Chapter X: , Note 2 (3)

_Lithium_, Li, is, like potassium and sodium, somewhat widely spread in siliceous rocks, but only occurs in small quantities and as mere traces in considerable masses of potassium and sodium salts. Only a very few rather rare minerals contain more than traces of it,[39] for example, spodumene and lithia mica. Many compounds of lithium are in all respects closely analogous to the corresponding compounds of sodium and potassium; but the _carbonate_ is sparingly soluble in cold water, which fact is taken advantage of for separating lithium from potassium and sodium. This salt, Li_{2}CO_{3}, is easily converted into the other compounds of lithium. Thus, for instance, the lithium hydroxide, LiHO, is obtained in exactly the same way as caustic soda, by the action of lime on the carbonate, and it is soluble in water and crystallises (from its solution in alcohol) as LiHO,H_{2}O. Metallic _lithium_ is obtained by the action of a galvanic current on fused lithium chloride; for this purpose a cast-iron crucible, furnished with a stout cover, is filled with lithium chloride, heated until the latter fuses, and a strong galvanic current is then passed through the molten mass. The positive pole (fig. 77) consists of a dense carbon rod C (surrounded by a porcelain tube P fixed in an iron tube BB), and the negative pole of an iron wire, on which the metal is deposited after the current has passed through the molten mass for a certain length of time. Chlorine is evolved at the positive pole. When a somewhat considerable quantity of the metal has accumulated on the wire it is withdrawn, the metal is collected from it, and the experiment is then carried on as before.[39 bis] Lithium is the lightest of all metals, its specific gravity is 0·59, owing to which it floats even on naphtha; it melts at 180°, but does not volatilise at a red heat. Its appearance recalls that of sodium, and, like it, it has a yellow tint. At 200° it burns in air with a very bright flame, forming lithium oxide. In decomposing water it does not ignite the hydrogen. The characteristic test for lithium compounds is the _red coloration_ which they impart to a colourless flame.[40]

[39] Thus _spodumene_ contains up to 6 p.c. of lithium oxide, and
_petolite_, and _lepidolite_ or lithia mica, about 3 p.c. of
lithium oxide. This mica is met with in certain granites in a
somewhat considerable quantity, and is therefore most frequently
employed for the preparation of lithium compounds. The treatment
of lepidolite is carried on on a large scale, because certain
salts of lithium are employed in medicine as a remedy for certain
diseases (stone, gouty affections), as they have the power of
dissolving the insoluble uric acid which is then deposited.
Lepidolite, which is unacted on by acids in its natural state,
decomposes under the action of strong hydrochloric acid after it
has been fused. After being subjected to the action of the
hydrochloric acid for several hours all the silica is obtained in
an insoluble form, whilst the metallic oxides pass into solution
as chlorides. This solution is mixed with nitric acid to convert
the ferrous salts into ferric, and sodium carbonate is then added
until the liquid becomes neutral, by which means a precipitate is
formed of the oxides of iron, alumina, magnesia, &c., as insoluble
oxides and carbonates. The solution (with an excess of water) then
contains the chlorides of the alkaline metals KCl, NaCl, LiCl,
which do not give a precipitate with sodium carbonate in a dilute
solution. It is then evaporated, and a strong solution of sodium
carbonate added. This precipitates lithium carbonate, which,
although soluble in water, is much less so than sodium carbonate,
and therefore the latter precipitates lithium from strong
solutions as carbonate, 2LiCl + Na_{2}CO_{3} = 2NaCl +
Li_{2}CO_{3}. _Lithium carbonate_, which resembles sodium
carbonate in many respects, is a substance which is very slightly
soluble in cold water and is only moderately soluble in boiling
water. In this respect lithium forms a transition between the
metals of the alkalis and other metals, especially those of the
alkaline earths (magnesium, barium), whose carbonates are only
sparingly soluble. Oxide of lithium, Li_{2}O, may be obtained by
heating lithium carbonate with charcoal. Lithium oxide in
dissolving gives (per gram-molecule) 26,000 heat units; but the
combination of Li_{2} with O evolves 140,000 calories--that is,
more than Na_{2}O (100,000 calories) and K_{2}O (97,000 calories),
as shown by Beketoff (1887). Oeuvrard (1892) heated lithium to
redness in nitrogen, and observed the absorption of N and
formation of Li_{3}N, like Na_{3}N (_see_ Chapter XII. Note 50).

LiCl, LiBr, and LiI form crystallo-hydrates with H_{2}O, 2H_{2}O,
and 3H_{2}O. As a rule, LiBr,2H_{2}O crystallises out, but
Bogorodsky (1894) showed that a solution containing LiBr +
3·7H_{2}O, cooled to -62°, separates out crystals LiBr,3H_{2}O,
which decompose at +4° with the separation of H_{2}O. LiF is but
slightly soluble (in 800 parts) in water (and still less so in a
solution of NH_{4}F).

[39 bis] Guntz (1893) recommends adding KCl to the LiCl in preparing Li
by this method, and to act with a current of 10 ampères at 20
volts, and not to heat above 450°, so as to avoid the formation of
Li_{2}Cl.

[40] In determining the presence of lithium in a given compound, it is
best to treat the material under investigation with acid (in the
case of mineral silicon compounds hydrofluoric acid must be
taken), and to treat the residue with sulphuric acid, evaporate to
dryness, and extract with alcohol, which dissolves a certain
amount of the lithium sulphate. It is easy to discover lithium in
such an alcoholic solution by means of the coloration imparted to
the flame on burning it, and in case of doubt by investigating its
light in a spectroscope, because lithium gives a red line, which
is very characteristic and is found as a dark line in the solar
spectrum. Lithium was first discovered in 1817 in petolite by
Arfvedson.

Bunsen in 1860 tried to determine by means of the spectroscope whether any other as yet unknown metals might not occur in different natural products together with lithium, potassium, and sodium, and he soon discovered two new alkali metals showing independent spectra. They are named after the characteristic coloration which they impart to the flame. One which gives a red and violet band is named _rubidium_, from _rubidius_ (dark red), and the other is called _cæsium_, because it colours a pale flame sky blue, which depends on its containing bright blue rays, which appear in the spectrum of cæsium as two blue bands (table on p. 565). Both metals accompany sodium, potassium, and lithium, but in small quantities; rubidium occurs in larger quantity than cæsium. The amount of the oxides of cæsium and rubidium in lepidolite does not generally exceed one-half per cent. Rubidium has also been found in the ashes of many plants, while the Stassfurt carnallite (the mother-liquor obtained after having been treated for KCl) forms an abundant source for rubidium and also partly for cæsium. Rubidium also occurs, although in very small quantities, in the majority of mineral waters. In a very few cases cæsium is not accompanied by rubidium; thus, in a certain granite on the Isle of Elba, cæsium has been discovered, but not rubidium. This granite contains a very rare mineral called _pollux_, which contains as much as 34 per cent. of cæsium oxide. Guided by the spectroscope, and aided by the fact that the double salts of platinic chloride and rubidium and cæsium chlorides are still less soluble in water than the corresponding potassium salt, K_{2}PtCl_{6},[41] Bunsen succeeded in separating both metals from each other and from potassium, and demonstrated the great resemblance they bear to each other. The isolated metals,[42] rubidium and cæsium, have respectively the specific gravities 1·52 and 2·366, and melting points 39° and 27° as N. N. Beketoff showed (1894), he having obtained cæsium by heating CsAlO_{2} with Mg([42 bis]).

[41] The salts of the majority of metals are precipitated as carbonates
on the addition of ammonium carbonate--for instance, the salts of
calcium, iron, &c. The alkalis whose carbonates are soluble are
not, however, precipitated in this case. On evaporating the
resultant solution and igniting the residue (to remove the
ammonium salts), we obtain salts of the alkali metals. They may he
separated by adding hydrochloric acid together with a solution of
platinic chloride. The chlorides of lithium and sodium give easily
soluble double salts with platinic chloride, whilst the chlorides
of potassium, rubidium, and cæsium form double salts which are
sparingly soluble. A hundred parts of water at 0° dissolve 0·74
part of the potassium platinochloride; the corresponding rubidium
platinochloride is only dissolved to the amount of 0·134 part, and
the cæsium salt, 0·024 part; at 100° 5·13 parts of potassium
platinochloride, K_{2}PtCl_{6}, are dissolved, 0·634 part of
rubidium platinochloride, and 0·177 part of cæsium
platinochloride. From this it is clear how the salts of rubidium
and cæsium may be isolated. The separation of cæsium from rubidium
by this method is very tedious. It can be better effected by
taking advantage of the difference of the solubility of their
carbonates in alcohol; cæsium carbonate, Cs_{2}CO_{3}, is soluble
in alcohol, whilst the corresponding salts of rubidium and
potassium are almost insoluble. Setterberg separated these metals
as alums, but the best method, that given by Scharples, is founded
on the fact that from a mixture of the chlorides of potassium,
sodium, cæsium, and rubidium in the presence of hydrochloric acid,
stannic chloride precipitates a double salt of cæsium, which is
very slightly soluble. The salts of Rb and Cs are closely
analogous to those of potassium.

[42] Bunsen obtained rubidium by distilling a mixture of the tartrate
with soot, and Beketoff (1888) by heating the hydroxide with
aluminium, 2RbHO + Al = RbAlO_{2} + H_{2} + Rb. By the action of
85 grams of rubidium on water, 94,000 heat units are evolved.
Setterberg obtained cæsium (1882) by the electrolysis of a fused
mixture of cyanide of cæsium and of barium. Winkler (1890) showed
that metallic magnesium reduces the hydrates and carbonates of Rb
and Cs like the other alkaline metals. N. N. Beketoff obtained
them with aluminium (see following note).

[42 bis] Beketoff (1888) showed that metallic aluminium reduces the
hydrates of the alkaline metals at a red heat (they should be
perfectly dry) with the formation of aluminates (Chapter XVII.),
RAlO_{2}--for example, 2KHO + Al = KAlO_{2} + K + H_{2}. It is
evident that in this case only half of the alkaline metal is
obtained free. On the other hand, K. Winkler (1889) showed that
magnesium powder is also able to reduce the alkaline metals from
their hydrates and carbonates. N. N. Beketoff and Tscherbacheff
(1894) prepared cæsium upon this principle by heating its
aluminate CsAlO_{2} with magnesium powder. In this case aluminate
of magnesium is formed, and the whole of the cæsium is obtained as
metal: 2CsAlO_{2} + Mg = MgOAl_{2}O_{5} + 2Cs. A certain excess of
alumina was taken (in order to obtain a less hygroscopic mass of
aluminate), and magnesium powder (in order to decompose the last
traces of water); the CsAlO_{2} was prepared by the precipitation
of cæsium alums by caustic baryta, and evaporating the resultant
solution. We may add that N. N. Beketoff (1887) prepared oxide of
potassium, K_{2}O, by heating the peroxide, KO, in the vapour of
potassium (disengaged from its alloy with silver), and showed that
in dissolving in an excess of water it evolves (for the
above-given molecular weight) 67,400 calories (while 2KHO in
dissolving in water evolves 24,920 cal.; so that K_{2}O + H_{2}O
gives 42,480 cal.), whence (knowing that K_{2} + O+H_{2}O in an
excess of water evolves 164,500) it follows that K_{2} + O evolves
97,100 cal. This quantity is somewhat less than that (100,260
cal.) which corresponds to sodium, and the energy of the action of
potassium upon water is explained by the fact that K_{2}O evolves
more heat than Na_{2}O in combining with water (_see_ Chapter II.
Note 9). Just as hydrogen displaces half the Na from Na_{2}O
forming NaHO, so also N. N. Beketoff found from experiment and
thermochemical reasonings that hydrogen displaces half the
potassium from K_{2}O forming KHO and evolving 7,190 calories.
Oxide of lithium, Li_{2}O, which is easily formed by igniting
Li_{2}CO_{3} with carbon (when Li_{2}O + 2CO is formed),
disengages 26,000 cals. with an excess of water, while the
reaction Li_{2} + O gives 114,000 cals. and the reaction Li_{2} +
H_{2}O gives only 13,000 cals., and metallic lithium cannot be
liberated from oxide of lithium with hydrogen (nor with carbon).
Thus in the series Li, Na, K, the formation of R_{2}O gives most
heat with Li and least with K, while the formation of RCl evolves
most heat with K (105,000 cals.) and least of all with Li (93,500
cals.). Rubidium, in forming Rb_{2}O, gives 94,000 cals.
(Beketoff). Cæsium, in acting upon an excess of water, evolves
51,500 cals., and the reaction Cs_{2} + O evolves about 100,000
cals.--_i.e._ more than K and Rb, and almost as much as Na--and
oxide of cæsium reacts with hydrogen (according to the equation
Cs_{2}O + H = CsHO + Cs) more easily than any of the oxides of the
alkali metals, and this reaction takes place at the ordinary
temperature (the hydrogen is absorbed), as Beketoff showed (1893).
He also obtained a mixed oxide, AgCsO, which was easily formed in
the presence of silver, and absorbed hydrogen with the formation
of CsHO.

Judging by the properties of the free metals, and of their corresponding and even very complex compounds, lithium, sodium, potassium, rubidium, and cæsium present an indubitable chemical resemblance. The fact that the metals easily decompose water, and that their hydroxides RHO and carbonates R_{2}CO_{3} are soluble in water, whilst the hydroxides and carbonates of nearly all other metals are insoluble, shows that these metals form a natural group of _alkali metals_. The halogens and the alkali metals form, by their character, the two extremes of the elements. Many of the other elements are metals approaching the alkali metals, both in their capacity of forming salts and in not forming acid compounds, but are not so energetic as the alkali metals, that is, they form less energetic bases. Such are the common metals, silver, iron, copper, &c. Some other elements, in the character of their compounds, approach the halogens, and, like them, combine with hydrogen, but these compounds do not show the energetic property of the halogen acids; in a free state they easily combine with metals, but they do not then form such saline compounds as the halogens do--in a word, the halogen properties are less sharply defined in them than in the halogens themselves. Sulphur, phosphorus, arsenic, &c. belong to this order of elements. The clearest distinction of the properties of the halogens and alkali metals is expressed in the fact that the former give acids and do not form bases, whilst the latter, on the contrary, only give bases. The first are true _acid elements_, the latter clearly-defined _basic or metallic elements_. On combining together, the halogens form, in a chemical sense, unstable compounds, and the alkali metals alloys in which the character of the metals remains unaltered, just as in the compound ICl the character of the halogens remains undisguised; thus both classes of elements on combining with members of their own class form non-characteristic compounds, which have the properties of their components. On the other hand, the halogens on combining with the alkali metals form compounds which are, in all respects, stable, and in which the original characters of the halogens and alkali metals have entirely disappeared. The formation of such compounds is accompanied by evolution of a large amount of heat, and by an entire change of both the physical and chemical properties of the substances originally taken. The alloy of sodium and potassium, although liquid at the ordinary temperature, is perfectly metallic, like both its components. The compound of sodium and chlorine has neither the appearance nor the properties of the original elements; sodium chloride melts at a higher temperature, and is more difficultly volatile, than either sodium or chlorine.

With all these qualitative differences there is, however, an important quantitative _resemblance between the halogens and the alkali metals_. This resemblance is clearly expressed by stating that both orders of elements belong to those which are univalent with respect to hydrogen. It is thus correct to say that both the above-named orders of elements replace hydrogen atom for atom. Chlorine is able to take the place of hydrogen by metalepsis, and the alkali metals take the place of hydrogen in water and acids. As it is possible to consecutively replace every equivalent of hydrogen in a hydrocarbon by chlorine, so it is possible in an acid containing several equivalents of hydrogen to replace the hydrogen consecutively equivalent after equivalent by an alkali metal; hence an atom of these elements is analogous to an atom of hydrogen, which is taken, in all cases, as the unit for the comparison of the other elements. In ammonia, and in water, chlorine and sodium are able to bring about a direct replacement. According to the law of substitution, the formation of sodium chloride, NaCl, at once shows the equivalence of the atoms of the alkali metals and the halogens. The halogens and hydrogen and the alkali metals combine with such elements as oxygen, and it is easily proved that in such compounds one atom of oxygen is able to retain two atoms of the halogens, of hydrogen, and of the alkali metals. For this purpose it is enough to compare the compounds KHO, K_{2}O, HClO, and Cl_{2}O, with water. It must not be forgotten, however, that the halogens give, with oxygen, besides compounds of the type R_{2}O, higher acid grades of oxidation, which the alkali metals and hydrogen are not capable of forming. We shall soon see that these relations are also subject to a special law, showing a gradual transition of the properties of the elements from the alkali metals to the halogens.[43]

[43] We may here observe that the halogens, and especially iodine,
may play the part of metals (hence iodine is more easily replaced
by metals than the other halogens, and it approaches nearer to the
metals in its physical properties than the other halogens).
Schützenberger obtained a compound C_{2}H_{3}O(OCl), which he
called chlorine acetate, by acting on acetic anhydride,
(C_{2}H_{3}O)_{2}O, with chlorine monoxide, Cl_{2}O. With iodine
this compound gives off chlorine and forms iodine acetate,
C_{2}H_{3}O(OI), which also is formed by the action of iodine
chloride on sodium acetate, C_{2}H_{3}O(ONa). These compounds are
evidently nothing else than mixed anhydrides of hypochlorous and
hypoiodous acids, or the products of the substitution of hydrogen
in RHO by a halogen (_see_ Chapter XI., Notes 29 and 78 bis). Such
compounds are very unstable, decompose with an explosion when
heated, and are changed by the action of water and of many other
reagents, which is in accordance with the fact that they contain
very closely allied elements, as does Cl_{2}O itself, or ICl or
KNa. By the action of chlorine monoxide on a mixture of iodine and
acetic anhydride, Schützenberger also obtained the compound
I(C_{2}H_{3}O_{2})_{3}, which is analogous to ICl_{3}, because the
group C_{2}H_{3}O_{2} is, like Cl, a halogen, forming salts with
the metals. Similar properties are found in iodosobenzene (Chapter
XI., Note 79).

The atomic weights of the alkali metals, lithium 7, sodium 23, potassium 39, rubidium 85, and cæsium 133, show that here, as in the class of halogens, the elements may be arranged according to their atomic weights in order to compare the properties of the analogous compounds of the members of this group. Thus, for example, the platinochlorides of lithium and sodium are soluble in water; those of potassium, rubidium, and cæsium sparingly soluble, and the greater the atomic weight of the metal the less soluble is the salt. In other cases the reverse is observed--the greater the atomic weight the more soluble are the corresponding salts. The variation of properties with the variation in atomic weights even shows itself in the metals themselves; thus lithium volatilises with difficulty, whilst sodium is obtained by distillation, potassium volatilises more easily than sodium, and rubidium and cæsium as we have seen, are still more volatile.