and the metal removed with a chisel.
Iron. — The chief ores of iron are the carbonate and the oxide. The former is practically always mixed with clay (clayband) or with coal (blackband), and generally contains sulphur and phosphorus in the form of calcium sulphate, CaSO4, and calcium phosphate, Ca3(PO4)9. The sulphur is sometimes present in the form of iron pyrites, FeS2. The ore is roasted to expel carbon dioxide, thus: 4FeCO3 + O2 = 2Fe2O3 + 4CO2. If it were then in its impure state smelted with coal, the iron would not flow, but would remain mixed with the clay. However, this process, if the ore is pure and charcoal is used as fuel, yields a mass of iron sponge, which can be heated and welded by hammering into a coherent mass.
20 MODERN CHEMISTRY
The process is still used by Africans, and was at one time universal. On the large scale, however, it is necessary to add lime in order to form a flux with the clay. Clay con- sists of a compound of silica, SiO0, and alumina, A12O3, and with lime it melts to a glassy slag. Alternate layers of coal, lime, and the roasted ore are fed in at the top of a blast-furnace, a tall conical erection of firebrick, strength- ened by being bound with iron hoops ; at the bottom there is a "crucible," or receptacle for the molten iron, which can be discharged when required by forcing a hole in its side with an iron bar. There are also holes which admit water-jacketed tubes or "tuyeres," which convey a blast of air heated to about 600° to increase the temperature of combustion of the coal. Here the reduction takes place in the upper part of the furnace, owing to the carbon monoxide formed by the combustion of the coal in the lower part of the heated mass ; it ac:s on the oxide of iron thus: Fe2O3+ 3CO = 2Fe + 3CO2. As the iron passes down the furnace it melts, and is met by the fused slag ; it then coheres and runs into the crucible, whence it is drawn off from time to time.
Carbon unites with molten iron, forming a carbide ; hence the product of the blast-furnace is not pure iron, but a mixture of iron with its carbide, and also with its sulphide and phosphide, if the ore has contained sulphates or phos- phates. When such impure iron is brought in contact with oxygen in a molten or semi-molten condition, the carbon, sulphur, and phosphorus are oxidised mostly before the iron. If lime be present, sulphate and phosphate of calcium are formed. The modern process of removing these impurities is to pour the molten metal into a pear- shaped iron vessel lined with bricks made of magnesia ; while it is molten, air is blown through the metal, and the carbon burns to carbon dioxide ; the sulphur and phosphorus are likewise oxidised and combine with lime, a layer of which floats on the surface of the molten metal. When these impurities have thus been removed in the " Bessemer
DISPLACEMENT BY OXYGEN 21
converter," the metal is poured into a mould. Steel is a mixture of iron with a trace of its carbide, and it is produced by mixing with the blown iron, before it is poured, a quantity of iron containing carbon and manganese (a metal which confers valuable properties on iron). The quantity of carbon in steel may vary between O.6 and 1.5 per cent. ; with the content of carbon varies also the quality of the steel ; that with a small proportion is soft, with a high proportion hard.
(e) Displacement by means of Oxygen. — Oxygen is used in Deacon's process to liberate chlorine from hydrogen chloride. The latter gas, mixed with air, is passed through a chamber kept between the limits of temperature 375°-4OO°, containing bricks soaked with cupric chloride, CuCl0. At this temperature the cupric chloride decomposes into cuprous chloride, CuCl, and free chlorine, but the cuprous chloride is reconverted into cupric chloride at the expense of the chlorine produced by the interaction of the hydrogen chloride and the air, thus : 4HC1 + O2 = 2H2O + 2CI2. The cupric chloride is again decomposed. This kind of action, where a limited quantity of a substance, itself not permanently changed, causes an apparently unlimited change in other reacting bodies, is termed " surface action," for its rate is dependent on the extent of the surface of the agent ; and the name "catalysis" is sometimes given to such an action. The action would take place independently of the catalytic agent, but at a very slow rate ; the presence of the catalyser has the effect of greatly increasing the rate at which the change takes place. The chlorine thus prepared is not pure, but mixed with the nitrogen and argon of the air, but it serves for some purposes. The rate of such action of oxygen in displacing bromine or iodine from their compounds with hydrogen is much greater, and at a high temperature the elements could be formed thus, but they are not usually produced in this way.
The preparation of nitrogen may be also regarded as a
22 MODERN CHEMISTRY
displacement by means of oxygen. Ammonia burns in oxygen, thus: 3NH3+ $O2= 3H2O + N2, but at the same time some of the nitrogen unites with the oxygen and forms NO9, nitric peroxide : this gas interacts with the ammonia, forming ammonium nitrate and nitrite, NH4NO3 and NH4NO9. If, however, the oxygen be not free, but in combination with an easily reduced metal, such as copper, it will combine with the hydrogen of the ammonia at a red heat, setting free the nitrogen. Another method involves the mutual displacement of nitrogen from its oxide by means of hydrogen, and from its hydride, ammonia, by oxygen: 2NH3 + N2O3= 3H2O + 2N2. This method is, however, usually represented by the equation NH4NO0 = 2H2O + N2; for ammonium nitrite, NH4NO2, may be regarded as a compound of N2O3 with 2NH3 and H2O. To obtain nitrogen by this method, since ammonium nitrite is not easily obtained, a solution of ammonium chloride may be warmed with one of sodium nitrite. The equation is then : NaNO2. Aq + NH4C1. Aq = 2H2O + N2 + NaCl. Aq. Another convenient method is to warm together solu- tions of sodium hypobromite and ammonium chloride ; the former loses oxygen readily, which combines with the hydrogen of the ammonia according to the equation : 3NaOBr.Aq. + 2NH4Cl.Aq. = sNaBr.Aq. + 3H0O +
2HCl.Aq. + N2.
Although sulphur, selenium, and tellurium burn in oxy- gen, still they may be displaced from their hydrides, H2S, H2Se, and H0Te, by means of oxygen at a red heat, provided the oxygen is present only in sufficient quantity to combine with the hydrogen, thus : 2H2S -f O2=2H2O + S9. Aqueous solutions of these compounds, too, are decomposed on standing, in contact with air, owing to similar displacement. Oxygen may displace mercury from its sulphide, cinnabar, HgS, which is the common ore of mercury ; here the sulphide is roasted in air, when the sulphur combines with the oxygen to form sulphur dioxide, a gas at ordinary temperature ; and mercury is liberated.
DISPLACEMENT OF ELEMENTS 25
also in the gaseous form, but condensing at temperatures below 358°.
(/) Displacement by use of Fluorine, Chlorine, and Bromine. — Fluorine, chlorine, and bromine may also be employed as displacing agents for nitrogen and oxygen.
A current of fluorine led through water displaces the oxygen, forming hydrogen fluoride ; but the oxygen is in an allotropic state (see Part i. ), called "ozone." Again, if a stream of chlorine is passed through, or if bromine- water be added to, a solution of ammonia, the hydrogen and chlorine combine, while the nitrogen is set free : 2NH3.Aq + 3C12 = 6HC1 + N9; but as ammonia com- bines with hydrogen chloride, the reaction 6NH3 + 6HC1 = 6NH4C1 occurs simultaneously; the complete equation is the sum of these two : 8NH3. Aq + 3d., = 6NH4C1. Aq + N2.
Chlorine, added to a solution of bromide or iodide of a metal, displaces the bromine or iodine ; here the non- ionised chlorine becomes ionised at the expense of the charge on the ionised bromine or iodine, while the latter
+ - + -
lose their charges, thus : 2KBr. Aq + C12. Aq. = 2KCl.Aq + Br2. Aq. Similarly, bromine displaces iodine from a soluble iodide. But iodine displaces chlorine from the nearly insoluble silver chloride. Here, the iodine is still less soluble than the chloride ; and as chloride dissolves, the less soluble and therefore non-ionised iodide is formed.
(g) Many metals are able to displace others. Thus, iron placed in a solution of a copper salt displaces the copper ; copper displaces silver ; silver, gold. In all these cases the action is doubtless an electrical one, and dependent on the replacement of a metal of lower by one of higher electric potential ; that of higher potential becomes ionised, while that of lower assumes the metallic state,
thus: CuCl2.Aq + Fe = FeCl2.Aq + Cu; 2AgNO3.Aq + Cu - Cu(NO3)2. Aq + 2 Ag.
24 MODERN CHEMISTRY
(h} There are some plans of obtaining elements which, though they can be referred to one or other of the three general methods exemplified already, are, on account of their complexity, better treated separately. Among these are the methods of separating hydrogen. The metals of the alkalies and alkaline earths attack water, forming hydr- oxides and liberating hydrogen : zNa + 2H9O = 2NaOH + H2; Ca+ 2H2O = Ca(OH)2 + H9. Magnesium powder, boiled with water, gives off hydrogen slowly ; but zinc requires the presence of an acid, and must not be pure, i.e. there must be a foreign metal present to serve as the anode. The impurity usually present in commercial zinc is lead ; the acid, for instance, sulphuric acid, is present in
r + +
dilute solution as ions of HH and SO4 ; the SO4 removes
+ +
the surface layer of the zinc as Zn, while the negative charge is transferred to the lead, which is in metallic contact with the zinc. This charge is neutralised by the » + +
positive charge of the HH, which, on being discharged, escapes in an non-ionised state. It may then be collected over water, in which it is very sparingly soluble. Hydro- gen, while it is on the point of discharging and is still in the ionised state, may be used to liberate certain elements from their oxides or chlorides. Zinc and hydrochloric acid,
I- \+~
for instance, in a solution of stannous chloride, SnCl2.Aq,
causes a deposition of tin owing to the exchange of charge ; the hydrogen retaining its charge instead of parting with it to the lead or other impurity in the zinc, while the tin is discharged in its stead. If zinc and hydrochloric acid are placed in contact with silver chloride, AgCl, which is an insoluble compound, the hydrogen remains charged, while the silver parts with the chlorine, the latter remaining in solution with negative charge. Lastly, if generated in a
solution of ferric chloride, Fe Clg.Aq, the zinc goes into solution as before ; and the positive electricity is provided
PROPERTIES OF ELEMENTS 25
by the loss of a positive charge provided by the ferric ions
+ + - changing to the ferrous ions of ferrous chloride, FeCl2. Aq,
+ -
and another molecule of HCl.Aq exists in solution. The valency of the iron is lowered. Such processes are gene- rally termed reduction ; the hydrogen is said to be in the " nascent state/' and is named the "reducing agent."
Metallic iron, manganese, cobalt, and nickel at a red heat remove oxygen from water with liberation of hydro- gen: 3Fe + 4H2O = Fe3O4 + 3H2; 2Co + 2H2O = CoO + O2. Conversely, a current of hydrogen passed over these oxides at a red heat will combine with their oxygen, reducing them to metal. This is an instance of mass- action. From the equations given above, it is seen that hydrogen is formed ; it does not remain in the tube to re-form water ; if it did, there would be a state of balance or equilibrium, all four substances remaining together in proportions depending on the temperature and on their nature ; in the current of steam, however, the hydrogen is carried on, and is no longer present to act on the oxide of the metal. And in the converse action the hydrogen conveys the steam away, so that it can no longer be deprived of oxygen by the metal.
As already remarked, carbon monoxide has a similar reducing action on the oxides of the more easily reducible elements. The product in this case is the dioxide, CO0, for example, Fe2O3+ 3CO = 2Fe + 3CO2. This action requires a red heat. Another reducing agent, applied by- fusing the oxide with it, is potassium cyanide, KCN ; it is converted into the cyanate, KCNO. The metal thallium may be prepared by its help, T19O + KCN = 2T1 + KCNO. As the cyanide is somewhat expensive, it is used only in special cases.
An instance has already been given of the mutual reduc- tion of two compounds in the case of nitrogen. Similar instances are known with lead and with sulphur. The chief ore of lead is the sulphide, a natural product termed
26 MODERN CHEMISTRY
galena. It is roasted, i.e. heated in contact with air to a red heat. After a portion has been oxidised to sulphate, PbS + 2O9 = PbSO4, the temperature is raised, when the sulphide and the sulphate mutually reduce each other : PbS + PbSO4 = 2Pb + 2SO2. With sulphur the partial burning of sulphuretted hydrogen may be explained in a similar manner; the reaction, 2H2S + O9 = 2H2O + S0, may be represented as the formation of water and sulphur dioxide by the complete combustion of one-half of the hydrogen sulphide, and its reaction with the remaining sulphide, thus: 2H2S + SO2= 2H2O + 38. And, as a matter of fact, that reaction does take place on mixing the two gases in the required proportion of two volumes of hydrogen sulphide with one of sulphur dioxide.
The Properties of the Elements. — It has been cus- tomary to divide the elements into two classes, the metals and the non-metals. As we have seen, this classification is a completely arbitrary one ; for there are some elements capable of existing in both states. The name " metal " was originally given to seven substances, all alike in possess- ing that bright lustre known as "metallic." These were gold, silver, mercury, copper, iron, lead, and tin. But in the Middle Ages bismuth and antimony were isolated in a fairly pure state, and these, together with zinc, were at first not received into the class, but were regarded as spurious ; for they were brittle and easily oxidisable. Although there is no reason for retaining the division, yet it is often convenient. Bodies which possess metallic lustre have the power of conducting electricity better than transparent bodies, and they are also relatively good conductors of heat.
The elements exist in various physical states. Those which are gases at the ordinary temperature, however, have all been condensed to the liquid state by sufficient reduction of temperature. The lowering of temperature is most easily produced by means of liquid air, now a cheap commodity. To liquefy air, it is compressed by a pump to a pressure of 150 atmospheres ; it then traverses a coil of copper pipe,
PROPERTIES OF ELEMENTS 27
and escapes from an orifice at the lower end. Now, compressed air has some resemblance to a liquid, for when it expands, as when a liquid changes to gas, heat is absorbed. The rapidly escaping air becomes cold, and in passing up over the coil of tube through which it has de- scended, it cools the pipe, so that the air passing down becomes colder and colder ; finally, it is so cooled that it liquefies, and escapes from the orifice in a liquid state. It may be poured from one vessel to another, with little loss by evaporation ; and if other gases be allowed to stream into a tube cooled by its aid, they too are liquefied. The principle of liquefying hydrogen is the same, for its boiling- point lies so low that it cannot be liquefied by the aid of liquid air. That of helium is still lower, but it too has yielded when compressed into a tube cooled by liquid hydrogen.
The elements which are gases at the ordinary temperature are hydrogen, helium, neon, argon, krypton, xenon, nitro- gen, oxygen and ozone, fluorine, and chlorine. The first seven are colourless, both in the gaseous and the liquid state. Oxygen is a colourless gas, but forms a pale blue liquid ; gaseous ozone has a blue colour ; fluorine is pale yellow ; and chlorine has a greenish-yellow colour. It forms a white solid, which, however, melts to a bright green liquid. Bromine is a dark red liquid at atmospheric temperature, but above its boiling-point, 59°, it is a deep red gas. Iodine is a blue-black solid, melting to a black liquid at 114°, and giving off a violet vapour. Ozone and the " halogens," as fluorine, chlorine, bromine, and iodine are called, have all a powerful odour, and act on the skin in a corrosive manner. Chlorine and bromine are soluble in water.
Among the other non-metallic elements are boron, a black, dusty, infusible powder ; carbon, in its ordinary form an amorphous (i.e. non-crystalline) black substance, of which the most familiar variety is charcoal ; carbon does not fuse, but at the enormously high temperature of the electric arc
28 MODERN CHEMISTRY
it volatilises ; silicon, a blackish-brown powder, melting at bright redness to a lustrous liquid, which solidifies in shining black lumps ; phosphorus, a waxy, pale yellow solid, melt- ing at 44.4° ; sulphur and selenium, yellow and brown-red solids, the former melting at 115° to a brown liquid, and boiling at 446° ; the latter forming a black liquid at 217°, and a black vapour at 665°.
The metals of the alkalies, as they are usually called, lithium, sodium, potassium, rubidium, and cassium, are soft white metals, at once attacked by water, and oxidised readily by air, caesium, indeed, taking fire spontaneously. To protect them from oxidation, they must be kept under rock-oil or ligroin, a compound which contains no oxygen. Of these, caesium has the lowest and lithium the highest melting-point. The metals calcium, strontium, and barium are sometimes named the " metals of the alkaline earths." They are hard white bodies, also, like those of the sodium group, oxidising readily on exposure to air, and at once attacked by water. Magnesium, zinc, and cadmium are noteworthy, inasmuch as their temperature of ebullition is not so high that it cannot be reached in an ordinary furnace; they can therefore be distilled. Magnesium and zinc are hard and brittle ; cadmium is softish, like lead, and of a somewhat greyer tint.
The remaining elements may be classed under the head- ings, " hard," " soft," " brittle," &c. This implies only their behaviour at ordinary temperatures ; at higher or lower temperatures the properties are materially changed. Mercury, for example, below -40°, is malleable; lead is brittle.
(a) Malleable metals :—
(1) White^ ductile, moderately hard: — beryllium, alumi- nium, gallium, indium, tin, silver, nickel. Red, copper. Tellow, gold.
(2) Grey-white, ductile, and moderately hard: — iron, manganese, cobalt.
METALS 29
(3) G 'rey- white and soft ; ductile: — thallium, lead ; some- what harder, and fusible only at a <very high temperature : — rnodium, ruthenium, palladium, platinum, iridium.
( b ) Liquid metal : — mercury.
(c] Brittle metals : —
(1) White, hard: — antimony, bismuth, tellurium, zirco- nium, didymium (a mixture), osmium, germanium. Less hard, arsenic.
(2) Grey, hard : — lanthanum, cerium, yttrium, uranium.
( 3 ) Grey powders, acquiring metallic lustre under the bur- nisher : — thorium, niobium, tungsten.
(4) Black powders : — tantalum, titanium.
The elements scandium, samarium, and gadolinium have not been prepared.
Although the external properties of the elements does not show any obvious relation to their order in the periodic table (see Part I.), yet it may be generally remarked that the density increases as each column is descended. Among the lightest of the elements are lithium, beryllium, magnesium, and aluminium, at least in the solid state ; whereas osmium, iridium, platinum, and gold are among the heaviest. But much more must be ascertained regarding their properties before a satisfactory comparison can be made.
CHAPTER II Classification of Compounds — The Hydrides.
Classification of Compounds. — Compounds of the elements may be divided conveniently into six classes : —
The Hydrides ;
The Halides ;
The Oxides and Sulphides (with Selenides and
Tellurides) ; The Nitrides and Phosphides (with Arsenides and
Antimonides) ;
The Borides, Carbides, and Silicides ; The Alloys.
Compounds can be prepared by many methods ; it is not so easy to classify them as it is to arrange into classes the methods of preparation of elements. As a rule, the pre- paration is carried out by one of the following methods : —
(a) The interaction of elements ;
(£) The action of an element on a compound ;
(c) The action of heat on a compound ;
(d) The interaction of compounds ;
(e) The addition of one compound to another.
These methods shall be considered in relation to each of the groups of compounds named above.
The Hydrides.
(a) The Interaction of Elements. — Lithium, sodium, and potassium, when heated to 300° in an iron tube in a
INTERACTION OF ELEMENTS 31
current of hydrogen, form white waxy compounds ; that of lithium has the formula LiH ; as the sodium compound has the formula Na9H, its existence is difficult to reconcile with the usual valency of either hydrogen or sodium, for these elements in all other compounds behave as monads. It would repay further investigation. It decomposes at 421°.
Iron, nickel, palladium, and platinum, when heated gently in hydrogen, absorb the gas. Meteoric iron, indeed, has been known to give off, on heating, 2.85 times its volume of gas. This natural variety of iron contains about 6 per cent, of nickel. Palladium, gently warmed in an atmosphere of hydrogen, absorbs over 900 times its volume of that gas, corresponding to 4.68 per cent, of the weight of the body produced. It is difficult to determine whether or not the palladium is in chemical combination with the hydrogen, or whether the hydrogen is in a state analogous to solution, for it is known that a solid can exert solvent power. There is a considerable rise of temperature accom- panying the absorption ; and if palladium, in a state of sponge, is placed in contact with a mixture of oxygen and hydrogen, the mixture may be made to explode. A ther- mometer-bulb coated with palladium sponge is a good test for the presence of an explosive mixture of marsh-gas and air in mines, for the rise of temperature produced is an in- dication of danger. These metals absorb hydrogen more readily if they are made the negative electrodes of a battery with which dilute sulphuric acid is electrolysed. Iron shows a very curious behaviour under these circum- stances. If a thin plate of iron is made to close the top of a barometer-tube full of mercury and a small cell be con- structed on it, hydrogen will pass through the iron, when the plate is made the kathode, and will depress the mer- cury in the tube. No other metal, so far as is known, shows this peculiarity ; it would appear that the hydrogen in the ionic state can penetrate the iron.
Carbon, heated to 1200° in an atmosphere of hydrogen, unites with it to form marsh-gas (methane), CH4. Only
32 MODERN CHEMISTRY
a small percentage of the hydrogen, however, enters into combination ; a balance soon establishes itself between the number of molecules of methane being formed and decom- posed in unit time. At a higher temperature, that of the electric arc, acetylene, C2H2, is formed, owing to the decomposition of the methane into that gas and free hydrogen: — 2CH4 = C9H9 + 3Ht>. Other compounds of carbon and hydrogen are formed simultaneously, and there again appears to be a state of equilibrium produced between the various hydrocarbons formed. With nitrogen, NH,, it appears to be impossible to induce hydrogen to enter into direct combination at such temperatures ; but if electric sparks be passed through a mixture of hydrogen and nitro- gen, combination to a limited extent ensues. Should the ammonia, NH3, be removed by having water, or, better, dilute sulphuric acid, present, the combination proceeds until all the gases, if they were originally present in the correct proportion — one volume of nitrogen to two volumes of hydrogen — have combined. Conversely, if sparks be passed through ammonia gas, there is nearly, but not quite, complete decomposition into its constituents. This enables the volume relations of ammonia to be demonstrated ; for it is found that two volumes of ammonia gas can be decom- posed into two volumes of nitrogen and six volumes of hydrogen. This is symbolised by the equation —
2NH3 = N2 + sH,
Weight 2 ( 14 + 3) 28 3 (2) grams. Volume 2(22.4) 22.4 3(22.4) litres.
The hydrogen can be nearly completely removed by ab- sorption with palladium-sponge, and the nitrogen remains.
Water, H20, is more completely formed than any one of the previously mentioned compounds by the interaction of its elements. A mixture of oxygen and hydrogen, in the proportion of one volume of oxygen to two of hydrogen, is exploded by heat ; this is most easily done by passing an electric spark through the mixture. While the position of
COMBINATION OF HYDROGEN 33
equilibrium for a mixture of nitrogen, hydrogen, and am- monia lies at such a point that very little of the compound is present, but chiefly the uncombined gases, the contrary is the case with hydrogen and oxygen. Here nearly all the oxygen and hydrogen combine, and only a trace remains uncombined. Combination may be made to take place slowly at much lower temperatures ; even at 300° slow combination occurs. Colloidal platinum, prepared by mak- ing an electric arc between poles of platinum under pure water, which appears to consist of very finely divided platinum disseminated through the water, has the power of causing union of oxygen and hydrogen left standing in contact with it, even at the temperature of the atmosphere. On the other hand, if water- vapour be raised to a very high temperature, above 1800°, decomposition into its consti- tuents takes place with considerable rapidity ; so that it is possible to obtain a mixture of oxygen and hydrogen by passing steam through a tube in which a spiral of platinum wire is kept at a white heat by means of an electric current. These actions are therefore termed " reversible," and they are expressed by such equations as —
CH4 ^ C + 2H9 ; 2H9 + O., ^ 2H20 ;
Hydrogen also combines with sulphur when passed through a flask containing boiling sulphur, and sulphuretted hydrogen, H2S, decomposes when raised to a low red heat.
Interesting relations are to be seen with the compounds of the halogens with hydrogen. In preparing fluorine by the electrolysis of hydrogen-potassium fluoride, KHF, in presence of hydrogen fluoride, H2F9, it is possible, by stop- ping the exit of the hydrogen, to cause a bubble to pass the bend of the U-tube and to rise into the fluorine ; the instant the gases unite there is a sharp explosion. This shows that these gases unite even in the dark to form H2F2. Chlorine and hydrogen, on the other hand, do not com-
VOL. II. C
34 MODERN CHEMISTRY
bine in the dark, but, when exposed to diffused daylight, slow but complete combination ensues ; in bright sunlight, or when illumined by the light from burning magnesium, the mixture of gases explodes, forming HC1. Bromine and hydrogen unite to form HBr when a current of hydrogen, having bubbled through a wash-bottle of bromine, passes through a red-hot tube ; with excess of hydrogen the union is practically complete. Iodine and hydrogen, on the contrary, unite very incompletely to produce HI ; and if hydrogen iodide be heated, a large proportion of it is decomposed into hydrogen and iodine. This change has been investigated much more completely than other changes of the same character already mentioned ; and as it is characteristic of all such reversible reactions, we shall con- sider it in somewhat greater detail.
The rate at which hydrogen iodide is produced from a mixture of hydrogen and iodine at any constant tempera- ture is much more rapid than that at which the reverse change of hydrogen iodide into iodine and hydrogen takes place. This rate was not difficult to determine. Weighed quantities of iodine were placed in a tube filled with hydro- gen, and after heating the sealed tube for a sufficiently long time for equilibrium to be established, it was opened under water. The hydrogen iodide formed at once dissolved in the water, and the residual hydrogen was measured. The amount of uncombined iodine remaining in the water was then estimated by known processes. It was thus possible to find the ratio of the combined to the uncombined hydro- gen. Now, it was discovered many years ago that the rate of chemical change depends on the amount of each of the reacting substances present in unit volume — a condition ex- pressed by the term "active mass." Thus, if we double the amount of hydrogen in the mixture of the gases men- tioned, we double its " active mass/' Let /9 denote the number of molecules in unit volume of the iodine gas, and ^2 that of the hydrogen, and let ^hi be that of the hydrogen iodide formed by their interaction. Then, as the rate of
HYDRIDES OF CARBON 35
formation of hydrogen iodide is proportional both to / and to />, it will be proportional to their product, h x /. And as H0 + I9= 2 HI, the rate of change of HI into H2 and I.? will be 2hi x ihi or 4(/>/)2. If we call the rate of forma- tion k, and that of decomposition /£', the proportion of these rates to each other will be kjk' = (h x i)/4(»i)2, if the gases are present in molecular proportions. At the tem- perature 440°, and at one atmosphere pressure, it was found that, taking the total hydrogen as unity, 0.28 was free and 0.72 combined, after a sufficient time had been al- lowed for the change to complete itself. Now, the iodine free must have been equal in number of molecules to the free hydrogen, i.e. 0.28, and the same number of atoms of iodine must have existed in combination as of hydro- gen in combination; hence 0.28x0.28/4(0.72x0.72) = 0.0375 = ^/y£'. This means that at 440° molecules of hydrogen iodide decompose into hydrogen and iodine at a rate only 0.0375 (or one twenty-sixth) of that at which combination takes place between the two gases.
(b] The action of an element on a compound leads to the formation of many hydrides. This process has been pretty fully treated in the description of the methods of preparation of elements. For example, on passing a current of hydrogen over hot cupric oxide, water, H.,0, is formed, while the oxide is reduced to copper, CuO + H9 = Cu + H2O. The oxides mentioned on p. 16 are thus reduced. It is not so usual for sulphides to lose sulphur on heating them in a stream of hydrogen ; indeed, it is only those sulphides which themselves decompose when heated that yield to such treatment ; but hydrogen fluoride, chlo- ride, bromide, and iodide are formed on heating the halides of many metals in a current of hydrogen. The process, however^ is not one which is used for the preparation of these hydrides.
(c) The third method — that of heating a compound — is also not in use as a means of preparing hydrides^ but it is often employed in order to produce the compound from
36 MODERN CHEMISTRY
which the hydride is separated. Thus, all compounds containing water of crystallisation, when heated, lose water when raised to a high temperature ; and double compounds of ammonia, too, lose ammonia on rise of temperature. Such compounds as calcium chloride, CaCl0, crystallise with water. The formula of the hydrated compound is CaClQ. 6H2O ; a similar compound with ammonia, CaCl0.6NH3, is also known ; compounds like these lose water or am- monia when heated. By this plan Faraday succeeded in liquefying ammonia, which at ordinary temperatures is a gas. Having sealed up the ammonio-chloride of calcium or of silver, AgCl.NH3, in an inverted U-tube, one leg was cooled with a freezing mixture, while the other was heated, and the gas liquefied under the combined influence of cold and pressure.
(^/) Most of the hydrides can be prepared by the fourth method — the interaction of compounds. The decom- posing agent is either water, an acid, or an alkali.
( i ) Water : — Marsh-gas, CH4, ethylene, C2H4, acety- lene, C0H0, ammonia, NH3, and phosphoretted hydrogen, PH3, may be produced by the action of water on some compounds of carbon, nitrogen, and phosphorus. Alumi- nium carbide, A14C3, yellow transparent crystals produced by heating a mixture of carbon and oxide of aluminium to whiteness in the electric furnace, on treatment with water yields pure methane, A14C3 + 1 2R,O - 3CH4 + 4A1 ( OH)g. Manganese carbide, black crystals produced by heating in the electric furnace a mixture of manganese oxide and carbon, yields a mixture of equal volumes of hydrogen and methane, Mn3C + 6H2O = 3Mn(OH)2 + CH4 + H2. Lithium, calcium, strontium, and barium carbides also formed in a similar manner in the electric furnace yield acetylene with water, Li9C.7 + 2H.7O = 2LiOH + C9H9 ; CaC2 + 2H2O = Ca(OH)2+~C2H2; The carbides of cerium, CeC2, lanthanum, LaC9, yttrium, YC2, and thorium, ThC2, yield a mixture of methane, ethylene, C.7H4, and acetylene, sometimes mixed with hydrogen ;
INTERACTION OF COMPOUNDS 37
and uranium carbide, U2C2, gives methane, ethylene, and hydrogen, but no acetylene.
Magnesium or calcium nitrides, prepared by heating metallic magnesium or calcium in a current of nitrogen, yield ammonia with water: Mg3N2 + 6H2O = 2NH3 -f 3Mg(OH).>, and calcium phosphide, produced by heat- ing lime with phosphorus, on treatment with water simi- larly gives off phosphoretted hydrogen : CagP2 + 6H2O = 3Ca(OH)0 + 2PH3. The sulphides of magnesium and aluminium, MgS and A19S3, are also decomposed by water, with production of hydrogen sulphide and the hydroxide of the metal : MgS + 2H.OH = Mg(OH)9 + H2S ; A1.,S3 + 6H.OH = 2Al(OH)g + 3H2S.
The halides of a certain number of elements are at once decomposed by water with formation of a hydride of the halogen and a hydroxide of the element. Boron, silicon, titanium, phosphorus, sulphur, selenium, and tellurium chlorides, bromides, and iodides are thus resolved. The method is practically made use of in preparing hydrogen bromide, HBr, and iodide, HI, by help of phosphorus. But the previous preparation of phosphorus bromide or iodide is unnecessary. It is sufficient to add bromine to water in contact with red phosphorus, and hydrogen bro- mide is evolved ; or to warm a mixture of iodine, water, and red phosphorus. The use of yellow phosphorus is not advisable, for the action is apt to take place too violently if it be used. It may be supposed that the phosphorus and halogen unite to form the pentahalide, which is then imme- diately decomposed by the water, thus : PBr.(or PI6) + 4H,O = H3PO4+5HBr(or sHI). The gaseous hydride may be collected over mercury or by downward displace- ment, or it may be dissolved in water and a solution of hydrobromic or hydriodic acid prepared.
A commercial method of producing hydrogen chloride, HC1, depending on the decomposition of magnesium chloride when heated in a current of steam, has been patented ; it results in the formation of a compound of
38 MODERN CHEMISTRY
oxide and chloride of magnesium, while the hydrogen of the water unites with a part of the chlorine ; the resulting gaseous hydrogen chloride is passed up towers, and comes into contact with water, thus yielding a solution of hydro- chloric acid.
(2) In many cases the compound from which the hydride is formed is not decomposed by water ; an acid, generally hydrochloric acid, must be present. The reason of this is
+ not easily explained ; it may be that the very few ions of H
and OH present in water are sufficient to effect the decom- position in some cases and not in others, and that when an acid is necessary the much larger number of ions of hydrogen present in its solution is required ; also it is known that the heat evolved during the decomposition of those compounds which are altered by water is greater than that which would be evolved by those which resist its action were they to be attacked by water. Many hydrides are prepared by the help of acids. Mag- nesium boride, Mg3B.,, yields with hydrochloric acid a trace of BH3 ; but as this compound is a very unstable gas, almost all of it decomposes into boron and hydrogen. The similar compound, Mg9Si, produced by heating a mixture of silica and magnesium powder to redness, when mixed with hydrochloric acid yields hydride of silicon, SiH4, as a colourless, spontaneously inflammable gas : — Mg2Si + 4HC1. Aq = 2MgCl2 Aq + SiH4. Arseniuretted hydrogen, AsH3, and antimoniuretted hydrogen, SbH3, are prepared from sodium or zinc arsenide or antimonide : Na3 As + 3HCl.Aq = 3NaCl.Aq + AsH3 ; Zn3Sb2 + 6HCl.Aq = 3ZnCl9.Aq + 2SbH3. These gases, however, may be ob- tained mixed with hydrogen if a solution of oxide of arsenic or antimony in hydrochloric acid, which yields chloride of arsenic or antimony, is treated with zinc. The first change is the replacement of the zinc by the arsenic or antimony,
"
thus : 2 AsCl3. Aq + 3Zn = 3ZnCl2 Aq + 2 As. Electrically
HYDRIDES 39
neutral zinc replaces positively charged arsenic, itself be- coming positively charged. The arsenic and the unattacked zinc form a couple, and the hydrochloric acid is electrolysed,
+ - + + ++ - +
2HC1. Aq + Zn = ZnCl2. Aq + 2H ; the hydrogen ion unites with the arsenic, negatively charged in the electric couple, forming electrically neutral hydride of arsenic, which escapes
+
as gas, 3 H + As = AsH3. An element in this form, capable of combination at the moment of liberation, is said to be in the nascent state, a word derived from " nascere," to be born. It differs from an ordinary element in being on the point of losing an electric charge, and it may either be evolved in
+ +
the free state by combining with itself, as H + H = H2, on giving up its charge, or it may enter into some other form of combination, as in the case explained. This process of preparing arsenic or antimony hydride is used as a test for the elements arsensic or antimony. It was devised by Marsh, and as the hydrides are very easily decomposed by a high temperature, the gas, if caused to pass through a red-hot tube, is decomposed, giving a deposit of arsenic (grey) or antimony (black). The former is more easily oxidised than the latter, and dissolves in a solution of bleaching-powder, in which the latter is insoluble. This process is particularly applicable where poisoning with arsenic or antimony is suspected.
H.2S, H2Se, H2Te. — Hydrogen sulphide, selenide, and telluride are prepared by treating a sulphide, selenide, or telluride with dilute sulphuric or hydrochloric acid : FeS
+ H9SO4.Aq - FeSO4.Aq + H2S ; Sb0S3 + 6HCl.Aq
= 2SbCl3.Aq + 3H2S. Na2Se.Aq + H2SO4.Aq =
Na2SO4.Aq + H2Se.
Adds. — Hyd'ride of fluorine, chlorine, bromine, and iodine, when dissolved in water, are termed "acids." As already mentioned, this name was originally applied to com-
40 MODERN CHEMISTRY
pounds which possess a sharp taste and change the colour of certain vegetable colouring matters. The word was later extended to apply to compounds similar in function, although not acid to taste, which attack the carbonates, causing them to effervesce, and which yield salts with the oxides of metals. All acids contain hydrogen, and it is now possible to define them in a very simple manner. An acid, in fact, is a compound which yields hydrogen ions when dissolved in water, or in some other solvent capable of causing ionisation. This definition applies to the hydrides of fluorine, chlorine, bromine, and iodine ; and also to those of sulphur, selenium, and tellurium ; for on
•f- + - +-
solution they ionise thus : HF.Aq ; HCl.Aq ; HBr. Aq ;
HLAq; H.SH.Aq; H.SeH. Aq ; H.TeH.Aq. But it is not confined to them, for the hydrogen may be united, not with a simple element, but with a complex group of
+ +
elements, as in H9SO4.Aq or HNO3.Aq. Now, in dilute solution, a solution of sulphuric acid is less ionised than one of hydrochloric acid, in about the proportion of 1:2, and it is therefore a weaker acid ; so that if a hydroxide, such as sodium hydroxide, be presented to a mixture of equal numbers of these molecules, in quantity requisite for only one of them, chloride of sodium will be formed in greater quantity than sodium sulphate ; yet, on heating a halide with sulphuric acid, because hydrogen chloride is a volatile compound, it removes itself from the sphere of action in a non-ionised state while the sodium remains as sulphate. Hence these hydrides may be thus prepared. Hydrogen fluoride, H2F2, is generally prepared by distilling calcium fluoride, a compound naturally occurring as " fluor-spar," with sulphuric acid in vessels of lead or platinum: CaF2 + H2SO4- BaSO4 + H2F2. The use of lead or platinum is obligatory on account of the action of hydrogen fluoride on glass or porcelain, the materials of which flasks and retorts are usually made ; for
HALIDES 41
hydrogen fluoride attacks the silica which they contain, forming with it silicon fluoride: SiO2 + 2H2F2 = SiF^ + 2H9O. Gold is almost the only other metal which resists the action of hydrogen fluoride. There is no such diffi- culty with the other halides. Hydrogen chloride, HC1, is prepared by distilling from a glass retort a mixture of common salt and oil of vitriol : NaCl + H.7SO4 = HNaSO4 + HC1. On a large scale this preparation is carried out in rotating circular furnaces, the mixture of salt and vitriol being delivered in through a hopper above, and at the high temperature the action goes further, and di-sodium sulphate is produced: 2 NaCl + H2SO4 = NaaSO4+ 2HCL The gas is passed up towers filled with coke, and exposed to a descending stream of water, in which it dissolves, forming a saturated solution of hydrochloric acid, or, as it used to be called, "muriatic acid" (from "muria," brine).
Hydrogen bromide, HBr, and iodide, HI, may similarly be produced by distilling together bromide or iodide of sodium or potassium with exactly the right weight of sulphuric acid for the equation 2KBr (or zKI) + H2SO4.Aq= K2SO4.Aq + 2HBr (or 2HI). But in these cases, the hydrogen bromide or iodide is very apt to exert a reducing action on the sul- phuric acid, depriving it of an atom of oxygen, thus : — H2SO4 + 2HI - H2SO3 + H2O + I2. Hence it is advis- abfe to use phosphoric acid, H3PO4, a compound not thus reduced :— H3PO4 + 2KI = H'K2PO4 + 2 HI.
All these halides come over as gases, and may either be collected over mercury or by " downward displacement," i.e. by delivering them to the bottom of a jar containing air, which owing to its less density is forced upwards, and escapes at the mouth of the jar. They cannot be collected over water, for they are readily soluble in it.
The compound HN3, termed hydrazoic acid (from the French term for nitrogen, "azote"), is also liberated in the gaseous form by warming its sodium salt with sulphuric acid. It, too, is readily soluble in water.
(3) Certain hydrides are set free by the action of an
42 MODERN CHEMISTRY
alkali, i.e. the hydroxide of one of the metals of the sodium or the calcium group. It is true that the change may be produced by other hydroxides, but they are not so efficient, and not so generally employed. Among these are ammonia, NHg, and hydrazine, NJH4. These bodies unite with acids ; for example, ammonia and hydrogen chloride form ammonium chloride, NH4C1, when mixed: — NH3 + HC1 = NH4C1. This compound is produced by a change in valency of the nitrogen atom ; in ammonia it is a triad, N'", but on union with hydrogen chloride the valency of the nitrogen becomes five, Nv. On distillation of a mixture of ammonium chloride with caustic soda or with slaked lime, either in presence or absence of water, the following change occurs :— NH4C1 + Na O H = NaCl + NH3 + H2O ; 2NH4Cl + Ca(OH)2 = CaCl2 + 2NH3+2H2O. The initial change is the formation of ammonium hydroxide, NH4OH ; this substance, being unstable when heated, decomposes into ammonia and water. Hydrazine, a com- pound of the formula N0H4, is similarly liberated from its chloride.
The usual source of commercial ammonia is coal-gas. On distillation of coal, all varieties of which contain nitrogen, it may be imagined that when methane, the principal consti- tuent of coal-gas, is strongly heated it splits into carbon and hydrogen. This hydrogen, at the moment of its formation, is in the nascent state, and it unites with the nitrogen, which is also in the nascent condition. As ammonia is very easily soluble in water, while the other constituents of coal-gas are sparingly soluble, the gas is deprived of ammonia by passing it through " scrubbers," pipes containing broken bricks kept moist with water. The ammonia dissolves, while the coal- gas passes on. The solution is next mixed with hydro- chloric acid and evaporated to dryness. The residue of ammonium chloride is then distilled with lime, as previously described. The ammonia is received in water, and brought into the market in the form of a concentrated solution, to which the name " liquor ammoniac " is given.
PROPERTIES OF HYDRIDES 43
(b) Certain double hydrides are formed by the addition of one hydride to another. Ammonia and hydrazine unite with hydrides of the halogens to form salts, such as ammo- nium chloride, NH4C1 ; but as these bodies show analogy with salts of the metals, they will be reserved until the latter are considered.
General Nature of the Hydrides. — Hydrides of lithium, sodium, potassium, iron, nickel, palladium, and platinum differ from the others in character ; they are solid bodies, decomposed by heat. Graham, indeed, who investigated that of palladium, was struck with the metallic nature of the substance, and was inclined to believe that it might be regarded as an alloy of a metallic form of hydrogen, to which he gave the name " hydrogenium ; " and it was for long believed that liquid hydrogen would show the character- istic property of metals, metallic lustre. But this anticipation has not been fulfilled. Liquid hydrogen is a colourless body ; and solid hydrogen is described as having a white crystalline appearance, like ice froth. But it must be confessed that hydrogen shows a marked similarity to metals in many of its compounds, as will be frequently seen in the sequel.
The remaining hydrides may be divided into three classes : — Those which react with neither acid nor bases, and which may therefore be described as neutral. To this class belong the hydrides of boron, carbon, silicon, arsenic, and antimony. That of phosphorus nearly falls into the same category, for its compounds with acids are very unstable. The next class — those which react with bases — comprises water and the hydrides of sulphur, selenium, and tellurium. The compounds are termed hydroxides, or, in the case of sulphur, hydrosulphides. These will be considered later, but an instance may be given here : — When lime is moistened with water it is slaked, with formation of calcium hydroxide, thus : CaO + H2O = Ca ( OH ) 2. The hydrides of fluorine, chlorine, bromine, and iodine also belong to this class ; but in their case an exchange takes place, thus : CuO + 2HCl.Aq = CuCl2.Aq
44 MODERN CHEMISTRY
= H9O. Hydrazoic acid is capable of similar reactions. Such hydrides, with the exception of water, are generally termed acids. The last group of hydrides, ammonia and hydrazine, and, in one or two isolated cases, hydrogen phos- phide, unite with acids, forming salts, thus : NH3 + HC1 = NH4C1; PH3 + HI = PH4I. It appears that the pre- sence of water is necessary for at least the first of these combinations ; for if perfectly dry hydrogen chloride is mixed with perfectly dry ammonia, no combination results. It is perhaps allowable to suppose that the presence of moisture leads to ionisation of the hydrogen chloride, and that the ionised molecule is capable of entering into com- bination, while the non-ionised molecule is without action on the ammonia. These compounds will be treated of under the heading of " salts."
The hydrides of boron, carbon, silicon, phosphorus, arsenic, and antimony are insoluble in water ; those of nitrogen, sulphur, selenium, tellurium, and the halo- gens are soluble. With the exception of certain hydrides of carbon, to be afterwards described, and water, all the rest are gases at atmospheric temperature. The fact that water is a liquid, and not, as might be expected, a gas, re- quires comment. It is noteworthy that water-gas possesses the density 9, corresponding to the molecular weight 1 8 ; hence there can be no doubt that in the gaseous state water has the formula H9O. But it is known that compounds of sulphur, which are in formulae, and in many properties analogous to compounds of oxygen, possess higher boiling- points than the corresponding oxygen compounds. For instance, bisulphide of carbon, CS2, boils at 44°, whereas carbon dioxide boils at about —80°. But water boils at 100°, and, contrary to expectation, its analogue, sulphuretted hydrogen, condenses to a liquid at a temperature much below o°. Now, it has been found by a method depend- ing on the rise of liquids in capillary tubes, that while the molecular weight of most substances in the state of liquid is identical with those which they possess in the gaseous state,
HYDROCARBONS 45
the molecular weight of water is considerably too great. The conclusion follows, therefore, that the molecular weight of water should be expressed by a more complex formula than H2O ; possibly by H4O2, or by one even more complex. Gaseous hydrogen fluoride, unlike its congeners, has a higher molecular weight than that ex- pressed by the formula HF ; determination of its density leads to the formula H2F2. These facts are probably to be explained by the view that oxygen may possess a higher valency than 2, and fluorine than I, at relatively low tem- peratures. It is not unlikely that the structural formula of
H\ /H
liquid water is >O=O< , and that of hydrogen
H/ \H
fluoride HF=FH, where oxygen acts as a tetrad and fluorine as a triad.
Hydrocarbons. — The hydrides of carbon, or " hydrocarbons," are very numerous, and form an im- portant group of substances. In many respects they are analogous to the metals, and they yield derivatives com- parable with those of the metals. The preparation of some of them has already been described ; but in order to give a more complete idea of their structure and functions, a short description of other methods of forming them is annexed.
Methane or marsh-gas, if mixed with its own volume of chlorine, and exposed to daylight — not sunlight, else the mixture would explode — undergoes the reaction CH4 + C12 = CH3C1 + HC1. The resulting gas, termed chloro- methane, is soluble in ether, a volatile liquid compound of carbon, hydrogen, and oxygen. If pieces of metallic sodium are added to the solution, the sodium withdraws chlorine from the chloromethane and a gas is evolved. On analysis, it gives numbers answering to the formula CH3. But if that were its formula, its molecular weight in grammes would occupy 22.4 litres; but 15 grammes occupy only 11.2 litres; hence its molecular weight must be 30, and not 15, and its formula cannot-be CH3, but must
46 MODERN CHEMISTRY
be C2H6. It is reasonable to suppose that the mechanism of
H\ : ; /H
the reaction is this : H— ,C— :C1 + Na Na + Cl— C— H ;
__ H/ \H
and that the two CH3 groups on liberation join together,
H\ /H
forming the complex group, H— ~C — C — H. Similarly,
H/ \H
mixing C9H6, which is named ethane, with its own volume of chlorine, a reaction takes place like that with methane, and chlorethane is formed, thus : C2H6 + C12 = C9H5C1 + HC1. Chlorethane dissolved in ether and treated with sodium yields not C2H5 but C4H10, and it may be supposed that the constitution of the new hydrocarbon,
H H H H butane, is HC — C — C — CH. A mixture of chloro-
H H H H methane and chlorethane gives with sodium an intermediate
H H H hydrocarbon, CJHL, propane, HC — C — CH. When
H H H
chlorine and propane are mixed in equal volumes, two chloropropanes result; they have identical formulas and molecular weights, and it is believed that the difference between them consists in the position of the entering atom of chlorine. In one case the chlorine replaces hydrogen attached to one of the terminal atoms of carbon, thus :
H H H
Cl C — C — CH, while in the other the medial hydrogen is H H H
H Cl H replaced : HC — C — CH. These two chloropropanes
H H H
yield in their turn two methylpropanes or butanes. Two such substances are said to be isomeric, or to exhibit isomerism with each other. The following list gives the names and formulae of some of this series of hydrocarbons ;
HYDROCARBONS
47
where the difference between their formulae is CH2, they are said to form a " homologous series."
H HCH
H Methane.
H H H
HC— C— CH
H H H
Propane.
H H H
HC— C— CH
H | H
HCH
H Isobutane.
H HCH
| H H HC— C— CH
| H H HCH
H Isopentane.
H H
HC— CH
H H
Ethane.
H H H H
HC— C— C— CH
H H H H
Butane.
H H H H H
HC— C— C— C— CH
H H H H H
Pentane.
H
HCH
H | H
HC— C— CH
H | H
HCH
H Tetramethyl-methane.
Chloromethane, if mixed with its own volume of chlorine and exposed to light, yields a dichloromethane, thus : CH3C1 + C12 = CH?C12 + HC1. This compound, which, like chloromethane, is also a gas soluble in ether, on treating its solution with sodium, loses chlorine and is converted into ethylene, thus: CH9C19 + 4Na + C19CH9 = 4NaCl
H H + C=C . The carbon atom, it will be observed, is still
H H a tetrad, but the two atoms are connected by a " double
48 MODERN CHEMISTRY
bond." Homologues of ethylene are known, of which the following are a few : —
H H H H H
C=C HC— C=C
H H H H
Ethylene. Propylene.
HHHH HHHH H H
HC— C— C=C HC— C=C— CH HC— C— CH
HHHH H H || H
HCH
Butylenes.
These hydrocarbons are characterised by the facility with which they combine with the halogens, forming oils ; they have, therefore, been termed "defines," or " oil-makers.'* They also unite with nascent hydrogen, and are converted into paraffins, as the members of the former group are termed. The equations which follow illustrate this : —
H2C Cl H9CC1 CH., CH3
1+1=1 || " + 2tt= |
H2C Cl H,CC1 CH2 CH3
By the further action of chlorine on dichloromethane, trichloromethane, or chloroform, CHC13, is produced. Chlorine can also be withdrawn from chloroform by sodium, and acetylene, C2H2, is formed: HCCl3 + 6Na + Cl3CH = 6NaCl + HC=CH. Here the two carbon atoms are represented as united by a treble bond, and each carbon atom is still believed to remain tetrad. Acetylene is also characterised by the ease with which it unites with chlorine, forming a tetrachlorethane : HC^CH + 2C12 = C19HC— CHC13. Here, also, other members of the series are known.
The passage of acetylene through ^ red-hot tube is attended by " polymerisation ; " that is, two or more mole- cules unite to form a more complex one. In this case, three
HYDROCARBONS 49
molecules of acetylene combine to form a molecule of the formula C6H6, a compound to which the name benzene is applied. It is produced in large quantity by the distillation of coal, and is separated from coal-tar oil by distillation. Its carbon atoms are imagined to form a ring, because, among other reasons, it yields only one mono-chloro-sub-
H H H
C— C=C stitution product : || gives on treatment with
C— C=C
H H H H H Cl C— C=C chlorine, C19, || | ; and as all the hydrogen atoms
^ c_ c=c
H H H
in the molecule are symmetrically arranged with respect to the carbon atoms, this condition is fulfilled.
The four first members of the methane series are gases ; those containing a greater number of atoms of carbon up to eleven are liquids, and the higher members are solids. The paraffin oil which is burned in lamps consists of a mixture of the liquid members, and paraffin candles largely consist of the solid members. They are all practically insoluble in water. The olefines have similar physical pro- perties, and benzene is a volatile liquid. Iodine, sulphur, and phosphorus dissolve in the liquid hydrocarbons.
These and other hydrocarbons may be considered as somewhat analogous to the metals ; the analogy appears in the methods of formation and formulas of their derivatives.
CHAPTER III
The Halides of the Elements— Double Halides — Endothermic Combinations — Hydrolysis — Oxidation and Reduction — Mass=Action.
The Halides. — Compounds of fluorine, chlorine, bro- mine, and iodine are thus named. They fall into classes when the elements are arranged according to the periodic system. Taking the chlorides as typical of the halides, we have the following table : —
|
LiCl NaCl |
BeCl2 MgCl2 |
BC13 A1C13 |
CC14 SiCl4 |
NC13 PC15 PC13 SF6 |
sci4' |
OC12 SC12 |
HC1 ... FC1? ... C1C1 |
|
KC1 RbCl CsCl |
CaCl2 SrCl2 BaCl2 |
ScCl3 YC13 LaCl3 |
TiCl4 ZrCl4 CeCl4 ThCl4 |
... AsCl3 ... SbCl5 SbCl3 ErCl3 ... BiCl3 |
SeCl4 TeCl4 |
TeCl2 |
IC13 IC1 |
|
CuCl AgCl |
ZnCl, CdCl2 GdCl, HgCf2 |
GaCl3 InCl3 Tici3 |
GeCl4 SnCl4 TbCl4 PbCl4 |
VC15 VC1, NbCl5NbCl3 ... ... PrCU ... TaCl3 ... WC16 |
MoCl4 WC14 |
CrCl2 MoCl2 NdCl2 WC12 |
MnCl3 ... |
|
FeCl3 Fed., RuCl3 RuCl2 |
CoCls CoCl2 ... RuCl3 ... |
PdCl4 |
NiCl2 PdCU |
OsCl4 OsCl3 OsCl.2 IrCl4 IrCl3 ... PtCl4 ... PtCl2
50
THE HALIDES 51
Besides these compounds, which present considerable regularity, others exist which have less claim to order. Thus, KI3 is also known ; it is unstable, but CsI3 is re- latively stable. Again, CuCl2 and AuClg exist, also HgCJ. In the next group, GaCl9, InCl, and InCl2 are also known, as well as T1C1. The following group contains SnCl., and PbCl2 ; PbCl4 is very unstable. Besides VCl/and VC13, VC14 and VC12 are also known ; and in the next group, CrCl3, MoCl3, and MoCl5, also WC15, UC13, and UC15. These compounds are difficult to classify.
The bromides and iodides, as well as the fluorides, corre- sponding to many of these chlorides in formula, are also known. Where they are of special interest, they will be alluded to in the sequel.
The characteristic of the halides of the elements of the lithium group is that they are all soluble white salts, crystallising in cubes. In dilute solution they are all ion- ised, and even in strong solution a large percentage of ions are present. Hence they all react as metal ions and as halo- gen ions. Thus, for instance, with silver nitrate, which is the usual test for ionic chlorine, the following reaction takes
place : — NaCl. Aq + AgNO3. Aq = NaNOg. Aq + AgCL Practically insoluble, and therefore practically non-ionised, silver chloride is precipitated, and free ions of sodium and the nitrate group remain in solution. If concentrated solutions are mixed, that portion which is ionised reacts ; and as it is removed from solution, the originally non- ionised molecules of sodium chloride are ionised, because the solution becomes more dilute as regards sodium chloride, and they, too, enter into reaction. In a similar way, the alkali metal ions react in presence of a suitable reagent. Another point to be noticed is that these salts are not hydrolysed, that is, do not react with water to give hydroxide and acid to any appreciable extent, and the usual method of preparing them depends on these facts. They may
52 MODERN CHEMISTRY
all be obtained by the addition of halogen acids to the hydroxides or carbonates of the metals dissolved in water,
thus: KOH.Aq + HBr.Aq-KBr.Aq + H2O. It will be noticed that the water is not ionised, nor does it hydro- lyse the potassium bromide ; hence, on evaporation, as concentration increases, the number of ions of potassium and bromine becomes fewer and fewer, and after the water has been removed the pure dry salt is left. With a' carbonate the action is similar. The equation is :
U2CO3.Aq + 2HI.Aq = 2LiL Aq + H2O + CO2. In dilute solution the acid H2CO3 would be liberated ; it is a very weak acid, /'.<?. it is comparatively very slightly ionised into
+
2H.Aq and CO3.Aq ; and, moreover, it readily decom- poses into H2O and CO2 ; hence it is removed from the sphere of action as it is formed, and on evaporation the salt is left behind, as in the previous example.
Sodium and potassium chlorides occur in nature ; the former in the sea, which contains from 3.8 to 3.9 per cent. Deposits, which have undoubtedly been formed by the drying up of inland seas, are found in many places. At Stassfurth in S. Germany there are large deposits of all the salts present in sea- water, including common salt, chlorides and sulphates of magnesium, potassium, and sodium, and calcium sulphate ; these have been deposited in layers in the order of their solubilities, the less soluble salts being deposited first. Bromides and iodides are also present in minute quantity in the residues from the evaporation of sea- water.
Solutions of the halides of the beryllium group of elements can also be made by acting on the hydroxides or carbonates of the metals with the halogen acid. To take barium chlo-
+ +- + - ++-
ride as an example, BaCO3.Aq+ 2HCl.Aq = BaCl2.Aq + H2O + CO2. Now barium carbonate is nearly insol- uble in water, but the portion which dissolves is ionised ; and, as explained above, when the portion which is ionised has
THE HALIDES 53
reacted, its place is taken by more of the carbonate entering into solution ; so that finally all is changed into chloride. With the hydroxides, the same kind of reaction takes place :
Ca(OH)9.Aq + zHBr.Aq == CaBr^Aq + 2H,O. These salts are also white and soluble in water. There is, however, one exception, namely, calcium fluoride, CaF9, which occurs native as fluor- or Derbyshire spar. It forms colourless cubical crystals, and is the chief compound of
+ + - fluorine. It is produced by precipitation: CaCl9.Aq +
q = CaF2H-2KCLAq. The calcium fluoride is non-ionised, and comes down in an insoluble form.
Water of Crystallisation. — The other halides of this group crystallise with water of crystallisation ; its amount varies from 7 molecules, as in BaI9.yH9O, to I as in ZnCl9.H9O. The retention of this so-called " water of crystallisation " has not yet been satisfactorily explained. It was for long believed that such compounds were " mole- cular/' as opposed to atomic ; that is, that the water molecules combined as wholes with the salt, and not by virtue of their atoms ; but it is more probably to be explained by the tetravalency of oxygen, although even with this assumption it is not easy to ascribe satisfactory constitutional formulas in all cases. It must at the same time be assumed that the halogen atoms are of a higher valency than unity ; possibly triad, or even pentad.
These salts are hydrolysed in solution to a small extent ; thus a solution of magnesium chloride, besides containing a large number of ions, has also reacted with the water to form hydroxide and hydrogen chloride: MgCl9 + 2H(OH) = Mg(OH)2 + 2HC1. As the solution becomes con- centrated on evaporation, the hydrogen chloride volati- lises with a part of the water ; and a mixture, or rather a compound, of the oxide and chloride remains. Hence these chlorides cannot be obtained in a pure state by evaporating their solutions. They exhibit another property, however,
54 MODERN CHEMISTRY
which makes it possible to obtain them in a pure state, namely, the power of forming " double halides." This pro- perty is not well marked with the halides of calcium, barium, and strontium, but the halides of beryllium, magnesium, zinc, and cadmium are notable in this respect. We have, for example, MgCl2.KC1.6H2O, ZnCl2.NH4Cl, and many similar bodies. In solution, such compounds are mainly ionised into their simple ions, but on evaporation the non- ionised salt separates in crystals, and is not subject to hydrolysis. Hence such salts can be dried without decom- position. The ammonium salts, when sufficiently heated, lose ammonia and hydrogen chloride by volatilisation, and the anhydrous halide is left]: MgCl2.NH4Cl = MgCl2 + NH3 -f HC1. The mode of combination of these double salts is possibly owing to the fact that the halogens are
capable of acting as triads ; thus Zn/ may be
taken as the constitutional formula of that particular salt.
The mono-halides of copper, silver, and gold may be attached to the first group ; and if that is done, the mono- halides of mercury must also be included. These com- pounds are all insoluble in water, and are consequently obtained by precipitation or by heating the higher halides, where these exist. Thus CuCl and AuCl are obtained by cautiously heating CuCl2 and AuCl3 ; they are white insoluble powders. Cuprous chloride is more easily obtained by removing half the chlorine from cupric chloride dissolved in concentrated hydrochloric acid, by digesting it with metallic copper: CuCl2.2HCl.Aq + zHCl.Aq + Cu = Cu2Cl2.4HCl.Aq, a brown compound, which is decomposed by water into Cu9Cl9 and 4HCl.Aq ; the cuprous chloride is thrown down as a snow-white powder. With silver and mercury, the chlorides AgCl and HgCl are formed by precipitation from the respective nitrates, AgNO3 and HgNO3, on addition of soluble chlorides. The bromides and iodides are similarly formed, and are also insoluble.
THE HALIDES 55
There are several interesting points connected with these halides. First, as regards their colour; the chlorides are white; cuprous bromide is greenish brown, while the brom- ides of silver, gold, and mercury are yellow ; and cuprous iodide is brownish, and the iodides of the other metals darker yellow than the bromides. It appears as if the colour was influenced both by the- metal and by the halogen. Next, the chlorides of copper and mercury give evidence of possessing the double formulae Cu2Cl2 and Hg2Cl2, which would imply that the metals were only pseudo-monads, and that the structural formulas should be Cl— Cu— Cu— Cl and Cl— Hg— Hg— Cl ; and this would correspond with the fact that the chlorides CuCl2 and HgCl2 are also known ; but, on the other hand, as AgCl in the state of gas has the simple formula given to it, it may be that it is the halogen which forms the bond of union between the two half- molecules, thus : CuCl=ClCu. Silver forms no higher halides.
The fluorides of these elements differ from the others in being soluble in water ; they are prepared from the oxides with hydrofluoric acid. They are very difficult to dry, for they undergo the reverse reaction, and are hydrolysed into oxide and hydrogen fluoride on evaporation.
Copper and mercury also function as dyads ; that is, their ions are capable of carrying a double electric charge under certain circumstances. What the mechanism of this change is, we do not know ; but the change in valency can be induced by presenting to the element a larger amount of halogen, if it is desired to increase the valency, or by remov- ing halogen if the opposite change is required. The addition of halogen to the mono-halide is in each case an exothermic change, and its converse is an endothermic one. Cuprous or mercurous chloride, heated in a current of chlorine changes to cupric or mercuric chloride, and the converse change can be brought about by heating the higher halide in a current of hydrogen, or by exposing the lower halide to the action of nascent hydrogen ; but it is difficult to
56 MODERN CHEMISTRY
prevent the action in the latter case from going too far and yielding the metal. A solution of cupric chloride saturated with sulphurous acid in presence of hydrochloric acid, and then diluted with water, gives a precipitate of cuprous chloride : 2CuCl2. Aq + H2SO3. Aq + 2HC1. Aq + H2O = Cu2Cl2.4HCl.Aq + H9SO4.Aq. The sulphurous acid re- moves oxygen from water, liberating hydrogen in presence of the cupric chloride, and the latter is deprived of half its chlorine and reduced to cuprous chloride. Similarly, stannous chloride forms a reducing agent for mercuric chloride : 2HgCl2. Aq + SnCl2. Aq = Hg2Cl2 + SnCl4. Aq. The converse change can be produced by exposing the lower halide in presence of halogen acid to the action of nascent oxygen : Cu2Cl2 + zHCl.Aq + O - iCuCL,. Aq + H2O. This oxygen in the case of copper may be molecular, O9, but for the formation of the higher halide of mercury, it must be derived from some substance capable of parting readily with oxygen, such as nitric acid.
Cupric iodide is very unstable, and readily yields up iodine, forming cuprous iodide. On mixing cupric chloride with potassium iodide, the cuprous iodide is precipitated :
2CuCl2. Aq + 4-KI. Aq - Cu2I2 + 4KCl.Aq + I2. It is to be noticed that the dyad cupric ions have lost two charges, ancl that these have neutralised the two negative charges of the iodine ions, causing them to be precipitated. (Inasmuch as the cuprous iodide is insoluble, it should not have had the ionic signs attached ; but they have been kept in order to show the changed valency. ) Mercuric iodide is an insoluble scarlet precipitate, and is therefore best produced by pre- cipitation. It dissolves, however, in a solution of potassium iodide, forming a double salt, of which more shortly.
Auric chloride contains triad gold, and thus has the formula AuCl3. It is not produced by the direct action of chlorine on gold, because the temperature of attack is above the temperature at which the compound is decomposed. But it is possible to volatilise gold in a current of chlorine,
THE HALIDES 57
because a few molecules escape decomposition and are volatilised along the tube through which the chlorine is passed, and on cooling the gold is deposited, owing to the decomposition of the chloride at a lower temperature. It may appear paradoxical that the chloride is stable at a higher temperature than that at which it decomposes ; but it is to be presumed that the difference of temperature between one favourable to an exothermic and to an endo- thermic action is very small ; and as endothermic substances increase in stability on rise of temperature, the chloride is capable of volatilisation ; on cooling it becomes unstable and undergoes decomposition with deposition of gold. The usual method of preparing this salt is to dissolve gold in a mixture of nitric and hydrochloric acids. This mixture yields ionic chlorine, the negative charge of which neutralises the positive charges of the gold ; but there are corresponding negative charges set free, which are transferred to the ion
NO3 of the nitric acid, converting it into 2O, with its four negative charges. The latter combines with the hydrogen,
+ - + - forming electrically neutral water: 3HC1 + HNO3.Aq +
Au = 2H2O + Au C13 . Aq + NO.
Auric chloride forms dark red crystals ; it is soluble in water, and when mixed with chlorides of the alkali metals forms a set of salts termed aurichlorides. The potassium salt, for example, has the formula K AuCl4 ; it is soluble in water, but, unlike the "double salts/' such as MgCl2,2KCJ, already alluded to, it is ionised by water, not into simple
+ ions like these, but into the ions K and the complex group
AuCl4. At the same time there exists in the solution a small number of simple ions, so that on electrolysis gold is deposited at the kathode, but the primary effect of the current is to send the aurichloric ions to the anode. The solution of mer- curic iodide in potassium iodide, of which mention was madebefore,is a half-way example of the same kind. Its solu-
58 MODERN CHEMISTRY
+ tibn contains ions of K and HgI3, but these are mixed with
+
a much larger proportion of the simple ions, K and I and i- +
Hg and I2. Ail grades of such salts are known ; indeed it is probable that the double salts, such as magnesium-potassium
chloride, contain a small number of complex ions of MgClg.
These halides have been considered at length because they form types of the others. Use will be made of the examples given in treating of the remaining halides.
We have seen that the halides may undergo either ionisa- tion or hydrolysis, or both at once. The ionisation may be more or less complete, and the hydrolysis is promoted by dilution and by a high temperature. The remaining halides display both these kinds of behaviour, and according as one or the other prevails, the methods of preparing them are affected. In certain cases, moreover, the halides form compounds with other halides, usually those of the alkali metals or hydrogen, which are less apt to be hydrolysed, and yield different complex ions. The halides of carbon and nitrogen belong to neither of these classes, for they are insoluble in and unacted on by water. As neither carbon nor nitrogen is acted on by the halogens (excepting that carbon burns in fluorine), they must be prepared indirectly by acting on one of their compounds with the halogen. Methane or carbon disulphide is chosen for the former, and ammonia in preparing the latter. By passing a current of chlorine saturated with the vapour of carbon disulphide through a red-hot tube, the chlorides of both carbon and sulphur are formed: CS2 + 3C12 = CC14 + S2C12. On treatment with water the sulphur chloride is decomposed, while the chloride of carbon may be distilled off ; it forms a colourless liquid boiling at 76.7°. Its smell resembles that of the closely allied chloroform, CHCJ3, and it is also possessed of anaesthetic properties. For the preparation of nitrogen chloride a jar of chlorine is inverted over a
THE HALIDES 59
saturated solution of ammonia in water ; oily drops are formed which settle to the bottom of the vessel : NH3. Aq + 3C12 = NC13 + 3HCl.Aq ; the HC1 unites with ammonia, forming ammonium chloride.
Endothermic Combination. — This body is fearfully explosive, for its formation is attended by great absorption of heat ; but during its formation the reagents do not grow cold ; for the formation of ammonium chloride is a highly exothermic reaction, and the amount of heat evolved by its formation is greater than that of the equivalent amount of chloride of nitrogen ; hence the change as a whole is accompanied by evolution of heat. It is thus that endo- thermic compounds are usually formed : by virtue of a simultaneous action in which heat is evolved. The slightest shock causes the decomposition of such endothermic bodies ; if one single molecule is decomposed, it evolves heat and brings about the decomposition of its neighbours ; and as all the molecules are in close proximity to each other, and as the products, nitrogen and chlorine, are both gases, and are, moreover, much raised in temperature by being set free, the decomposition is accompanied by sudden and enormous ex- pansion. Nitrogen iodide, prepared by adding a solution of iodine to aqueous ammonia, is a black solid of the formula NI3.NH3 ; it is also explosive.
The fluorides of boron and silicon are both produced by the action of a strong solution of hydrofluoric acid on the oxides ; but it is necessary to have some agent present to withdraw water, such as concentrated sulphuric acid. These compounds are both gaseous. Their formation is shown by the equations: B2O3 + 6HF = 2BF3 + 3H2O ; SiO2 + 4HF = SiF4+ 2H3O. If the water is not with- drawn, combination ensues between the fluoride and hydrogen fluoride, with formation of HBF4 or H9SiF6, named re- spectively hydroborofluoric, and hydrosilicifluoric acids thus:
2H2SiF6 + H2SiO3. These compounds ionise into H-
60 MODERN CHEMISTRY
ions, and the complex ions BF4 and SiFfi ; and many salts are known in which metals replace the hydrogen. They are similar in kind to potassium aurichloride.
The other halides of boron and silicon, and also of phosphorus, sulphur, selenium, tellurium, and iodine, react at once with water, forming hydrogen halide and an acid. The equations are as follows : —
BCl3+3H9O.Aq = B(OH)3.Aq+3HCl.Aq;
SiCl4+3H;O.Aq - O=Si(OH)9 + 4HCJ.Aq;
PCJ3+3H9O.Aq = P(OH)3.Aq+3HCl.Aq;
4H;O.Aq = 0=P(OH)3.Aq+5HCl.Aq; 2 + 2H0O.Aq = O=S(OH)2.Aq
"
9O.Aq:= O=Te(OH)2.Aq + 4HCl.Aq
+ Te; + 3H9O.Aq = O=S(OH)2.Aq + 4HCl.Aq;
H2O.Aq = 3HIO3.Aq+i5HCl.Aq + I2.
It is to be noticed that where a hydroxy-compound corresponding to the halide is capable of existence, it is formed ; if not, excess of the element is set free. Hence none of these halides can be prepared by acting on the hydroxide with a halogen acid ; they are all made either by the direct action of the halogen on the element, or by what comes to the same thing, the action of the halogen on a strongly heated mixture of the oxide of the element with carbon. Boron, silicon, and phosphorous chlorides are vola- tile liquids ; they fume in the air owing to their action on the water-vapour. S9C12 is a yellow liquid ; when saturated with chlorine at a low temperature, SC19 and SC14 are successively formed ; but on rise of temperature they dis- sociate into the lower chloride. IC1 is a black solid, converted by excess of chlorine at a low temperature into IC13, a yellow solid, which easily dissociates into IC1 and C19 ; and PC15 is a pale yellow solid, volatile at a high temperature in a perfectly dry atmosphere without
THE HALIDES 61
dissociation, but resolved by the least trace of moisture into PC13 and C12.
Valency of Elements. — We may remark here the gradual increase of valency as we pass from left to right in the periodic table. Lithium is a monad, with its congeners ; the elements of the beryllium group are dyads ; boron a triad ; carbon a tetrad ; phosphorus acts as pentad as well as triad ; sulphur, as a pseudo-monad, a dyad, and a tetrad ; and Moissan has lately shown that sulphur burns in fluorine, forming a very stable hexafluoride, SFC ; while iodine forms a monochloride and a trichloride, and probably also a pentafluoride and a heptafluoride.
Passing back to the boron group, if it is desired to form anhydrous chloride, it is necessary either to heat the element, or its oxide mixed with charcoal, in a current of chlorine, or, except in the case of boron, to prepare a double salt of the chloride with ammonium chloride, and to volatilise the latter after driving off the water ; the aqueous chlorides are formed by dissolving the oxides or hydroxides in hydrochloric acid. Thallium forms monohalides, sparingly soluble in cold water, and thereby attaches itself to the copper group.
Almost the same remarks apply to the elements of the carbon group ; solutions of the chlorides, with exception of those of carbon and silicon, are obtained from the element and hydrochloric acid or from the hydroxide, and they cannot be dried without reacting wholly or partially with water. For instance, titanium chloride, on careful addition of water, can become ClTi(OH)3, Cl2Ti(OH)2, ClgTi(OH), all of which are intermediate products between the tetrachloride and the tetrahydroxide ; such compounds are termed "basic chlorides." Anhydrous stannic chloride is a fuming liquid, formed by the distillation of a mixture of the metal with mercuric chloride or by heating the metal in a stream of chlorine. Lead tetrachloride is a very unstable liquid, formed from the tetracetate, Pb(C2H3O2)4, by converting it into the double ammonium salt with a
62 MODERN CHEMISTRY
mixture of ammonium chloride and concentrated hydro- chloric acid; this salt, (NH4)t>PbC]r>, is then decomposed by concentrated sulphuric acid, when the tetrachloride separates as a heavy liquid. It at once decomposes into PbCl2 + C10 on warming ; hence PbO2, when warmed with hydrochloric acid, undergoes the change: PbO,-t- 4HC1. Aq - PbCl, + Aq + C12.
Tin and lead resemble elements of the zinc group in forming dichlorides. On dissolving tin in hydrochloric acid the dichloride is formed ; and a solution of the tetra- chloride, when exposed to the action of nascent hydrogen, yields the lower chloride. This action may be thus for- mulated :— S n C14. Aq + 2H = SnCl2. Aq + HC1. Aq. Stannous chloride is a white, soluble salt, crystallising with water of crystallisation. Lead dichloride, on the other hand, is sparingly soluble in cold water ; it is formed when a soluble lead salt, such as the nitrate, is mixed with the solution of a chloride: Pb(NO3)2.Aq+ 2NaCl.Aq = PbCl2 + 2NaCl. Aq. The bromide and the iodide are also sparingly soluble, and are similarly produced.
With arsenic and the remaining members of that group we may notice the same characters : the anhydrous chlorides produced by the action of chlorine on the element, or, when it is not available, on a mixture of the oxide with carbon at a red heat ; the aqueous solution produced by dissolving the oxide or hydroxide in hydrochloric acid., Basic chlorides are also known, e.g. ClAsO, CISbO, and ClBiO, from the trichlorides ; and OPC13, and OSbCl3,, from the pentachlorides, on reacting with a small amount of water.
Mass=Action. — The action of mass, that is, the quan- tity of a compound in unit volume, is well illustrated by the action of water on antimonious chloride. A solution of this salt in hydrochloric acid gives a precipitate on adding water: SbCl3.nHCl.Aq + H2O = OSbCl + (n + 2)HCl.Aq. Here the increase in the number of molecules of water
THE HALIDES 63
causes the precipitation of the basic chloride ; on adding more hydrochloric acid, however, so as to increase its active mass, the reaction is reversed, and the precipitate re- dissolves : OSbCl+(n + 2)HCl.Aq = SbCl3.Aq.nHCl-h Ht,O. Above a certain concentration of water SbOCl is stable ; above a certain concentration of hydrogen chloride, SbCl3.
The higher halides of molybdenum, tungsten, and uranium, themselves prepared by the action of halogen on the element, yield tue lower halides on heating. They are volatile, coloured bodies, soluble in water ; the higher ones are decomposed by water.
The elements chromium, manganese, iron, cobalt, and nickel, although not all belonging to the same series, show, nevertheless, a gradation of properties. The dihalides of all are known in the dry state ; they are most readily obtained by heating the metal in a current of hydrogen halide, if required anhydrous ; or if in solution or crystal- lised with water, by dissolving the oxide or carbonate in the halogen acid and evaporating until crystallisation ensues. As examples: Fe + 2HC1 = FeCL + H.,0 ; MnCO -f 2HBr. Aq - MnBr2. Aq + H2O + CO2.
The trihalides are best made by heating the elements, in a current of halogen, if required anhydrous ; if in solu- tion, by dissolving the oxide or hydroxide in the halogen acid. The trihalides of manganese and cobalt are very unstable ; and if the corresponding oxides be treated with halogen acid, a portion of the halogen is evolved, thus : Fe9O3 + 6HCl.Aq - 2FeCl3.Aq + 3HQO ; Mn.,O8 +- 6HC1. Aq - 2MnCl3Aq + sH2O. But MnCirAq gradu- ally decomposes, especially if temperature is raised, thus : 2MnCl3.Aq=2MnCl2.Aq + Cl2. And if MnO2 be em- ployed, chlorine is evolved from the outset : 2MnO2 -f- 8HCLAq=2MnCl3.Aq + 4H2O + CJ2; the MnCl3 de- composing further on standing or on rise of temperature. With Co2O3 a transient brown coloration is noticeable on. adding hydrochloric acid, implying the momentary forma-
64 MODERN CHEMISTRY
tion of CoCl8.Aq ; but it is at once resolved into CoCl.2.Aq and free chlorine.
Oxidation and Reduction. — As already remarked, the raising of the valency of an element is often spoken of as " oxidation ; " the reducing of the valency, as " reduc- tion." The tendency of chromous halides to transform into chromic compounds is so great, that it is not possible to expose them to air without the change taking place, and consequently the reduction of chromic compounds to chromous is a difficult operation. But with iron, both classes of compounds have nearly equal stability ; hence oxidation and reduction play a great part in their formation. The action of nascent hydrogen from any source reduces ferric halide into ferrous: FeCl3.Aq + H = FeCl9.Aq + HC1. Aq. Similarly, a ferrous halide, in presence of halogen acid and either free or nascent oxygen, is oxidised to a ferric : aFeCl,. Aq + 2HC1. Aq + O = 2FeCl3. Aq + H,O. Or the halogen itself may be used to effect the change : 2FeCl2.Aq + Q2 = 2FeCJ3.Aq. On evaporating these solutions, hydrolysis takes place partially ; thus ferric chloride yields compounds of a basic character, such as (OH)FeCl2, (OH)2FeCl, which are partly hydroxide, partly chloride. This statement applies to the halides of all these metals.
Colour of Ions. — The triad and dyad ions in the case of these metals exhibit remarkable differences of colour. Thus chromous ions are blue, chromic, green ; basic ferric ions are orange-yellow, ferrous, pale green ; manganic, brown, manganous, pale pink ; cobaltous, red, and nickelous, grass- green. Hence a change in the ionic charge of the metallic ion is accompanied by a striking colour-change.
The halides of the palladium and platinum groups of metals closely resemble in character those of gold, which have already been described. The dihalides of the palladium group are all soluble, save PdI0, which is pre- pared by precipitation with potassium iodide. Nitro-hydro- chloric acid yields the higher chloride ; it remains on evapo-
HALIDES OF COMPLEX GROUPS 65
ration. These form with chlorides of the alkalies double salts, e.g. RuCl3.2HCl, RhCl3.2HCl, and PdCl4.2HCl;
the latter are probably ionised as KK and PdCl6, &c. Chlorine also acts directly on red-hot metals of the platinum group, forming a mixture of chlorides ; these, on heating, lose chlorine, giving lower chlorides. Solutions of the halides can also be prepared by the action of the halogen acid on the respective oxides. On heating to a high temperature, all these halides are decomposed into the metal and halogen. The compounds K2PtCl6 and (NH4)2PtCl6 require special mention ; they are orange salts, nearly insoluble in water, and are used as tests for potassium and ammonium, and also as a precipitant in estimating these ions. Their existence is probably to be ascribed to the power possessed by chlorine of sometimes acting as a triad, and the structural formula is be- KCi=Cl\ /Cl NH4C1=C1\ /Cl
lievedtobe >Pt< and >Pt< .
KC1= Cr XC1 NH4Cl=C2i/ XC1
Halides of certain complex groups are also known. When these contain oxygen or hydroxyl, (OH), they are generally termed basic salts or halo-acids ; they will be considered later. The others may be divided into two classes : those like ammonium halides, and those derived from hydrocarbons.
Ammonium and phosphonium halides. — These hal- ides, which are formed by direct addition of the hydrogen halide to ammonia or to phosphine, closely resemble in colour, in crystalline form (cubic), and in reactions, the halides of the lithium group of metals. On mixing a solution of ammonia and hydrochloric acid, for example, the combina- tion occurs: NH3.Aq + HCl.Aq==NH4Cl.Aq ; and on evaporating the solution to dryness, ammonium chloride is left in an anhydrous state. From the conductivity of ammonia solution, it is known to contain a certain amount of NH4OH in an ionised condition ; and the equation may be
written: NH4OH.Aq + HCLAq = NH4Cl.Aq + H2O.
VOL. II. E
66 MODERN CHEMISTRY
As the hydroxyl ion is removed from the solution by the for- mation of practically non-ionised water, more and more am- monium hydroxide is formed to maintain equilibrium between the NH3. Aq and the NH4OH. Aq ; and the whole is ulti- mately transformed. The rate of transformation, however, is a very rapid one. Combination has been shown not to take place between perfectly dry ammonia and dry hydrogen chloride ; hence it does not seem unlikely that ionisation may occur, either in the gaseous state, or more probably on the surface of the vessel in the condensed layer of moisture which appears always to adhere to all solid surfaces. Once started, combination occurs continuously until the reaction is complete. Ordinarily "dry" ammonia, however, at once gives a dense cloud with hydrogen chloride, bromide, or iodide. Again, perfectly dry ammonium chloride has the vapour-density 26.25, corresponding to the molecular weight (N= I4 + H4 = 4 + C1 = 35.5) = 53.5 ; whereas, if moist, the density is half that amount, corresponding with a mixture of NH3 = 1 7 and HC1 = 36.5. These compounds have densities of 8.5 and 18.25 respectively, and a mix- ture in equal proportions of each has a density the mean of the two, viz., 13.125. It appears necessary that ionisation
+
into NH4 and Cl should take place before dissociation into NH3 and HC1 is possible. The electrically neutral body
+
NH4C1 can volatilise unchanged ; the ions NH4 and Cl are incapable of volatilisation as such, and in volatilising unite their opposite charges, and form the two electrically neutral compounds HC1 and NHg.
PhospMne, PH3, also unites with hydrogen chloride, but only under high pressure, at the ordinary temperature. On the other hand, phosphonium iodide, PH4I, is pro- duced by the union of phosphine with hydrogen iodide under atmospheric pressure ; it forms white, cubical crystals, which, like ammonium chloride, dissociate when heated. The hydrides of arsenic and antimony form no such compounds.
PHOSPHINE GROUPS 67
It must be assumed that these compounds are formed with change of valency of the nitrogen or phosphorus ; the triad becomes pentad ; the NmH3 becomes H4NVC1. On distilling with sodium hydroxide or slaked lime, water is formed, and the element is reduced to its original triad con- dition, thus: NH4Cl.Aq + NaOH.Aq = NH4OH.Aq +
NaCLAq, and NH4OH.Aq = NH3 + H2O.Aq, two electrically neutral bodies.
Carbon shows no such tendency to change valency. The hydrocarbons of the methane series are " saturated/' i.e. they have no tendency to take up any other element. Hence halogen must replace hydrogen. This can be done either directly, by the action of the halogen on the hydro- carbon, as, for instance, CH4 + C12 = CH3C1 + HC1 ; or indirectly, by the action of the halogen acid on the hydr- oxide : CH3OH + HC1 = CH3C1 + H2O. Such hydr- oxides are termed alcohols ; that derived from ethane, C2H6, is the ordinary anhydrous alcohol of commerce ; its formula is C9H5OH, and the corresponding chloride of ethyl is C0H,C1. It will be remembered that the struc-
H\ /H
tural formula of ethane is H-^C — QJ-H, and that of H/ ^H
H\ /H ethyl chloride is H-^C— C^-C1. There is, however, a
H/ \H
difference between the formation of ethyl chloride, for example, and of sodium chloride. Whereas sodium chloride is ionised in solution in water, ethyl chloride is insoluble, and is therefore non-ionised. Hence the action is a slow one ; the alcohol is saturated with hydrogen chloride, allowed to stand for some hours, and distilled ; ethyl chloride, being volatile, passes over ; it is a gas, condensing at about 12° to a mobile colourless liquid. It is probable that the hydrogen chloride is ionised in solution in alcohol ; the alcohol is also possibly ionised to a minute extent ;
68 MODERN CHEMISTRY
water is formed by the union of the hydrogen and hydroxyl
+ ions, and non-ionised ethyl chloride distils over : C0H.OH
+ HCl.Alc = H2O + C,H5Cl. But this suggestion, it must be admitted, is somewhat speculative, and is based only on analogy with reactions of a more familiar nature.
The formation of some of the halogen compounds of the olefines, and of hydrocarbons of the acetylene and benzene series, has already been alluded to on p. 48.
CHAPTER IV
Hydroxides and Acids — "Insoluble Substances" — Indicators— Preparation of Basic Oxides— Pro- perties of the Basic Oxides and Hydroxides — Sulphides — The ' ' Solubility-Product ' ' — Basic Oxides and Hydroxides of Complex Groups: Alcohols, Aldehydes, Ethers; and Sulphines, Amines and Phosphines.
The Oxides and Hydroxides, Sulphides and Hydrosulphides, Selenides and Tellurides. —
Owing to the dyad valency of oxygen, sulphur, selenium, and tellurium, compounds of these elements are more numerous than those of the halogens. And whereas double halides of hydrogen and other elements are not numerous, being confined to such bodies as H2SiF6, HBF4, H.7PtCl6, and a few others, those of the oxides are very numerous, and form two important classes, the " hydroxides " and the "acids."
Hydroxides and Acids. — Members of both these classes may be regarded as hydroyxl, that is, water minus one atom of hydrogen, -OH, in combination with elements, but they differ radically in that the true hydroxides ionise
+ - ++ -
into element and hydroxl, thus: NaOH.Aq, Ca(OH)9.Aq, + + + - Bi (OH)3; whereas acids ionise into hydrogen and a
+ • + -
negatively charged radical, thus : HOCl.Aq, HNO3.Aq, + - - + -
H(HSOJ.Aq, H0SO4.Aq, and many others. There 69
70 MODERN CHEMISTRY
are certain hydroxides in which the ionisation may take either form ; such compounds are said to be either "basic " or " acid " according to circumstances ; thus, aluminium
hydroxide, A1(OH)3, is basic; with hydrochloric acid
+ + + - it reacts in the following manner: Al(OH)3.Aq +
3HCl.Aq = Al Cl3.Aq + 3H2O; on the other hand, when caustic soda is presented to aluminium hydroxide, it forms sodium aluminate, NaAlO2.Aq, a derivative of the acid HAlOi;.Aq, which is formed from A1(OH)3 by loss of water : "A1(OH)3 = O=A1OH + H0O. The ions in the
+ latter case are H and A1O9, and the reaction takes place
between HA1O2 and NaOH, thus: HAlO2.Aq +
NaOH.Aq = NaAl62.Aq + H2O. It is generally the case that the acids are derived from hydroxides which have lost a portion of their hydrogen as water. They are, like O=A1OH, partly oxide, partly hydroxide.
"Insoluble Substances." — The hydroxides are, with some exceptions, generally spoken of as insoluble in water. The word " soluble " is a relative term ; it is probable that very few, if any, substances are absolutely insoluble. Silver chloride is usually regarded as wholly insoluble in water, but pure water shaken up with that salt acquires increased conductivity, showing that some chloride must have gone into solution. In one of the equations given above, A1(OH)3 is followed by "Aq," implying that it is dissolved and ionised in solution. This method of writing is perfectly correct for the portion which is dis- solved, but that constitutes only a very minute fraction of the whole. What is dissolved, however, is ionised and enters into reaction, and when it has been removed, as in the equation given, with formation of practically non-ionised water-molecules, its place is taken by more : equilibrium tends to establish itself between the dissolved portion and
HYDROXIDES AND ACIDS 71
the portion remaining undissolved. We know well that if excess of common salt be placed at the bottom of a vessel of water it will not all dissolve, but, as the dissolved portion diffuses away into the upper layers of water, its place is taken by fresh salt, which dissolves, until, if sufficient time be given, the whole solution becomes saturated with salt. Similarly, the removal of the aluminium — as ions, it may be — and of the hydroxyl of the aluminium hydroxide, A1(OH)3, as water, on treatment with an acid, causes a fresh portion of the hydroxide to go into solution, and this continues to go on until all has undergone reaction.
The hydroxides of the elements may be classified like the halides ; the analogy between the formulae is seen on comparing the tables on pp. 72, 73, with those on p. 50.
Oxygen compounds of fluorine are wanting.
[C1(OH)7], [C1(OH)5], [C1(OH)3], Cl(OH).
[OC1(OH)5], [OC1(OH)3], OCl(OH), C12O.
[02C1(OH)3]> 02C1(OH), [C1203].
03C1(OH), [C1205]. [C1207].
The formulae enclosed in brackets are of unknown sub- stances. The whole scheme is given in order to show the gradual loss of water of an ideal heptoxide.
The compounds I(OH)6(ONa), OI(OH)5, O2I(OAg)3, and O3I(OAg) are known, corresponding to the theoretical perhalic acids. Those corresponding to the halic acids are O2Br ( OH ) bromic acid, and I ( OH ) 5 and O2I (OH ) , iodic acids. Br(ONa) and I(ONa), named respectively hypo- bromite and hypoiodite of sodium, are also known.
It may be noticed that the formulae of some of the com- pounds of chromium are analogous to those of sulphur and of molybdenum ; other compounds, on the contrary, show more resemblance to those of iron. While manganese, like chromium, also shows analogy with iron, it too forms O2Mn(OK)2, like O2S(OH)2 or O 2Mo( OH )2, termed potassium manganate, as the others are hydrogen sulphate and hydrogen molybdate ; and also MnO0, analogous to
MODERN CHEMISTRY
0
o
:0
:0 ..IS?:
^ ^O
ox o o"oT
UO ^c OU ""'
ffl
:0
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O tn O 'Oj2O 'O 'O ' 'OO ^•^ in' "o'c/^ r? TT v? ~^r.£* W W fflffl
^-q Sr^o od'< gdffi1
:O
P9
OH
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:O w
OH
fflOCQ
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bjo b/) SS
:§0
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UU WC/5 CQffi
lu"
OXIDES AND HYDROXIDES 73
9, °? Q
I-IH HH , -*
c ^ : : : : *
O* O 9 § o* O
u u ^ ^ £> D
1 . i „ ^ 1 1 . i ' I
^ 9^o ^ ^ o 5- ^ duo^ oo^oo
as * « •: «
o 9. » d5 c? ' c? 6s
>* > 9 j? ^* rt cT
O O > Z X H H
£
1 1
O OH
I : o : o o:
O ^ « OH
£ : : £
o o o o o • o
C C *O T3 T3 bO
N N U O O X
° : : O q,
3 bO tuO
U < <
74 MODERN CHEMISTRY
SO2 and MoO2 ; but manganese also forms O3Mn(OK), termed potassium permanganate, which is analogous in formula as well as in crystalline form with potassium perchlorate, O3C1(OK). It is convenient, however, also to include chromium and manganese in the iron group of elements.
Cr(OH)3 Mn(OH)8 Fe(OH)3 Ni(OH)3 Co(OH)3
OCr(OH) OMn(OH) OFe(OH)
Cr2O3 Mn2O3 Fe2O3 Ni2O3 Co2O3
Cr(OH)2 Mn(OH)2 Fe(OH)o Ni(OH)2 Co(OH)2
CrO MnO FeO NiO CoO
Elements of the palladium group have a very wide range of valency ; hence they form a large group of compounds.
OsO4 IrO3
Rh03 02Ru(OK)2 ... 02Os(OH)2 CX>Ir(OK)2
Rh(OH)4 Ru(OH)4 Pd(OH)4 Os(OH)4 Ir(OH)4 Pt(OH)4
Rh02 Ru02 Pd02 Os02 Ir02 PtO2 Rh(OH)3 Ru(OH)3
Rh0O3 Ru2O3 ... OsoO3 Ir2O3
Pd(OH)2 ... ... Pt(OH)2
RhO RuO PdO OsO IrO PtO
The hydroxides of lithium, sodium, potassium, rubidium, and caesium are all soluble white compounds, melting to colourless liquids at a red heat. They do not lose water, even at the highest temperatures, hence the oxides cannot be prepared from them ; indeed, the oxides are produced only by the action of the metal on the hydr- oxide, at a high temperature ; for instance, 2NaOH + 2Na = Na2O + H2. They are white, non-crystalline sub- stances, combining at once with water to form the hydr- oxides Na2O + H2O = 2NaOH. The hydroxides are prepared from the carbonates by boiling a solution with slaked lime (calcium hydroxide) : Na2CO3.Aq + Ca(OH)2.Aq = 2NaOH.Aq + CaCO3; or by heating to redness a mixture of the carbonate with ferric oxide, when the ferrite is formed : K2CO3 + Fe2O3 - 2KFeO2 + CO2. On treatment with water, potassium ferrite is decomposed, thus: KFeO2 + 2H2O.Aq = KOH.Aq + Fe(OH)3.
INDICATORS 75
In either case, the solution of hydroxide is evaporated to dryness in an iron vessel and fused.
These hydroxides are said to be basic, for they are neutralised by acids, forming salts. Thus, with hydro- chloric acid, KOH.Aq + HCl.Aq = KCl.Aq + H20, the point of neutralisation — that is, when the acid and base are present in theoretical quantity to form the salt and water — is determined by the addition of an " indicator."
Indicators. — The most important indicators are litmus, phenol-phthalein, and methyl-orange. Litmus is, a weak acid, red in colour, the salts of which are blue. When dissolved in water, the molecule is hardly at all ionised, hence the red colour of the acid is alone visible. If a base such as sodium hydroxide is added, which, in aqueous solu-
+
tion, is largely ionised into Na.Aq and OH.Aq, the hydroxyl ions combine with that portion of the hydrogen ions of the litmus acid which exist in solution ; when these are withdrawn, more hydrogen ions take their place, and the solution acquires the colour of the ion of the litmus acid, viz., blue. Conversely, if an acid be added to a base in which the blue litmus ions are present, the hydrogen ions of the acid combine with the hydroxyl ions of the base, forming water, so long as any are present ; after they are all in combination they convert the ion of the litmus acid into the red acid, non-ionised, and there is a marked colour-change. As the colours of the litmus acid and of its ion are both very bright, the presence of a mere trace of the indicator suffices. Phenol-phthalein, like litmus, is also a weak acid, that is, it is hardly ionised at all in dilute solution ; the acid is colourless, but the ions are pink, hence the addition of a trace of free alkali causes the colourless solution to become pink. But this indicator gives results only with strong bases, like the hydroxides of the alkalies ; with ammonium hydr- oxide, present in a solution of ammonia in water, it is not a good indicator, for NH4OH is too weak a base, i.e. the hydroxyl and ammonium ions are present in too small amount
76 MODERN CHEMISTRY
to liberate the ions of the phenol-phthalei'n, unless much ammonium hydroxide is present in solution. Hence the presence of a trace of free ammonium hydroxide is not revealed by that indicator. Phenol-phthalei'n is therefore serviceable only with strong bases, but it may be used for weak acids. Methyl-orange, on the other hand, is a com- paratively strong acid ; with a weak base it forms the ions of a salt, and it may therefore be used for weak bases like ammonium hydroxide, or for strong bases like the hydrox- ides of the alkali metals ; but it is too strong an acid to serve well as an indicator of excess of a weak acid, such as carbonic or acetic acid. Its colour-change is from orange to orange-pink.
Preparation of Basic Oxides. — The hydroxides of the metals of the sodium group, as already mentioned, do not lose water on heating, and the oxides, therefore, cannot be thus obtained. Neither do their carbonates lose carbon dioxide, nor their nitrates oxides of nitrogen, save at an im- practicable temperature. But all other basic oxides may be prepared by heating the hydroxides, carbonates, or nitrates of the metals, and a few may be obtained by heating the sulphates. Calcium and strontium oxides are generally prepared from the carbonates, which are found as minerals, named limestone and strontianite respectively. The opera- tion of preparing " quicklime " or calcium oxide is techni- cally, but wrongly, called " burning." Alternate layers of lime and coal are placed in a tower of brick or stone, termed a limekiln ; the coal is set on fire, and its heat expels the carbon dioxide from the carbonate : CaCO3 = CaO+CO9. If calcium carbonate be heated in a closed vessel, however, so that the carbon dioxide does not escape, the dissociation proceeds until the amount of carbon dioxide in the vessel has reached a certain proportion, which is per- fectly definite for each temperature, or until the carbon dioxide has attained a certain " concentration." The reaction then stops. But if the carbon dioxide be removed as it is formed, the reaction goes on to the end, until all carbon
BASIC OXIDES 77
dioxide has escaped. The draught in the kiln removes the carbon dioxide, hence the product is calcium oxide. Stron- tium carbonate is causticised in the same way as limestone ; but the temperature for witherite (BaCO3) is inconveniently high ; baryta is consequently prepared by heating the nitrate, Ba(NOg),, ; it may be supposed to split into BaO and N.,O5 ; the latter, however, decomposes even at moderate temperatures into NO2 and O ; hence the equation is : 2Ba(NO3),= 2BaO + 4NO2 + O2. These oxides are whitish-grey solids, volatile at the temperature of the elec- tric arc, and combining with water with great rise of tem- perature to form hydroxides. The hydroxides are sol- uble in water — barium hydroxide most, calcium least. An aqueous solution of the former deposits crystals of a hydrate, Ba(OH)2.8H2O.
The sparing solubility of calcium hydroxide makes it possible to precipitate it by the addition of caustic alkali to a soluble salt of calcium, provided too much water is not pre- sent : CaCl2.Aq + 2NaOH.Aq = Ca (OH), + aNaCl.Aq. Of course, a saturated solution of calcium hydroxide remains, hence the precipitation is not complete. This plan is appli- cable to the preparation of all hydroxides which are insoluble in water, unless they dissolve in excess of the caustic alkali ; if they do, they are said to display " acid " properties. Beryllium and magnesium hydroxides are thus precipitated : MgCl2.Aq + 2KOH.Aq = Mg(OH)2+2KCl.Aq. The hydroxide may be filtered off and dried, and the white mass, on ignition, leaves the oxide as a white powder. By this means, too, the hydroxides of copper (cupric, Cu(OH)2), silver, AgOH, zinc, Zn(OH)2, cadmium, Cd(OH),, aluminium, A1(OH)3, scandium, yttrium, lanthanum. and ytterbium, gallium, indium, and thallium, with similar formulae, titanium, zirconium, thorium, with formulas OM(OH)2, (where M stands for an element of that group; germanium, tin (stannous, Sn(OH)2, and stannic, SnO2.nH,O), lead (plumbous, Pb(OH)2), bis- muthous, Bi(OH)3, chromic, Cr(OH)3, and chfomous,
78 MODERN CHEMISTRY
Cr(OH)2, manganic and manganous, ferric and ferrous, cobaltous and nickelous : in short, from all elements which form " basic " hydroxides. And from almost all these the oxides may be obtained by heating the hydrates to redness. Excess of the precipitant, however, must be avoided in many cases, for some of these hydroxides display acid properties if in presence of excess of alkali. Thus, for example, if excess of sodium hydroxide or potassium hydroxide be added to the solution of a soluble salt of zinc, such as the chloride, nitrate, or sulphate, the first change, as already shown, is the precipitation of the hydroxide ; but on additi :>n of excess of alkali, the precipitate redissolves, forming the
+ + compound Zn(OK)9.Aq, of which the ions are K, K, and
ZnO9 ; this compound is thrown down by alcohol, in which it is insoluble. It is generally named zincate of potassium. Cadmium forms a similar compound, but that of aluminium has the formula OAl(OK) ; the hydroxide, A1(OH)3, on losing water is transformed into the condensed hydroxide, OAl(OH), which may be termed aluminic acid, of which the hydrogen atom is exchangeable for metals. Stannous and plumbous hydroxides dissolve in excess of alkali, doubt- less forming compounds similar to that of zinc ; and chromic hydroxide is soluble in cold solution of caustic alkali, form- ing, no doubt, a compound analogous to that of aluminium ; but it is decomposed on warming, with reprecipitation of the hydroxide. The hydroxides of all these elements may also be precipitated by a solution of ammonium hydroxide, and some of them are redissolved ; but the compounds formed are of a different nature from those described in the case of zinc, &c., and will be afterwards considered.
Properties of the Hydroxides. — As regards the properties of the hydroxides, that of copper (cupric) is light blue, and of silver, brown ; chromic hydroxide is grey- green, and chromous, yellowish ; manganic, brown, and manganous, very pale pink ; ferric, rust brown, and ferrous, white when pure, but usually dirty green ; cobaltous, dingy
PROPERTIES OF HYDROXIDES 79
red, and nickelous, grass-green. The others are all white amorphous bodies, and they all yield oxides on heating. Cupric oxide is black ; even when boiled with water the hydroxide loses water and changes colour. Argentous oxide is brown ; when heated to redness it loses oxygen, leaving a residue of metallic silver. Zinc oxide is yellow when hot and white when cold ; cadmium oxide is a brown powder ; the oxides of aluminium, scandium, yttrium, lanthanum, ytterbium, gallium, and indium, and of titanium, zirconium, thorium, germanium, and stannic oxide are white powders ; thallium oxide is a yellow powder ; tin monoxide is a black powder ; that of lead (litharge, massi- cot) is a yellow substance, fusible at a red heat ; bismuth sesquioxide is a yellowish powder ; chromic, ferric, and manganic are respectively green, rust-red, and brown ; chromous oxide is unknown, for any attempt to dry it results in the decomposition of water, the absorption of its oxygen by the chromous oxide which becomes chromic oxide, and the evolution of hydrogen. Ferrous hydroxide can be dried, but only with rigid exclusion of air ; it is a black powder. Manganous oxide is greyish-green, cobaltous, olive-green, and nickelous also greyish-green. Manganous hydroxide must also be dried in absence of air.
These hydroxides and oxides are named bases. There are some basic oxides, which are precipitated by adding a hydroxide, such as that of sodium, to a soluble salt, and to which there is no corresponding hydroxide. This is the case with the oxides of mercury. On referring to the table of halides on p. 50, it will be seen that the chloride of mercury has the formula HgCl9. This compound, commonly called corrosive sublimate, when treated in solu- tion with sodium hydroxide, gives a precipitate, not of hydroxide, as might be expected, but of oxide : HgCl0. Aq + 2NaOH.Aq = HgO + 2NaCl.Aq + H2O. Similarly, a soluble mercurous salt, such as mercurous nitrate, Hg0(NO3)0, on treatment with an alkali gives a precipitate
8o MODERN CHEMISTRY
of mercurous oxide: Hg2(NO3)0.Aq + 2NaOH.Aq = Hg2O + 2NaNO3.Aq + H26. These are cases of relative stability ; for, as has been already remarked, on boiling a solution from which cupric hydroxide has been precipitated, the blue hydroxide is changed into black oxide ; other hydroxides lose their water at a still higher temperature ; while those of the alkaline metals may be volatilised without decomposing.
Oxides produced by Heating Carbonates.— Most of the basic oxides may also be prepared by heating the carbonates, a class of salts afterwards to be discussed. The carbonates of the alkali metals, however, are not thus decomposed ; like their hydroxides, they may be volatilised without decomposition. But all other carbonates are de- composed by exposure to a red heat. The process has already been described as a method of manufacturing quick- lime. Most carbonates, however, do not require the same high temperature ; a dull red heat suffices. And the oxides do not, as a rule, recombine with the carbon dioxide expelled, as does lime ; hence there is no danger of re-car- bonating the oxide.
Oxides produced by Heating Nitrates. — The nitrates, too, of nearly all the basic metals, yield the respec- tive oxides when they are heated to bright redness. The nitrates of the alkali metals in this instance, as in others, do not behave in this way. When heated they lose oxygen, but only at a very high temperature, forming the nitrites, a class of salts afterwards to be described. Thus, potassium nitrate undergoes the decomposition : 2KNO3 = 2KNO9 + O2. The product of heating the other nitrates, however, is the oxide, while a mixture of oxides of nitrogen is evolved. This may be supposed to take place in two stages : first, the nitrate may be imagined to decompose into the oxide and nitrogen pentoxide, thus : Zn(NO8)2 = ZnO + N2O5 ; the last compound is very easily decomposed by heat, and yields a lower oxide of nitrogen : 2N2O. = 4NO2 + O2 ; while if the temperature is over 600°, which is usually
SULPHIDES AND HYDROSULPHIDES 81
exceeded in decomposing the nitrates, the nitric peroxide is partly further decomposed into nitric oxide and oxygen : 2NO2 = 2NO + O2. The products, therefore, are NO2, NO, and O2.
A metal which forms two oxides, one containing more oxygen than the other, if the nitrate of the lower oxide is heated, yields the higher oxide. Cases of this are mercury, tin, and iron. Mercurous nitrate, carefully heated, gives, not mercurous oxide, Hg9O, but mercuric oxide, HgO : HgNO3 - HgO + NO2 ; similarly Sn(NO3)2 yields SnO2, and not SnO ; and Fe(NO3)2, Fe2O3, and not FeO.
Oxides produced by Heating Sulphates. — The sulphates require a higher temperature than the nitrates for their decomposition, consequently they are not generally used as a source of oxides. But the equivalents of mag- nesium, zinc, and some other metals have been determined by estimating the weight of oxide obtainable on heating a weighed amount of sulphate ; and ferrous sulphate has been distilled in fireclay retorts for many years past at Nord- hausen, in Saxony, for the purpose of making " Nordhausen sulphuric acid," H2S2OK, and red oxide of iron, Fe2O3, which, made in this way, has a fine colour, and is used as a paint. When a sulphate is heated, the gas which escapes is not entirely SO3, as might be imagined from the equation : MgSO4 = MgO + SO3 ; the high temperature decomposes most of the sulphur trioxide into the dioxide, SO0, and oxygen ; and the oxygen, in the case of ferrous sulphate, oxidises the FeO into Fe2O3.
Sulphides and Hydrosulphides. — The analogy between the elements oxygen and sulphur is well shown by comparing the sulphides of the elements of which the oxides have been described. Elements of the lithium group form both hydrosulphides and sulphides ; thus we know sodium hydro sulphide, NaSH, analogous to the hydroxide NaOH, and sulphide, similar in formula to the oxide Na9O, Na9S. Hydrogen sulphide is a weak acid ; hence, on passing hydrogen sulphide through
82 MODERN CHEMISTRY
a concentrated solution of sodium hydroxide at 95° until saturation is complete, white crystals of NaSH.2H2O
+ - deposit on evaporation. The equation is : NaOH.Aq
+ HSH.Aq = NaSH.Aq + H,O. On mixing the solu- tion with an equivalent quantity of sodium hydroxide and
+ evaporating, the sulphide is produced : Na SH . Aq +
+ - + - -
NaOH.Aq = Na2S.Aq + H2O. Here it must be sup- posed that the hydrogen of the hydrosulphide is present as an anion, and that it reacts with the hydroxyl of the caustic soda, forming water, while the sodium sulphide remains in solution in an ionised form, and can be recovered on evaporation in crystals with 9H2O. Similar compounds exist with potassium.
Calcium, strontium, and barium also form hydro- sulphides and sulphides, analogous in formula to the hydroxides and oxides. They are similarly prepared to the sodium compounds, but, as the metals are dyads, their formulas are M(SH)2 and MS; and there is an inter- mediate compound between the hydroxide and hydro- sulphide, having, in the case of calcium, the formula HSCaOH. They are also soluble in water. Magnesium, too, forms a hydrosulphide, probably Mg(SH)2; it is prepared by passing sulphuretted hydrogen into water in which magnesium oxide is suspended. It is unlike the hydrosulphides of the alkalies, for while they do not decompose with water, it, on the contrary, when its solution is heated to 80°, reacts with water, yielding hydroxide and sulphuretted hydrogen: Mg(SH)2.Aq + 2HOH = Mg(OH)2 -f 2H2S. The probable explanation of this change is that water is not wholly non-ionised, but that there are present some hydrogen ions ; these are not so inconsiderable in number, compared with those of the weak acid H2S ; on raising temperature, a certain amount of hydrogen sulphide is liberated, and, being volatile, it escapes,
THE SOLUBILITY-PRODUCT 83
and is no longer present to act on the magnesium hydroxide and reconvert it into sulphide.
Sulphides of boron, aluminium, chromium, and silicon are at once decomposed by water, and cannot, therefore, be produced in aqueous solution. They are white substances formed by heating the elements to a high temperature in a current of sulphur vapour.
The sulphides of copper, silver, gold, cadmium, mercury, indium, thallium, tin, lead, arsenic, anti- mony, and bismuth, and of the metals of the palladium and platinum groups, are all insoluble in water, or, to be more accurate, very sparingly soluble. They form no hydrosulphides. Hence they are precipitated from soluble salts of these metals by addition of sulphuretted hydrogen ; they form flocculent precipitates, usually characterised by striking colours, and are therefore generally used as a means of recognising the metal. CuS, Ag9S, Au0S3, HgS, T12S, T12S3, PbS, PtS2, and the other sulphides of the platinum group of metals are black ; CdS, SnS2, and As.,S3 are yellow ; ln9S3, SnS, and Bi9S3 are brown, and Sb9S3 is orange. These sulphides are not attacked by dilute acids. On the other hand, the sulphides of zinc, manganese, iron, cobalt, and nickel are not precipitated by hydrogen sulphide, but they are thrown down by a soluble sulphide or hydrosulphide, such as those of ammonium or sodium. They, too, form flocculent precipitates ; ZnS is white, MnS pink, and FeS, CoS, and NiS are black. The reason of the difference in the behaviour of the two classes of sulphides is an interesting one, and will be now explained.
Solubility** Product. — It has already been mentioned on p. 14 that the rate of chemical change depends on the amount of each of the reacting substances present in unit volume. This last is generally termed the " concentration " of these substances, for the more concentrated the solution the greater the mass present in unit volume. Now, if two
+ kinds of ions, such as Na and Cl, are present in solution,
84 MODERN CHEMISTRY
necessarily in equal numbers, the solution will also contain a certain number of molecules of non-ionised NaCl, formed by their union, and the relative number of ions and mole- cules will depend on the concentration ; the number of ions in proportion to the number of non-ionised molecules will be greater, the greater the dilution. For each dilutk (and for each temperature) a state of balance will result; the position of this equilibrium will depend on the relative rate at which ionisation and union of ions to form mole- cules go on ; if ionisation takes place twice as quickly as combinations of ions to form molecules, then two-thirds of the dissolved substance will exist as ions, the remain- ing third being non-ionised molecules. If the solution is made more concentrated by evaporation, the conditions are changed, and the rate of ionisation is reduced compared with the rate of union of ions with each other. Suppose that concentration be pushed so far that solid salt separates out ; the limit of concentration will be reached, since it is now impossible to alter the number of ions and of molecules in unit volume of the solution. The ratio will now remain constant, and if c and c be the concentrations of the ions (and they are, of course, equal), and if C be that of the non-ionised molecules, then c.c = LC9 k being a factor ex- pressing the relative proportions of the non-ionised mole- cules. If k is very small, then there are many molecules and few ions present ; if, on the contrary, k is large, the ions are numerous and the molecules few. The expression LC is termed the "solubility-product."
To take a specific case : — A solution of ammonia in water consists partly of the ions NH4 and OH, and partly of non-ionised molecules of NH4OH ; it is a weak base — that is, the number of non-ionised molecules is much greater than that of the ions. In a solution con- taining 1.7 grams of ammonia per litre (one- tenth normal solution), only 1.5 per cent, of the total number of mole- cules exist in the ionic state. Hence a solution of ammonia, unlike one of caustic soda or potash, gives no precipitate
INSOLUBILITY OF SULPHIDES 85
of hydroxide when added to a solution of salts of the relatively strong bases, such as calcium, strontium, or barium chlorides. With salts of the weaker base magnesia, however, ammonia produces a precipitate of magnesium hydroxide. It is possible still further to reduce the ion- isation of ammonia solution ; this can be done by the addition of an ammonium salt, such as the chloride, which, like most such salts, is highly ionised. The reason is, that
+
while (concentration of NH4) x (concentration of OH) = k x (concentration of NH4OH), if more ammonium ions be added, the number of hydroxl ions will diminish by union with NH4 ions, forming non-ionised ammonium hydroxide, because the increase of the number of ammonium ions will increase the value of the product on the left-hand side of the equation, and in order that it may balance that of the right, the relative number of molecules of NH4OH must be in- creased ; and we may see that if ammonium chloride is added to a solution of magnesium chloride, ammonia solution will no longer produce a precipitate of magnesium hydroxide ; the ammonia is too weak a base, that is, it contains too few hydroxyl ions, which are the reason of its basic nature.
Let us now return to the consideration of the insolubility of sulphides of the copper group in acids and the solubility of such sulphides as that of zinc. No substance, as has been before remarked, is wholly insoluble in water ; zinc sulphide, how- ever, belongs to the very sparingly soluble compounds.
+ +
Hence the product r(Zn) xr'(S) has a very small value, for it is equal to >£.£(ZnS), which must necessarily be very small, seeing that the compound is so sparingly soluble.
+ +
Now, the ions of H2S are H, H, and S ; but though the ionisation is very small, hydrogen sulphide being a very weak acid, they are yet sufficient to reach the value of the very small solubility-product ^.C'(ZnS). If, however,
the concentration of the S-ions is still further diminished
86 MODERN CHEMISTRY
+
by addition of some compound rich in H-ions, such as + - + +
HCl.Aq, then the product r(Zn) xc( S ) will be less than >£.£(ZnS), and there will be no precipitate ; or if hydrochloric acid be added to precipitated zinc sulphide, it will be dissolved. On the other hand, the addition of acetic acid, a weak acid, and poor in hydrogen ions, does not bring about solution of zinc sulphide ; indeed, the pre- cipitation of zinc from a solution of its acetate by hydrogen sulphide is almost complete.
The solubility-product of copper sulphide, and of the other sulphides which are not soluble in dilute acids, is still less ; hence hydrogen sulphide precipitates them from
acid solution, for the concentration of the S-ions of the hydrogen sulphide may be very much diminished without
+ +
the product r(Cu) x c ( S ) becoming less than ^.£(CuS), for CuS is still less soluble in water than ZnS.
Oxides and Hydroxides of Complex Groups. — The oxides and hydroxides of complex groups show analogy in their formulae, and often in their methods of preparation with the basic oxides and hydroxides. A few instances of these will now be given.
Ammonia (see p. 42) is very soluble in water ; at the ordinary temperature, no less than 800 volumes of the gas dissolve in one volume of water, forming a very pungently smelling solution named liquor ammonia. This solution con- sists for the most part of a mixture of liquid ammonia with water ; it probably also contains ammonium hydroxide, NH4OH, and, as already mentioned, less than 1.5 per cent.
+ of the ions NH4 and OH. It is, therefore, a weak base.
Hydrazine, N2H4, also forms a hydrate, N2H5OH, a fuming liquid with slight smell (and consequently in all probability fairly highly ionised) ; it boils at 119°, and is very corrosive, attacking wood, cork, and even glass. It has a strong reducing action, so that if added to a solution of
ALCOHOLS 87
cupric sulphate which contains cupric ions, Cu, it gives an immediate precipitate of cuprous oxide, Cu2O, nitrogen being evolved. Like ammonia, it precipitates such hydr- oxides as that of aluminium, iron, &c. Hydroxylamine, NH0OH, is a somewhat similar body, produced by passing nitric oxide, NO (see p. 97), through a mixture of granu- lated tin and hydrochloric acid, to which a little platinic chloride has been added ; the nascent hydrogen reduces the nitric oxide to hydroxylamine ; it unites with the hydrochloric acid, forming hydroxylamine hydrochloride, NH3OHC1. After removal of the tin by addition of sodium hydroxide and filtration, the solution is evaporated to dryness and mixed with alcohol ; hydroxylamine hydro- chloride dissolves, while sodium chloride remains. A solution of the base may be obtained by addition of silver hydroxide : NHgOHCl. Aq + AgOH. Aq = AgCl + NH3OH. Aq. If sodium methoxide (see p. 88) be added to a solution of the hydrochloride in methyl alcohol, the base is liberated, and can be separated from the alcohol by fractional distillation ; it is a volatile white solid. This compound is interesting, because the OH group is under no circumstances an ion ; its
solution in water must contain ions of NH3OH and OH, since it reacts like ammonium hydroxide.
Alcohols. — The hydroxides of the hydrocarbon radicles are, as mentioned on p. 67, termed alcohols.1 Of these there are very many, but a few only will be chosen to serve as examples : methyl alcohol, CHgOH, ethyl alcohol, CH3— CHQOH, as types of monohydric alcohols, which may be taken as the analogues of the
CH2OH hydroxides of the monad metals ; glycol, , a
CH2OH dihydric alcohol, may be likened to barium hydroxide,
1 A special class of such hydroxides derived from benzene, C6H6, are termed phenols. "Carbolic acid," C6H3OH, is the best known of these.
88 MODERN CHEMISTRY
CH2OH Ba(OH)9; and glycerine (glycerol), CHOH , is a
CH2OH
trihydric alcohol, as aluminium hydroxide is a trihydroxide. These substances differ from the hydroxides, however, by their being non-electrolytes, and therefore non-ionised. Or perhaps it is more correct to say that their conductivity is of the same order of magnitude, but less in value, than that of pure water. The corresponding halides, for ex- ample, CH3C1, C2H4C12, and C3H5C13, are also regarded as non-ionised ; they are practically insoluble in water. Nevertheless, methyl chloride has been transformed into methyl alcohol by heating with water to a high temperature in a sealed tube under pressure — CH3C1 + HOH = CH3OH + HC1 ; and the others, but preferably the bro- mides, may be similarly changed into hydroxides by heating with silver hydroxide, or with silver oxide and water : CH2Br CH2OH
CHBr + sAgOH.Aq = CHOH .Aq + 3 AgBr. Is it CH2Br CH2OH
possible that at a higher temperature the ionisation is suffi- cient (though it must be exceedingly small) to produce the interaction ?
The metals sodium and potassium dissolve in the alcohols, with evolution of hydrogen, forming compounds somewhat analogous to the hydroxides ; instead of hydrogen, however, they contain a hydrocarbon group : sodium methoxide, for example, has the formula Na(OCH3). Such sub- stances are white solids, like caustic soda.
Aldehydes. — The alcohols, if oxidised by boiling them with chromic acid, yield a class of bodies analogous to the oxides, termed aldehydes : CH3— CH2— OH + O = (CH3 - CH)"O + H2O. It will be noticed that ethane, CH3 — CH3, has lost two hydrogen atoms, and that the residue, CH3 — CH0, is now a dyad group, capable of com- bination with an atom of dyad oxygen. The aldehydes are volatile liquids, with strong odour, and those containing few
AMINES AND PHOSPHINES 89
atoms of carbon are miscible with water. They form easily decomposable compounds with water, which are di-hydr-
oxides ; e.g. ordinary aldehyde forms CH3
OH
they are called aldehydrols. When brought into contact with solutions from which hydrogen is being evolved, the aldehydrols lose oxygen, and are converted into alcohols :
/OH CH3CH/ + 2H = CH3-CH2OH + H2O.
The alcohols cannot be termed basic substances ; still, it is evident that they show analogy with the true bases in many respects.
Amines and Phosphines. — Derivatives of nitrogen, phosphorus, sulphur, and even of iodine and of oxygen, containing hydrocarbon groups, are however known, which are true bases, though weak ones. If ammonia in alcoholic solution be heated with excess of methyl iodide, tetra- methyl -ammonium iodide is formed : NH3 + 4CH3I = N(CH3)4I + sHI. This iodide, digested with water and silver hydroxide, exchanges iodine for hydroxyl, and after removal of the silver iodide by filtra- tion the solution may be evaporated to dryness. The residue is a white solid, of the formula N(CH^)4OH ; it is termed tetra -methyl- ammonium - hy dioxide ; in its reactions it shows great analogy with caustic potash, having a caustic taste, and producing precipitates with the usual salts of the metals. In solution it is more ionised than ammonium hydroxide, though less than that of potassium.
Phosphine, as remarked on p. 66, combines with hydrogen iodide, forming a salt, PH4I, phosphonium iodide, resembling ammonium chloride. But as it is decomposed by water into phosphine, PH3, and hydrogen iodide, an attempt to convert it into phosphonium hydroxide, PH4OH, cannot be made. Substituted phosphonium compounds,
90 MODERN CHEMISTRY
however, are known, in which a hydrocarbon radicle, such as methyl, replaces hydrogen. Sodium and phosphorus combine when heated together under an oil called xylene, forming PNa3 ; this body, treated with methyl iodide, yields trimethyl phosphine, P(CH3)., ; with more methyl iodide P(CH3)4I is formed; and its solution in water, which is not decomposed by the solvent, yields with silver hydroxide tetra-methyl-phosphonium hydroxide, P(CH3)4OH, a base resembling the corresponding ammonium compound.
These compounds exist owing to the double valency of nitrogen and of phosphorus, which can function either as triad or pentad. Double valency is to be noticed also with oxygen and with sulphur, although with the former tetrad combinations are far from stable, while with the latter both dyad and tetrad compounds can be formed.
Ethers. — Oxide of methyl and oxide of ethyl, which are usually named methyl and ethyl ethers, are formed by mixing solutions in alcohol of methyl or ethyl iodide with sodium methoxide or ethoxide : CH3I.Alc + NaOCHg. Ale = Nal + HgCOCHg. Ale. The ether has a low boiling- point, and can be separated by fractional distillation from the alcohol in which it is dissolved. Methyl ether is a gas; ethyl ether a volatile liquid, boiling at 37°. Such compounds can also be prepared more readily by distilling a mixture of the alcohol with sulphuric acid, which yields HCH^SO4, hydrogen methyl sulphate, with the alcohol : HCH3SO4 + CH3OH = HgCOCHg + H2SO4. Now, methyl ether and hydrochloric acid combine at a low
CH3\ /H temperature, yielding /O\ > but it is impossible to
CH/ NCI replace the chlorine by hydroxyl.
Similar sulphur compounds, however, are stable. Methyl sulphide, produced by the action of methyl iodide on potas-
CH3I 1C KI CH3,
sium sulphide, + J>S — + J>S, unites with
CHJ K/ KI CH/
ETHERS 91
^ methyl iodide, forming /S\ ? a compound con-
CH/ \I
taining tetrad sulphur ; with silver hydroxide it yields the corresponding tri-methyl-sulphonium hydroxide,
CH3 CH3
"
, a compound exhibiting basic properties. OH
From iodine, too, iodonium compounds have been prepared, in which the iodine functions as a triad ; and a hydroxide with basic properties is known.
CHAPTER V
Neutral Oxides — Peroxides— Action of Nitric Acid on Metals; on Oxidisable Substances — Complexity of Oxides — Spinels and Simi- lar Compounds.
THE properties of all chemical compounds show gradation ; and there is a slow transition from basic oxides and hydr- oxides, like those which we have been considering in the last chapter, to acid oxides and hydroxides. The transition takes place along two paths ; first, there are some oxides which are neither basic nor acid ; and second, a number of oxides exist which are either basic or acid, according to circumstances. We shall consider first the neutral oxides.
Peroxides. — In the potassium and calcium groups of elements, peroxides are known. When sodium is burned in air a light yellow powder is formed, sodium dioxide, of the formula Na2O2 ; potassium yields a tetroxide, K2O4. Both of these substances react with water, giving off oxygen ; but if they are very slowly added to the water, so that the temperature does not rise much, a solution is obtained. The corresponding barium compound is formed when barium monoxide is heated under pressure in air (see p. 13). On addition to water it forms a hydrate, probably Ba=O(OH)2.7H2O. On treatment with acid, hydrogen dioxide, H2O2, is formed ; and if sulphuric acid be added in theoretical amount to the barium dioxide, nearly insoluble barium sulphate is formed, along with a fairly pure solution of hydrogen dioxide : BaO=(OH)2.Aq
NEUTRAL OXIDES 93
+ H2SO4. Aq = BaSO4 + O=OH2. Aq. It can be purified, and indeed obtained anhydrous by distillation under very low pressure. It then forms a somewhat viscous, colourless liquid, with a sharp taste.
There is some doubt as to the constitution of hydrogen dioxide, and consequently of the dioxides from which it is derived. It is unlikely that barium ever acts as a tetrad, and much more probable that this character is to be attri- buted to oxygen ; hence the formula of its dioxide is more likely to be Ba=O=O, than O=Ba=-O ; and consequently hydrogen dioxide has more probably the formula O=OH2, than HO=OH. Indeed, hydrogen dioxide is possibly a weak acid, since the hydrated dioxides of calcium and barium are precipitated on addition of concentrated solu- tions of hydrogen dioxide to the hydroxides suspended in water. These substances have all bleaching power, for they readily part with their second atom of oxygen, and it is capable of oxidising coloured insoluble substances to colour- less soluble ones.
Neutral Oxides, Class I. — The next neutral oxides met with are carbon monoxide, CO, nitrous oxide, N2O, and nitric oxide, NO. These are all gases, but condense at low temperatures to colourless liquids, and at still lower, freeze to white solids.
Carbon monoxide is prepared by burning carbon in a supply of oxygen insufficient to convert it into the dioxide ; or by passing the dioxide over a layer of carbon, heated to redness. It appears that the monoxide is always the first product ; for if moisture be excluded during the combustion of carbon in oxygen, the amount of dioxide relatively to the monoxide is very small ; and it is known that if water-vapour be absent, carbon monoxide cannot be induced to explode with oxygen. If even the minutest amount of moisture be present, on passing a spark the union takes place with explosion. This phenomenon is not easily accounted for ; it is readily repre- sented by the equation 2CO +H2O + O2= 2CO2 + H2O. Can it be that at the very low pressure of the water-vapour
94 MODERN CHEMISTRY
+
a trace is ionised into H and OH, and that the OH furnishes the oxygen for the CO, the hydrogen recombining with oxygen to re-form the molecule of water ? For it has been found that no moisture is requisite to promote the union of oxygen and hydrogen if these gases be heated together. Phosphorus and sulphur, too, show reluctance in uniting with oxygen, in absence of moisture. In ordinary moist air, carbon monoxide burns with a blue flame. It is nearly insoluble in and has no action on water.
Other methods of preparing carbon monoxide are : by withdrawing the elements of water from formic acid by adding it drop by drop to warm concentrated sulphuric O
acid ; HC - OH + H2S04 = CO + H2SO4.H2O ; by heat- ing a mixture of oxalic acid with concentrated sulphuric
CO. OH acid ; | + H9S04 = CO + C02 + H9SO4.H2O ;
CO.OH
the carbon dioxide is separated from the monoxide by bubbling the mixture of gases through a solution of caustic potash, which absorbs the dioxide, allowing the monoxide to pass ; and lastly, by heating a mixture of potassium ferro- cyanide and fairly concentrated sulphuric acid ; K4Fe(CN)6 + 6H.2S04 + 6H20 = 2K?S04 + FeSO4 + 3 (NHJ,SO4 + 6CO. In the last reaction, it may be taken that hydro- cyanic acid, HCN, is first liberated, and that it reacts with water, forming ammonia and carbon monoxide : HCN + H2O = NH3 + CO ; the ammonia subsequently combines with the sulphuric acid.
If carbon monoxide is passed over metallic nickel or iron in a fine state of subdivision produced by reducing their oxides, volatile compounds are formed of the formulas Ni(CO)4, and Fe(CO)5; on exposing the latter to light gold-coloured crystals are formed, of the formula Fe2(CO)^. The nickel carbonyl boils at 43°, and the iron penta- carbonyl at 103°; di-ferro-hepta-carbonyl decomposes
NITROUS OXIDE 95
when even moderately heated. At 180° these compounds are decomposed into metal and carbon monoxide, the metal being deposited as a mirror on the hot surface.
Nitrous oxide, N2O, is most readily prepared by heating ammonium nitrate, NH4NO3 ; the equation is : NH4NO3-N2O + 2H2O. It is somewhat soluble in water, and is best colfected by downward displacement. The aqueous solution has a sweetish taste ; and the gas, if breathed, produces insensibility ; it is therefore frequently employed by dentists as an anaesthetic. If a mixture with air is respired, it produces with some persons a state of excitement, which has procured for it the name " laughing- gas." It is an endothermic compound, and if submitted to sudden shock it explodes with violence. It may be sup- posed that the fulminate used to explode it decomposes some molecules in the neighbourhood ; these, on decompos- ing, evolve heat, and decompose their neighbours, and the explosion rapidly travels throughout the gas ; the products are nitrogen and oxygen. A candle will burn in nitrous oxide, for the temperature of the flame is sufficiently high to decompose the gas, and the combustion proceeds as in dilute oxygen. Although nitrous oxide is not acted on by water or bases it has claims to be regarded as the anhydride of hyponitrous acid, from a solution of which it is liberated
N-OH N, by heat: || - li \O+H9O. As neither ammo-
N-OH N/
nium nitrate nor hyponitrous acid can be reproduced by bringing together nitrous oxide and water, its production by heating one of these compounds is termed an " irreversible reaction."
Action of Nitric Acid on Metals. — The product of the action of nitric acid on metals varies according to the metal acted on, the concentration of the acid, and the temperature. The acid in aqueous solution is more or less
-f ionised, the ions being H and NO3. If a metal of which
96 MODERN CHEMISTRY
the ions are highly electropositive is presented to these ions of nitric acid the hydrogen ions impart their charge to the non-ionised metal, which metal enters into solution as ions, while hydrogen is evolved. This is the case when nitric acid acts on magnesium, and theoretically also on aluminium, manganese, zinc, cadmium, iron, cobalt, and nickel, for all these metals in the ionic state have higher electro-affinity than hydrogen, and that in the order given. It may be termed the normal action of acids on metals, and repre-
+ +
sented thus : M + 2H = M + H2. But along with this action others take place in which the nitric ion is " re- duced " or deprived of oxygen. Some examples of this will now be given.
When silver is attacked by nitric acid, nitric peroxide, NO9, is produced, and partly evolved as gas. The react-
H NO3
ing substances are Ag, and + and ; one of the
H N03
NO3 groups loses oxygen, being converted into electrically
neutral NO2 and an ion of oxygen, O, which combines with the two hydrogen ions, forming water, non-ionised, H2O. But this leaves a negatively charged nitrate group without a corresponding positively charged partner ; more- over, the charge of the decomposed nitrate group is still available. An atom of silver, therefore, goes into solution as a positively charged ion, and restores electric equilibrium in the solution. With less concentrated acid the nitrate ion parts with two atoms of oxygen, requiring three negative electrons, in addition to the one originally attached to the
group NO3 ; to effect this three positive electrons must attach themselves to three atoms of silver, which then go
+ into solution as ions, hence the charge is : 3 Ag + 4-H +
4NO3 = NO + 2H2O + 3 Ag + 3NO3, the balance of elec-
OXIDES OF NITROGEN 97
trie charge not having been disturbed, although one nega- tive and one positive electron have disappeared. With metals yielding kations of higher potential, the reduction of the nitrate ion goes still farther ; nitrous oxide, N2O, nitrogen, and even ammonia may be produced, in relative amounts depending on the metal, on the concentration, and on the temperature. It may be taken that the lower the tem- perature, the less the concentration, and the higher the metal stands in the electro-negative series, the greater the reduc-
+ tion. The equations are : qM" + loH + ioNO3 = N2O
i2NO = N
2
6H,0 + 5 M+ i oN03 ; 4M" + i oH + i oNO3 = NH4 + "
. All these changes may proceed simultaneously ; but copper and moderately strong nitric acid yields fairly pure nitric oxide ; if more concentrated acid be employed, a mixture of varying proportions of nitric oxide and peroxide are evolved ; while by using zinc or magnesium and very dilute acid, nitrous oxide, nitro- gen, hydrogen, and ammonium nitrate are the main products.
Oxidation by means of Nitric Acid. — Action of the same nature occurs when an element capable of chang- ing its valency, i.e. the number of electrons associated with its ionised atom, is treated in the ionic condition with nitric
+ + acid. For example, the ferrous ion, Fe, on treatment
+ + + with nitric acid at 100° becomes ferric, Fe, while nitric
+ + +
oxide is evolved : 3 Fe + 6R + 4H + 4NO3 = NO + 2H2O
+ + +
+ 3Fe+ 3NO3 + 6R ; R being any monovalent anion. Such operations are usually spoken of as " oxidations in the wet way/'
Nitric oxide is a colourless gas, very sparingly soluble in water ; on bringing it into contact with oxygen, unless
VOL. II. G
98 MODERN CHEMISTRY
moisture is absolutely excluded, union takes place to form nitric peroxide, NO2, along with a trace of N2O3, nitrous anhydride. On sufficiently cooling nitric oxide it con- denses to a colourless liquid, and at a still lower temperature it forms a white solid.
Nitrous anhydride, strictly speaking, belongs to the class of acid-forming oxides ; its formula is N2O3, When nitric oxide and nitric peroxide are brought together, only a minute quantity of N2O,} is formed ; that is, because on converting it into the gaseous state it decomposes almost completely into these products. On cooling such a mixture, however, a blue liquid condenses, which has the formula N9Og. It will be afterwards alluded to.
Nitric peroxide, as usually seen mixed with air at ordinary temperatures, is an orange-coloured gas. When pure it condenses to an orange-red liquid, boiling at 22° ; it freezes at — 10° to a colourless solid. The liquid has a molecular weight corresponding to the formula N2O4, and the gas, at temperatures not much exceeding the boiling- point, consists mainly of the same substance. But as the temperature rises the colour grows darker, until, at 140°, it forms a blackish-red gas, consisting wholly of NO2. With progressive increase of temperature NO2 dissociates in its turn into NO and O2, and at 600° the change is complete. As temperature falls the action is reversed.
Neutral Oxides, Class II. — The next class of oxides comprises those which may be termed neutral, because they can act either as bases or as acids, according as they are treated with an acid or with a base. Their hydroxides may be comprised in the same class. A case of this kind has already been explained on p. 70 ; it is there shown that aluminium hydroxide, when treated with acids, yields salts of aluminium, while with bases aluminates are formed.
Complexity. — It appears probable that such oxides have molecular formulae more complex than those usually ascribed to them; for instance, aluminium oxide is certainly
NEUTRAL OXIDES 99
more complex than is implied by the usual formula A19O3 ; it may be A14O6 or A16O9, but there is no means at present of determining the degree of complexity of the molecule. The argument in favour of this view is the very high melting-points and boiling-points of such oxides. It is a well-known fact that as the molecular weight of compounds increases the boiling-point rises. Examples to illustrate this are best drawn from carbon compounds, where " poly- merism" is not infrequent ; that is, where compounds exist having the same percentage composition but molecular formulas, of which the higher ones are multiples of the lower one. We are acquainted with a series of com- pounds of carbon and hydrogen, of which the first member is ethylene, C2H4 ; bodies of the formulae C4Hg, C6H12, C8H16, C10H20, &c., are also known ; and the boiling-point increases with the molecular weight. Now, the chlorides of the elements are, as a rule, easily volatile, and have low melting-points ; and where it happens that both chloride and oxide have a simple molecular formula, as, for example, carbon tetrachloride, CC14, and carbon dioxide, CO9, the chloride has always a higher boiling-point than the oxide. It would appear to follow, therefore, that if the oxides of the metals had as simple molecular formulae as the chlorides they would show more volatility than the latter. As this is not the case, the presumption is that the oxides possess more complex formulae than we are in the habit of ascribing to them. This probability will be dealt with as occasion arises.
Among the oxides and hydroxides which exhibit the power of acting both as acid and basic compounds are cupric hydroxide, Cu(OH)2, which dissolves in a con- centrated solution of potassium hydroxide with a dark blue colour ; zinc and cadmium hydroxides, which dissolve in excess of alkali ; sodium zincate has been separated by addition of alcohol, and is precipitated in white needles of the formula Na2ZnO2.8H2O; and aluminium hydroxide, which dissolves in alkali, forming an aluminate, MA1O2 ;
ioo MODERN CHEMISTRY
stannous and plumbous hydroxides, Sn(OH)2 and Pb(OH)2, dissolve in alkalies, forming Compounds no doubt analogous to zincates. Chromous, ferrous, manganous, cobaltous, and nickelous hydroxides are not thus soluble. Chromic hydroxide, however, is soluble in soda, probably forming a compound like sodium aluminate ; unlike the latter, chromium hydroxide is thrown down on boiling the solution.
But such compounds, when they do not contain sodium or potassium, are often insoluble in water, and then they cannot be prepared by the action of the one hydroxide on the other. The oxides combine when heated together in the dry condition, and sometimes when the compound formed is decomposed by water (hydrolysed) it is con- venient to prepare it either from the oxide or from the carbonate.
Spinels. — A considerable number of compounds, analogous to the aluminates, is produced in this way, and many of them are found in nature as minerals. To this class belong the " spinels," so called because one of their number, the native aluminate of magnesium, had received this name. Viewed as a combination of oxides, such com- pounds possess the general formula M2OrMO, and they can be prepared by heating the sesquioxide (a name given to oxides when the proportion between the metal and the oxygen is as one to one and a half, or, more correctly, as two to three) with the monoxide. The spinels all crystal- lise in regular octahedra ; they are therefore said to be isomorphous with each other. Viewed as aluminates, they may be written M"(MO9)9 ; compare NaAlO.,. Among them are true spinel, Mg(AlO2)2; franklinite, Zn(FeO2)2; chrysoberyl, Be(AlO2)2; and chromite, or "chrome-iron ore/' Fe(CrO2)2. But it is not neces- sary that the metals of a spinel should be different ones ; if a metal is capable of existing in two forms, e.g. as dyad and triad, it may form a similar compound. Such are magnetite^ or "magnetic iron ore/7 Fe"(Fe'"O2)2, and
SPINELS • • 101
hausmanite, Mn"(Mn'"C,)<,, th^ lirst atom- of .Iron or manganese being dyad, like magnesium, and the second triad, like aluminium.
Reasoning by analogy, it would appear not unlikely that native oxides, such as alumina (corundum, ruby, sapphire), or iron sesquioxide (haematite), may be in reality an aluminate of alumina, A1(A1O0)3, or ferric ferrite, Fe(FeO,)8.
A common test for zinc and aluminium is to heat together before the blow-pipe the salt suspected to contain the metal with cobalt nitrate ; it is probable that the green colour produced by zinc is due to the formation of a cobalt zincate, Co(ZnO0), and the blue colour shown by alumina to a similar body, Co(AlO2)2.
When lead is heated to redness in air the first product of its oxidation is litharge, PbO ; on continuing the ap- plication of heat, at a carefully regulated temperature, the yellow litharge becomes red, and the product of the action is minium or " red lead," PbgO4. Now, on treating red lead with dilute nitric acid, lead nitrate dissolves, while lead dioxide, hydrated, remains as an insoluble residue. Red lead, therefore, may be regarded as a compound between two molecules of monoxide and one of dioxide, 2 PbO + PbO2 ; the former reacts with nitric acid forming the nitrate, while the latter remains. Now, if the dioxide be heated with caustic potash it dissolves, forming potassium plumbate, K.JPbO3 ; and red lead may be regarded as a
/Pb"\ iv basic plumbous plumbate, O<^ VPbOg) ; "basic,"
\pb"/
because the first written atoms of lead are partly oxide, partly salt ; they are dyad, while the second atom of lead is tetrad.
It is possible to regard nitric peroxide in this light as a nitrate of nitrosyl, O=N— NO3 ; but its easy decom- position into NO2 when heated militates against the view. Compounds of antimony and bismuth, having the
102 MODERN CHEMISTRY
formula* Sfe-2O4 '.arid Ei,jO44 inay be similarly regarded as O=Sb(SbO3) and G>=Bi(BiO3) ; of this, however, there is no proof.
Manganese and chromium also form " dioxides," to which the simple formulas MnO2 and CrO9 are usually
o ' ,o
attributed ; they, too, may be written x/Ci\ /Cr
O^ XX
o, TI xx
and VMn/ />Mn". They would then be termed
O^ NX
chromous chromate and manganous manganate. Such ideas must be regarded as speculative, but there can be little doubt that the formulae are more complex than they are usually written. The former is a snuff-coloured powder, produced by the action of nitric oxide on a chromate ; the latter, formed by oxidising and precipitating a manganous salt simultaneously, is best prepared in a hydrated state by the action of a hypobromite on a manganous salt: MnCl9.Aq 4-NaOBr.Aq + 2NaOH.Aq - O=Mn(OH)2 + NaBr.Aq + 2NaCl.Aq. It is a common black mineral in the anhy- drous state, and is known as pyrolusite. It will be re- membered that the ordinary method of preparing chlorine is to heat this mineral with dilute hydrochloric acid, and also that on heating alone it furnishes oxygen, being itself con- verted into Mn3O4, a brown powder, which may be formu- lated as a spinel, viz. (O=Mn— O)2=Mn.
In concluding this chapter on neutral oxides, it may be mentioned that there are a few which, acting generally as feeble bases, yet display feebly acid properties if in the presence of a strong base like soda or potash. Such are the oxides of gold, the metals of the platinum group, and of titanium, zirconium, and thorium. The chlorides of these elements are soluble in water, as also the sulphates and nitrates of the last three. Sulphates of gold and plati- num, however, are hydrolysed by water, giving oxides and sulphuric acid, thus : Pt(SO4)2 + zHOH = PtO2 + 2H2SO4.
NEUTRAL OXIDES 103
Salts of these elements, on treatment with soda, yield no pre- cipitate, for they are dissolved by the alkali ; the compounds formed are indefinite, but it may be supposed that they contain aurate, MAuO2.Aq, or platinate, titanate, zirconate, or thorate, MPtO3.Aq, &c. Iron and calcium titanates occur native ; FeTiO3 is termed ilmenite, and CaTiO3 perowskite. The first is isomorphous with and crystal- lises along with native ferric oxide ; the ore is known as " titanic iron ore." It is the commonest compound of titanium.
CHAPTER VI
Anhydrides — Acids and Salts — Basic and Acid Chlorides— The Borates—The Carbonates and Thiocarbonates — Other Acids containing Car- bon; their Salts with Alcohol Radicals— The Silicic Acids and the Silicates.
Basic Salts. — Many compounds are known which are at the same time chloride and oxide, or chloride and hydr- oxide of elements. Where the element with which the oxygen and chlorine is combined is one which forms a basic oxide, the compounds in question are termed basic chlorides. Similarly, there are basic bromides and iodides. For ex- ample, zinc oxide heated with zinc chloride forms oxychlo- rides, of which the simplest example is Cl~~Zn~~O~Zir~Cl; aluminium chloride, evaporated with water, has its chlo- rine gradually replaced by hydroxyl, forming successively C12=A1(OH), C1-A1=(OH)2, and finally, A1(OH)8, though at a temperature sufficient to complete the reaction, the aluminium would probably form the condensed hydr- oxide O=A1OH instead of the trihydroxide. We shall see later that other groups, playing a part analogous to that of the chlorine in a basic salt, may also exist in basic salts.
Acid Chlorides. — Another class of double oxides and chlorides exists, most of which are easily volatile, and which therefore are of known molecular weight. These are the so-called "acid chlorides" — oxychlorides of those elements which form acids. These are related to acids, in as much as by replacement of their chlorine by hydroxyl,
BORATES 105
acids are formed. It will therefore be convenient to con- sider them along with the acids to which they are related.
A general idea has already been given of the nature of acids in describing the hydroxides of zinc and of aluminium. As a rule, acids are condensed hydroxides ; that is, hydroxides which, having lost the elements of water, are partly oxides, partly hydroxides. They also possess the property of ionising into one or more hydrogen ions and an electro-negatively charged radical. In following the order of the periodic table, after such feebly acidic hydroxides as those of zinc and aluminium, hydroxide of boron claims attention.
Borates. — In certain lakes in California the water, when evaporated, deposits crystals of the formula Na2B4OrioH2O ; this substance is named borax. It is a white, crystalline salt, easily soluble in hot water, but sparingly soluble in cold. When mixed with sulphuric acid nacreous scales separate of the formula B(OH)3, or, as it is usual in writing the formulas of acids to place the hydrogen atoms first, H3BO3. Boracic acid hardly deserves the name of acid ; in aqueous solution it exists almost entirely in the non-ionised state. No ions are volatile ; but this compound issues in Tuscany and in the Lipari Islands along with steam from cavities in the ground, termed suffioni ; it is easily recognised, for it imparts a green colour to a flame held in the steam. When heated to 100° boracic acid loses water and is changed into metaboracic acid, O=B-OH, a vitreous substance ; and at a red-heat boron oxide, Bi;O3, is left as a transparent, colourless glass. Its con- stitution is O=B— O— B=O.
The borates of the alkalies are prepared by mixing boracic acid with the hydroxide of the alkali metal ; although there are very few hydrogen ions in an aqueous solution of boracic acid, however dilute, yet some of those present combine with the hydroxyl ions of the alkali, forming
water, thus: H3BO3.Aq + 3NaOH.Aq = Na3BO3.Aq +
io6 MODERN CHEMISTRY
3H2O. But there are so few ions present, that those of the water, which, it will be remembered, is ionised, although to an extremely minute extent, are yet sufficiently numerous to bear some proportion to those of the boracic acid ; hence the reaction given above is perceptibly reversed, and on dissolving borax in water it is " hydrolysed," that is, split "by the hydrogen and hydroxyl ions of the water into non- ionised boracic acid and caustic soda, the latter, of course, largely ionised as usual. It is therefore possible to estimate the sodium of borax by addition of a solution of a strong acid, such as hydrochloric or sulphuric acid of known con- centration, just as if no boracic acid were present, provided methyl-orange be used as an indicator. (See p. 75.) Thus the addition of 36.5 grams (H=i; 01 = 35.5) of hydrochloric acid, dissolved in a litre of water (such a solution is termed a "normal solution"), to 191 grams of a solution of crystallised borax in a litre of water (1/2 [Na2 = 46 + B4 = 44 + O7 = 112 + ioH2O = 180]) ; (in all 1/2 of 382) gives a solution which is neutral to methyl- orange.
Fused borax has the property of dissolving oxides of the metals, forming complex borates ; certain of these are coloured, and their formation is often made use of for detecting the presence of such metals as copper (blue), silver (yellow), chromium (green), ferric iron (yellow), ferrous iron (bottle-green), manganese (amethyst, when heated in a flame containing excess of oxygen), cobalt (blue), and nickel (reddish). Borax is also used for soldering easily oxidisable metals, such as iron, copper, or brass ; the film of oxide which prevents the metal touching and alloying with the solder is thus removed. Both borax and boracic acid have considerable antiseptic properties, and are used for preserving eggs, milk, and other animal and vegetable substances.
Carbonates and Thiocarbonates. — The car- bonates and the thio carbonates are derivatives of carbon dioxide (or ' rather of carbon oxy-hydroxide, commonly
CARBONATES 107
called carbonic acid), and of carbon disulphide. Carbon is a tetrad, and. the analogue of carbon tetrachloride would be the tetrahydroxide, C(OH)4; but this body is unstable, and its first anhydride, O=C(OH)2, is known only in aqueous solution. However, carbonyl chloride, O=CC12, exists ; it is produced by the direct union of carbonic oxide with chlorine, when a mixture of both gases is exposed to sunlight ; it was formerly known as " phosgene gas," meaning "made by light"; but it is more conveniently prepared by passing a mixture of the two gases over animal charcoal heated to redness. It condenses to a liquid, boiling at 8.4°. It is immediately decomposed by water, thus: O=CC12 + 2HOH = O=C(OH)2+ 2HC1; if sufficient water is present, the carbonic acid can remain in solution. The existence of the oxychloride establishes the formula of carbonic acid.
Carbonic acid is a very easily decomposable substance ; if liberated, unless a great deal of water be present, it splits into its anhydride, CO2, and water: O=C(OH)2 = CO0 + H2O. The anhydride is a colourless gas, which con- denses to a solid at about —80° ; it can be liquefied only under pressure. Carbon dioxide, or carbonic anhydride, is produced by heating a carbonate ; as already remarked, all carbonates, except those of the alkaline metals, are decomposed by heat, forming oxides, and evolving carbon dioxide. It is also produced when carbon or carbon monoxide is burned with excess of oxygen. Lastly, it is produced in large quantities during the process of fermenta- tion. Glucose, or grape sugar, either produced by the hydrolysis of starch or extracted from fruits like grapes, when mixed in dilute aqueous solution with yeast, a vegetable organism, decomposes into ethyl alcohol and carbon dioxide, thus: C6H12O6 = 2C2H5OH + 2CO2. The carbon di- oxide being heavier than air, collects in the fermenting tuns ; it is now often collected and compressed until it liquefies ; and the liquid on expansion solidifies to a snow- like solid, used for producing low temperatures.
io8 MODERN CHEMISTRY
A solution of carbonic anhydride in water contains carbonic acid, O=C(OH)9, which is a very weak acid owing to the small extent of its ionisation. It is probable, too, that liquid carbon dioxide exists in the solution, mixed, but not combined with the water. Carbonic acid reacts with sodium, potassium, calcium, or barium hydroxide,
f - - + •
forming carbonate of the metal: H2CO3.Aq + zNaOH.Aq
= Nt2C03.Aq + 2H20 ; H2C03.Aq + Ca(OH)2.Aq - CaCO3 + 2H9O. In such actions it is only the ionised portion of the acid which reacts, and the hydrogen ions form water ; when these are removed another portion becomes ionised in order to restore equilibrium ; it reacts in its turn until all has become transformed. On evaporation of the solution the alkaline carbonate is left as a white crystal- line salt; hydrated sodium carbonate, Na9CO3. ioH2O, is ordinary washing-soda. All other carbonates are insol- uble in water, and are consequently thrown down as precipitates on adding a solution of sodium carbonate to any ionised solution of other metals. They form flocculent precipitates, generally possessing the colour of the ion of the metal ; thus copper carbonate is blue, ferrous green, cobalt pink, and so on. But with the exception of the carbonates of the metals of the sodium and calcium groups all other precipitated carbonates are " basic," that is, they are partly hydroxides, partly carbonates. Copper carbonate, for example, may be assigned the formula
/O-Cu-OH O=C<f ; it will be noticed that each atom
X0-Cu-OH
of copper is combined with the oxygen of the carbonic residue on the one hand, and with hydroxyl on the other. The paint known as " white lead " consists of a basic carbonate of lead, more complex than the example given above, of the formula
HO-Pb-0-(CO)-0-Pb-0-(CO)-0-Pb-OH.
Native Carbonates. — Many carbonates exist in the
ACID CARBONATES 109
native state ; some are widely distributed minerals. Among these are Iceland or calc-spar, arragonite, limestone, chalk, and marble, all of them calcium carbonate; strontianite, SrCO3 ; witherite, BaCO3 ; spathic iron ore, FeCOs, also named clay-band when contaminated with clay, and black-band when mixed with shale. Magnesite is MgCO3 ; dolomite, a mixture of magne- sium and calcium carbonates ; calamine, ZnCO3 ; and cerussite, PbCO3. Malachite and azurite are basic car-
/O-Cu-OH bonates of copper, O=C\ , and
X0-Cu-OH O O
HO-Cu-0-C-O-Cu-O-C-O-Cu-OH.
We see here again that with weak bases, such as the hydroxides of most metals, the carbonates tend to become basic, that is, to be hydrolysed. This is why the preci- pitates obtained on adding a soluble carbonate to a salt of such metals are basic, and not normal carbonates.
"Acid" Carbonates. — The name "acid carbon- ate " is given to a double carbonate of hydrogen and a metal. Such bodies are prepared by the method which always is adopted for the preparation of double salts — by
/ONa mixture. Hydrogen sodium carbonate, O=C\ ;
XOH
the corresponding potassium compound ; hydrogen cal- O O
I! II
cium carbonate HO-C- O-Ca-O-C-OH, a ferrous carbonate of similar formula, and many others are all formed when carbonic acid and the respective normal car- bonate are mixed, the mixture being kept cold. On raising the temperature of all of these, carbon dioxide escapes, and the neutral carbonate is again formed. " Acid " carbonate of sodium is the common " baking-soda ; " hydrogen calcium
no MODERN CHEMISTRY
carbonate is a constituent of many natural waters, and is the cause of what is termed " temporary hardness " ; for on boiling the water the neutral carbonate is precipitated, and the water ceases to be "hard." The same result may be effected, paradoxical as it may appear, by the addition of lime-water ; for then sufficient calcium hydroxide is present to form normal calcium carbonate with the hydrogen carbon- ate, thus : Ca(HCO3)2.Aq + Ca(OH)2.Aq = 2CaCO3 + 2H2O.Aq. Hydrogen ferrous carbonate is a constituent of chalybeate springs ; on exposure to the atmosphere the iron is oxidised to ferric hydroxide, and the carbonic acid, being too weak an acid to form a carbonate with such a weak base as that, escapes: 2Fe(HCO3)2.Aq+ 5H2O + O = 2Fe(OH)3 + 4H2CO3.Aq. The ferric hydroxide is deposited as a brown scum on the banks of the streams flowing from such wells.
Carbonates of Radicals. — Although normal hydr- oxide of carbon is unknown, yet if the hydrogen be replaced by ethyl, — C2H5, the compound is stable. The compound, which is produced by the action of carbon tetrachloride on sodium ethoxide, CC14 + 4Na-O-Ck>H5 = 4NaCl + C(O-C2H5)4, is the analogue of C(OH)4. It is a volatile liquid, and is named ethyl orthocarbonate. And a corresponding carbonate of ethyl, O=C(OC2H5)2, the analogue of carbonic acid, O=C(OH)2, is formed by treating carbonyl chloride 'with alcohol: O=CCU + 2HO-C2H5= O=C(OC2H5)2 + 2HC1. These compounds are volatile, and can be weighed in the state of vapour, hence their molecular weights are known, and this is an additional proof of the correctness of the formulae ascribed to carbonic acid and the carbonates.
Thiocarbonates. — The sulphocarbonates, or thiocar- bonates (from the Greek theion, sulphur) form a class of salts analogous to the carbonates, both in their formulas and in the method of their preparation. Carbon disulphide, a volatile liquid, boiling at 46°, possessing a disagreeable smell, is produced when sulphur vapour is led over charcoal
THIOCARBONATES in
heated to redness in a fireclay tube ; in fact, the carbon is burned in sulphur gas. When shaken with a concentrated aqueous solution of the sulphide of sodium or potassium, it dissolves, forming the compound Na2CS3, or K9CS3. These thiocarbonates, like the carbonates, are white, crystal- line salts ; on adding acid, thiocarbonic acid separates as an oil ; it slowly decomposes, especially if warmed, into hydrogen sulphide and carbon disulphide. Many of its salts are insoluble, and may be prepared by precipitation.
The formula of carbon dioxide is CO2, that of carbon disulphide CS2 ; and it is evident that an intermediate substance should exist of the formula COS. This sub- stance is carbon oxysulphide. It is a gas, prepared by heating thiocyanic acid, HSCN, the ammonium salt of which is produced when ammonia is passed through a mixture of carbon disulphide and alcohol : CS2 + 2NH3.Alc = H2S + (NHJSCN.Alc. On evaporation of the alcohol the ammonium thiocyanate crystallises out. This salt, dis- tilled with sulphuric acid, yields in passing the acid HSCN, which, on account of the high temperature, reacts with water, forming ammonia (which yields ammonium sulphate with the sulphuric acid) and carbon oxysulphide, COS : HSCN + H2O - NH3 + COS.
Like nitrous oxide, carbon disulphide is an endothermic compound, and can consequently be decomposed by shock ; when a fulminate is exploded in it, it is resolved into carbon and sulphur. On the other hand, carbon dioxide and oxy- sulphide are exothermic compounds, heat being evolved during their formation.
Acids containing Carbon. — An enormous number of acids containing carbon is known, in which the acidic carbon atom is combined with oxygen and hydroxyl, and also with hydrocarbon residues, such as methyl or ethyl, or with some more complex group of carbon atoms. The simplest
O
of these is formic acid, H— C— OH. Acetic acid is
ii2 MODERN CHEMISTRY
O
II
methyl-formic acid (CH3)— C— OH ; ethyl-formic acid is
O ll named propionic acid ; its formula is CH3— CH2— C— OH.
0=>C-OH Oxalic acid is to be regarded as di-carboxyl, I ,
O=C-OH
the name carboxyl being a contracted form of" carb(onyl hydr)oxyl" ; it is commonly written —CO— OH.
Formic acid (from formica, an ant) is contained in ants and stinging nettles. Sodium formate is produced when carbon monoxide is left in contact with sodium hydroxide ; the reaction takes a considerable time: CO+NaOH = H— CO— ONa. It is also formed by heating oxalic acid, better in presence of glycerine: (CO— OH)2 = CO0 + H— CO— OH. It is a colourless, pungently smelling liquid, boiling at 99°, and a fairly strong acid in aqueous solution ; it is poisonous. Its salts are crystalline, and possess the colours of the metallic ions which they contain. When warmed with concentrated sulphuric acid, or with other substances capable of withdrawing water, it yields carbon monoxide. Yet CO is not the true anhydride of formic acid, seeing that an anhydride can be obtained only from loss of the elements of water from hydroxyl groups, for formic acid contains the group H— C= ; the real anhydride
o o
would be H— C— O— C— H ; it is unknown.
Acetic acid is the acid constituent of vinegar, and is a solid, melting at 17° to a liquid, boiling at 118°. It can be formed synthetically by bringing into contact carbon dioxide and sodium methide, a compound of the formula Na-CH3 ; the equation is : Na-CH3 + CO2 = H3C— CO— ONa ; the sodium salt, distilled with sulphuric acid, yields acetic acid. It is produced on a large scale by the distillation of wood ; the distillate consists mainly of
"ORGANIC ACIDS" 113
acetic acid and methyl alcohol ; it is neutralised with lime, and distilled, when the alcohol passes over, leaving behind the calcium acetate ; this is evaporated to dryness, and heated, so as to char tarry matters, also produced when wood is distilled ; the calcium salt is finally distilled with sulphuric acid. Acetic acid is also formed by the oxidation of aldehyde (p. 88), which is itself an oxi- dation-product of alcohol. The connection between these bodies is : CH3-CH2-OH, CH3-CH=O, and O ii
CH3— C— OH. Aldehyde may be regarded as the anhydride of CHg— CH=(OH)2, and acetic acid of CH3 = C(OH)3. The usual oxidising agent is chromic acid ; if the product of oxidation is conveyed away as it is formed by sloping the condenser downwards, aldehyde is obtained ; if the aldehyde is returned to the oxidising mixture by sloping the condenser upwards, and cooling with ice and water, the product is acetic acid. The oxidation is also effected by an organism called " mother of vinegar " ; sour wine or beer is allowed to trickle down a cask filled with shavings of beech-wood, on which the slimy masses of the organism are growing ; oxygen enters, and the vinegar flows out at the bottom of the cask.
On distilling acetic acid with phosphorus pentachloride, hydroxyl is exchanged for chlorine : 4CH3— CO— OH + PC15 = 4CH3-CO-C1 + H8PO4 + HC1. the compound obtained is named acetyl chloride ; acetic acid may be regarded as hydroxide of the group (CHg— C=O)— , and on treating acetyl chloride with water it is at once formed : CH3-COC1 + H-OH = CH3-CO-OH + HC1. And aldehyde may be regarded as a hydride of acetyl, (CH3— CO)— H. A similar body cannot be made from formic acid, for it decomposes into carbon monoxide and hydrogen chloride : H-CO-C1 = CO + HC1.
Oxalic acid is contained as hydrogen-potassium salt in the plants sorrel and rumex. It can be prepared by the
ii4 MODERN CHEMISTRY
oxidation of sugar with concentrated nitric acid, or by heating sawdust with a mixture of caustic soda and potash in shallow trays ; on treating the charred residue with water, sodium oxalate, a comparatively insoluble salt, remains, while the excess of alkali dissolves ; the sodium oxalate is extracted with boiling water, and calcium chloride is added ; this precipitates the almost insoluble calcium oxalate ; and on digesting it with the equivalent amount of sulphuric acid, sparingly soluble calcium sulphate remains, while oxalic acid dissolves. The filtered solution, when
C(OH)3 evaporated, deposits crystals of ortho-oxalic acid, |
C(OH)3 CO-OH which, at 100°, dehydrate to I . Oxalic acid is
CO-OH
a di-basic acid, and its salts, like those