Chapter I., Note 65. 100 parts of water dissolve

0° 20° 40° 60° 80° BaO 1·5 3·5 7·4 18·8 90·8 SrO 0·3 0·7 1·4 3 9

Supersaturated solutions are easily formed.

The anhydrous oxide BaO fuses in the oxyhydrogen flame. When ignited in the vapour of potassium, the latter takes up the oxygen; whilst in chlorine, oxygen is separated and barium chloride formed.

[55 bis] Brugellmann, by heating BaH_{2}O_{2} in a graphite or clay crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in a platinum crucible--in crystals belonging to the cubical system, sp. gr. 5·74. SrO is obtained in the latter form from the nitrate. The following are the specific gravities of the oxides from different sources:--

MgO CaO SrO from RN_{2}O_{6} 3·38 3·25 4·75 " RCO_{3} 3·48 3·26 4·45 " RH_{2}O_{2} 3·41 3·25 4·57

[56] The property of barium oxide of absorbing oxygen when heated, and giving the peroxide, BaO_{2}, is very characteristic for this oxide (_see_ Chapter III., Note 7). It only belongs to the anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides of calcium and strontium may be obtained by means of hydrogen peroxide. Barium peroxide is insoluble in water, but is able to form a hydrate with it, and also to combine with hydrogen peroxide, forming a very unstable compound having the composition BaH_{2}O_{4} (obtained by Professor Schöne), which in course of time evolves oxygen (Chapter IV., Note 21).

Barium oxide is decomposed when heated with potassium; fused barium chloride is decomposed, as Davy showed, by the action of a galvanic current, forming metallic _barium_; and Crookes (1862) obtained an amalgam of barium from which the mercury could easily be driven off, by heating sodium amalgam in a saturated solution of barium chloride. Strontium is obtained by the same processes. Both metals are soluble in mercury, and seem to be non-volatile or only very slightly volatile. They are both heavier than water; the specific gravity of barium is 3·6, and of strontium 2·5. They both decompose water at the ordinary temperature, like the metals of the alkalis.

Barium and strontium as saline elements are characterised by their powerful basic properties, so that they form acid salts with difficulty, and scarcely form basic salts. On comparing them together and with calcium, it is evident that the alkaline properties in this group (as in the group potassium, rubidium, cæsium) increase with the atomic weight, and this succession clearly shows itself in many of their corresponding compounds. Thus, for instance, the solubility of the hydroxides RH_{2}O_{2} and the specific gravity[57] rise in passing from calcium to strontium and barium, while the solubility of the sulphates decreases,[58] and therefore in the case of magnesium and beryllium, as metals whose atomic weights are still less, we should expect the solubility of the sulphates to be greater, and this is in reality the case.

[57] Even in solutions a gradual progression in the increase of the specific gravity shows itself, not only for equivalent solutions (for instance, RCl_{2} + 200H_{2}O), but even with an equal percentage composition, as is seen from the curves giving the specific gravity (water 4° = 10,000) at 15° (for barium chloride, according to Bourdiakoff's determinations):

BeCl_{2} : S = 9,992 + 67·21_p_ + 0·111_p_^2 CaCl_{2} : S = 9,992 + 80·24_p_ + 0·476_p_^2 SrCl_{2} : S = 9,992 + 85·57_p_ + 0·733_p_^2 BaCl_{2} : S = 9,992 + 86·56_p_ + 0·813_p_^2

[58] One part of calcium sulphate at the ordinary temperature requires about 500 parts of water for solution, strontium sulphate about 7,000 parts, barium sulphate about 400,000 parts, whilst beryllium sulphate is easily soluble in water.

Just as in the series of the alkali metals we saw the metals potassium, rubidium, and cæsium approaching near to each other in their properties, and allied to them two metals having smaller combining weights--namely, sodium, and the lightest of all, lithium, which all exhibited certain peculiar characteristic properties--so also in the case of the metals of the alkaline earths we find, besides calcium, barium, and strontium, the metal magnesium and also _beryllium_ or _glucinum_. In respect to the magnitude of its atomic weight, this last occupies the same position in the series of the metals of the alkaline earths as lithium does in the series of the alkali metals, for the combining weight of beryllium, Be or Gl = 9. This combining weight is greater than that of lithium (7), as the combining weight of magnesium (24) is greater than that of sodium (23), and as that of calcium (40) is greater than that of potassium (39), &c.[59] Beryllium was so named because it occurs in the mineral _beryl_. The metal is also called glucinum (from the Greek word [Greek: glykys], 'sweet'), because its salts have a sweet taste. It occurs in beryl, aquamarine, the emerald, and other minerals, which are generally of a green colour; they are sometimes found in considerable masses, but as a rule are comparatively rare and, as transparent crystals, form precious stones. The composition of beryl and of the emerald is as follows: Al_{2}O_{3},3BeO,6SiO_{2}. The Siberian and Brazilian beryls are the best known. The specific gravity of beryl is about 2·7. Beryllium oxide, from the feebleness of its basic properties, presents an analogy to aluminium oxide in the same way that lithium oxide is analogous to magnesium oxide.[60] Owing to its rare occurrence in nature, to the absence of any especially distinct individual properties, and to the possibility of foretelling them to a certain extent on the basis of the periodic system of the elements given in the following chapter, and owing to the brevity of this treatise, we will not discuss at any length the compounds of beryllium, and will only observe that their individuality was pointed out in 1798 by Vauquelin, and that metallic beryllium was obtained by Wöhler and Bussy. Wöhler obtained _metallic beryllium_ (like magnesium) by acting on beryllium chloride, BeCl_{2}, with potassium (it is best prepared by fusing K_{2}BeF_{4} with Na). Metallic beryllium has a specific gravity 1·64 (Nilson and Pettersson). It is very infusible, melting at nearly the same temperature as silver, which it resembles in its white colour and lustre. It is characterised by the fact that it is very difficultly oxidised, and even in the oxidising flame of the blowpipe is only superficially covered by a coating of oxide; it does not burn in pure oxygen, and does not decompose water at the ordinary temperature or at a red heat, but gaseous hydrochloric acid is decomposed by it when slightly heated, with evolution of hydrogen and development of a considerable amount of heat. Even dilute hydrochloric acid acts in the same manner at the ordinary temperature. Beryllium also acts easily on sulphuric acid, but it is remarkable that neither dilute nor strong nitric acid acts on beryllium, which seems especially able to resist oxidising agents. Potassium hydroxide acts on beryllium as on aluminium, hydrogen being disengaged and the metal dissolved, but ammonia has no action on it. These properties of metallic beryllium seem to isolate it from the series of the other metals described in this chapter, but if we compare the properties of calcium, magnesium, and beryllium we shall see that magnesium occupies a position intermediate between the other two. Whilst calcium decomposes water with great ease, magnesium does so with difficulty, and beryllium not at all. The peculiarities of beryllium among the metals of the alkaline earths recall the fact that in the series of the halogens we saw that fluorine differed from the other halogens in many of its properties and had the smallest atomic weight. The same is the case with regard to beryllium among the other metals of the alkaline earths.

[59] We refer beryllium to the class of the bivalent metals of the alkaline earths--that is, we ascribe to its oxide the formula BeO, and do not consider it as trivalent (Be = 13·5, Chapter VII., Note 21), although that view has been upheld by many chemists. The true atomic composition of beryllium oxide was first given by the Russian chemist, Avdéeff (1819), in his researches on the compounds of this metal. He compared the compounds of beryllium to those of magnesium, and refuted the notion prevalent at the time, of the resemblance between the oxides of beryllium and aluminium, by proving that beryllium sulphate presents a greater resemblance to magnesium sulphate than to aluminium sulphate. It was especially noticed that the analogues of alumina give alums, whilst beryllium oxide, although it is a feeble base, easily giving, like magnesia, basic and double salts, does not form true alums. The establishment of the periodic system of the elements (1869), considered in the following chapter, immediately indicated that Avdéeff's view corresponded with the truth--that is, that beryllium is bivalent, which therefore necessitated the denial of its trivalency. This scientific controversy resulted in a long series of researches (1870-80) concerning this element, and ended in Nilson and Pettersson--two of the chief advocates of the trivalency of beryllium--determining the vapour density of BeCl_{2} = 40, (Chapter VII., Note 21), which gave an undoubted proof of its bivalency (_see_ also Note 3).

[60] Beryllium oxide, like aluminium oxide, is precipitated from solutions of its salts by alkalis as a gelatinous hydroxide, BeH_{2}O_{2}, which, like alumina, is soluble in an excess of caustic potash or soda. This reaction may be taken advantage of for distinguishing and separating beryllium from aluminium, because when the alkaline solution is diluted with water and boiled, beryllium hydroxide is precipitated, whilst the alumina remains in solution. The solubility of the beryllium oxide at once clearly indicates its feeble basic properties, and, as it were, separates this oxide from the class of the alkaline earths. But on arranging the oxides of the above-described metals of the alkaline earths according to their decreasing atomic weights we have the series

BaO, SrO, CaO, MgO, BeO,

in which the basic properties and solubility of the oxides consecutively and distinctly decrease until we reach a point when, had we not known of the existence of the beryllium oxide, we should expect to find in its place an oxide insoluble in water and of feeble basic properties. If an alcoholic solution of caustic potash be saturated with the hydrate of BeO, and evaporated under the receiver of an air pump, it forms silky crystals BeK_{2}O_{2}.

Another characteristic of the salts of beryllium is that they give with aqueous ammonia a gelatinous precipitate which is soluble in an excess of ammonium carbonate like the precipitate of magnesia; in this beryllium oxide differs from the oxide of aluminium. Beryllium oxide easily forms a carbonate which is insoluble in water, and resembles magnesium carbonate in many respects. Beryllium sulphate is distinguished by its considerable solubility in water--thus, at the ordinary temperature it dissolves in an equal weight of water; it crystallises out from its solutions in well-formed crystals, which do not change in the air, and contain BeSO_{4},4H_{2}O. When ignited it leaves beryllium oxide, but this oxide, after prolonged ignition, is re-dissolved by sulphuric acid, whilst aluminium sulphate, after a similar treatment, leaves aluminium oxide, which is no longer soluble in acids. With a few exceptions, the salts of beryllium crystallise with great difficulty, and to a considerable extent resemble the salts of magnesium; thus, for instance, beryllium chloride is analogous to magnesium chloride. It is volatile in an anhydrous state, and in a hydrated state it decomposes, with the evolution of hydrochloric acid.

In addition to the above characteristics of the compounds of the metals of the alkaline earths, we must add that they, like the alkali metals, combine with nitrogen and hydrogen, and while sodium nitride (obtained by igniting the amide of sodium, Chapter XII., Note 44 bis) and lithium nitride (obtained by heating lithium in nitrogen, Chapter XIII., Note 39) have the composition R_{3}N, so the nitrides of magnesium (Note 14), calcium, strontium, and barium have the composition R_{3}N_{2}, for example, Ba_{3}N_{2}, as might be expected from the diatomicity of the metals of the alkaline earths and from the relation of the nitrides to ammonia, which is obtained from all of these compounds by the action of water. The _nitrides_ of Ca, Sr, and Ba are formed directly (Maquenne, 1892) by heating the metals in nitrogen. They all have the appearance of an amorphous powder of dark colour; as regards their reactions, it is known that besides disengaging ammonia with water, they form cyanides when heated with carbonic oxide; for instance, Ba_{3}N_{2} + 2CO = Ba(CN)_{2} + 2BaO.[61]

[61] Thus in the nitrides of the metals we have substances by means of which we can easily obtain from the nitrogen of the air, not only ammonia, but also with the aid of CO, by synthesis, a whole series of complex carbon and nitrogen compounds.

The metals of the alkaline earths, just like Na and K, absorb hydrogen under certain conditions, and form pulverulent easily oxidisable metallic hydrides, whose composition corresponds exactly to that of Na_{2}H and K_{2}H, with the substitution of K_{2} and Na_{2} by the atoms Be, Mg, Ca, Sr, and Ba. The _hydrides of the metals of the alkaline earths_ were discovered by C. Winkler (1891) in investigating the reducibility of these metals by magnesium. In reducing their oxides by heating them with magnesium powder in a stream of hydrogen, Winkler observed that the hydrogen was absorbed (but very slowly), _i.e._ at the moment of their separation all the metals of the alkaline earths combine with hydrogen. This absorptive power increases in passing from Be to Mg, Ca, Sr, and Ba, and the resultant hydrides retain the combined hydrogen[62] when heated, so that these hydrides are distinguished for their considerable stability under heat, but they oxidise very easily.[63]

[62] As the hydrides of calcium, magnesium, &c. are very stable under the action of heat, and these metals and hydrogen occur in the sun, it is likely that the formation of their hydrides may take place there. (Private communication from Prof. Winkler, 1894.) It is probable that in the free metals of the alkaline earths hitherto obtained a portion was frequently in combination with nitrogen and hydrogen.

[63] Thus, for instance, a mixture of 56 parts of CaO and 24 parts of magnesium powder is heated in an iron pipe (placed over a row of gas burners as in the combustion furnace used for organic analysis) in a stream of hydrogen. After being heated for 1/2 hour the mixture is found to absorb hydrogen (it no longer passes over the mixture, but is retained by it). The product, which is light grey, and slightly coherent, disengages a mass of hydrogen when water is poured over it, and burns when heated in air. The resultant mass contains 33 per cent. CaH, about 28 per cent. CaO, and about 38 per cent. MgO. Neither CaH nor any other MH has yet been obtained in a pure state.

The acetylene derivatives of the metals of the alkaline earths C_{2}M (Chapter VIII., Note 12 bis), for instance, C_{2}Ba, obtained by Maquenne and Moissan, belong to the same class of analogous compounds. It must here be remarked that the oxides MO of the metals of the alkaline earths, although not reducible by carbon at a furnace heat, yet under the action of the heat attained in electrical furnaces, not only give up their oxygen to carbon (probably partly owing to the action of the current), but also combine with carbon. The resultant compounds, C_{2}M, evolve acetylene, C_{2}H_{2}, with HCl, just as N_{2}M_{3} give ammonia. We may remark moreover that the series of compounds of the metals of the alkaline earths with hydrogen, nitrogen and carbon is a discovery of recent years, and that probably further research will give rise to similar unexpected compounds, and by extending our knowledge of their reactions prove to be of great interest.

Thus the analogies and correlation of the metals of these two groups are now clearly marked, not only in their behaviour towards oxygen, chlorine, acids, &c., but also in their capability of combining with nitrogen and hydrogen.

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